1st Year Chem Notes 2018.pdf

PROFESSOR SERIES

Professor easy notes for

11 CHEMISTRY TH

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Written By: Muhammad Shahid

Muhammad Shahid (M.Sc Chemistry)

A-One Professors Academy of Sciences Jauharabad Designed by: Mukhtar Graphics

Muhammad Asif Mukhtar (CA & OP)

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A-One Professors Academy of sciences ِِ ‫پ‬ ‫نزد ُ ی‬ )‫سزریئاکجلوجرہٓاابد(وخاشب‬

All Rights reserved: No part of this publication may be reproduced, stored in a retrieval system, or transmitted by any means, mechanical, electronic, recording, photocopies, or otherwise without the prior written permission of the publisher.

Acknowledgement: A tribute to the Great teacher, The Great personality & The Great human Prof. Khwaja Khalid Iqbal (late) Associate Professor Govt. Post Graduate College Jauharabad. The founder of

A-One Professors Academy of Sciences Jauharabad

A Brief introduction of academy: The academy was started by two great professors. Sir Khwaja Khalid (Late) Associate professor of Chemistry and Sir Akhtar Hussain (Late) Associate Professor of Physics in 1996 in Jauharabad District Khushab.

Preface: The subject of chemistry has undergone a remarkable advancement at national and international levels. These notes are according to NEW SYLLABUS & PATTERN recommended for 11th class.

Important features of notes:  These notes covers not only the whole book but also prepares students for the examination according to new pattern.  All the exercises & important questions of textbook have been solved in an easy an understandable way.  Short and long questions of previous SARGODHA BOARD papers are included i.e.2007-2018.  I wish you all the best of luck for your exams.

Do you know? There are five states of matter.    

Solid Liquid Gas Plasma (Matter exists at very high temperature)  Bose Einstein condensate (Matter exists at very low temperature)

Contents: Chapters

Sr. NO Unit 1

Pages

1

Unit 3

Basic concepts Experimental techniques in chemistry Gases

17

Unit 4

Liquids and solids

29

Unit 5

Atomic structure

45

Unit 6

Chemical bonding

55

Unit 7

Thermochemistry

67

Unit 8

Chemical equilibrium

75

Unit 9

Solutions

83

Unit 10

Electrochemistry

91

Unit 11

Reaction kinetics

101

Unit 2

9

Things to remember:

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 1

“Basic concepts” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 1 BASIC CONCEPTS

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PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) (1) What is atomic mass unit (a.m.u)? Give its value in grams. (2008) (2009) The unit used to express the relative atomic mass is called atomic mass unit. It is 1/12th of the mass of one carbon atom. 1 amu = 1.661× 10-27 kg or 1 amu = 1.661× 10-24 g (2) Law of conversation of mass has to be considered during stoichiometric calculations. (2008) Stoichiometric calculations are based on balanced chemical equation. During this it is assumed that matter is neither created nor destroyed but changes from one form to another. So Law of conversation of mass has to be considered during stoichiometric calculations otherwise with unbalanced chemical equations stoichiometric calculations will not be possible. (3) One mole of N2, CO2 and H2 contains equal number of molecules. Give reason. (2008) 1 mole of CO2 = 6.023 × 1023 molecules 1 mole of N2 = 6.023 × 1023 molecules 1 mole H2 = 6.023 × 1023 molecules Since all the gases has same number of moles i.e. N2, H2 and CO2 have their one mole so they contains equal number of molecules. (4) No individual neon atom in the sample of the element has a mass of 20.18 amu. (2017) (2018) OR Atomic masses may be expressed in fractions, why? (2009) 20

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A sample of neon consists of 10Ne, 10Ne and 10Ne in the percentages of 90.92%, 0.26% and 8.82% respectively. The average atomic mass of Ne can be calculated as (20×90.92)+(21×0.26)+(22×8.82)

Average atomic mass = = 20.18 amu 100 Hence 20.18 amu is the average atomic mass of neon and no individual neon in the sample that has a mass of 20.18 amu. (5) NaCl has 58.5 amu as formula mass and not the molecular mass, justify it? (2009) NaCl is an ionic compound and does not exist in molecular form. Therefore NaCl is a formula unit and has 58.5 amu as formula mass and not the molecular mass. (6) Define molecular ion. How it is generated? OR What is molecular ion? Give one example. (2009) (2010) (2012) (2013) When a molecule loses or gains an electron, molecular ion is formed. For example CH4+, CO+, N2+. Cationic molecular, ions are more abundant than anionic ions. These can be generated by passing high energy electron beam as α-particles or X-rays through a gas. The breakdown of molecular ions obtained from the natural products can give important information about their structure. (7) Calculate the mass in grams of 2.74 moles of KMnO4. (2010) No of moles of KMnO4 = 2.74 moles Molar mass of KMnO4 = 158 g/mol [(39) + (55) + (16×4) = 158] Mass of KMnO4 = No. of moles × Molar mass of KMnO4 Muhammad Shahid = 2.74 × 158 = 432.92gram A-One Professors Academy of (8) Define ions. Give two examples. (2011) Sciences Jauharabad OR What are ions? Under what conditions they are produced? (2017) Ions are those species which carry either positive or negative charge. By ionization of an ionic compound in water. e.g. NaCl Na+ + ClPositive ions are produced by passing high energy electron beam, α-particles or X-rays through a gas. e.g. Na  Na+ + e-

CHAPTER # 1 BASIC CONCEPTS

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Negative ions are produced by the addition of an electron to a neutral species. e.g. Cl + e-  Cl(9) Write the formulas used for the determination of percentage of C & H. (2012) The percentage of carbon can be calculated by using following formula: Mass of CO2 12.00 % of carbon = × × 100 Mass of organic compund

44.00

The percentage of hydrogen can be calculated by using following formula: Mass of H2 O 2.016 % of hydrogen = × × 100 Mass of organic compund

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(10) What is limiting reactant? How it is identified?(2012) (2013) (2015) The reactant which controls (limits) the amount of product formed during a chemical reaction is called limiting reactant. It can be identified by following steps:  Calculate the number of moles from the given amount of reactant.  Find out the number of moles of the product with the help of a balanced chemical equation.  Identify the reactant which produces the least amount of product as limiting reactant. (11) What is stoichiometry? Give its assumptions. (2012) (2013) (2016) Stoichiometry is a branch of chemistry which gives a quantitative relationship between reactants and products in a balanced chemical equation. It has two assumptions:  All the reactants are completely converted into the products.  No side reaction occurs. (12) What is Avogadro’s number? Give an example. (2012) Avogadro’s number is the number of atoms, molecules and ions in one gram atom of an element, one gram molecule of a compound and one gram ions of a substance, respectively. For example 1.008 g of Hydrogen = 1 mole of H = 6.02 × 1023 atoms of H 18 g of H2O = 1 mole of H2O = 6.02 × 1023 molecules of H2O 2296 g of SO4 = 1 mole of SO4 = 6.02 × 1023 ions of SO42(13) Molecular formula is integral multiple of empirical formula. Give comments. (2012) The empirical formula gives only the simplest whole number ratio of atoms of each element in a compound, while molecular formula tells the actual number of atoms of each element in a molecule. So empirical formula is multiplied with a suitable integer “n” to get the molecular formula. Where “n” is the ratio molecular mass and empirical formula mass. For example benzene 6 × (CH) = C6H6 (14) Define (a) Limiting reactant (b) Atomicity. (2012) Limiting reactant: The reactant which controls (limits) the amount of product formed during a chemical reaction is called limiting reactant. For example burning of coal to form CO2. Coal is limiting reactant. Atomicity: The number of atoms present in a molecule is called its atomicity. For example the atomicity of O2 is two. (15) What is meant by atomicity? Explain with two examples. (2015) (2017) The number of atoms present in a molecule is called its atomicity. The molecule can be monoatomic, diatomic and triatomic etc. If the molecule contains one atom it is monoatomic, if it contains two atoms it is diatomic, and if it contains three atoms it is triatomic. Molecules of elements may contain one two or more same type of atoms. For example He, Cl2, O3, P4, S8. The molecules of compounds consist of different kind of atoms. For example HCl, NH3, H2SO4, C6H12O6. (16) A compound may have same empirical and molecular formula. Justify it. (2015) There are many compounds, which have same empirical formula and molecular formula. For example H2O, CO2, NH3 and C12H22O11 have the same empirical and molecular formulas. Their simple multiple “n” is unity. Actually value of “n” is the ratio of molecular mass and empirical formula mass.

CHAPTER # 1 BASIC CONCEPTS

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(17) Distinguish b/w actual yield and theoretical yield. (2015) Actual Yield: The amount of the products obtained as a result of a chemical reaction is called actual yield. It is based on experiment. Theoretical (Experiment) Yield: The amount of the products calculated from the balanced chemical equation is called theoretical yield. (18) What is percentage yield? Give its significance. (2015) Percentage yield is actually the efficiency of a reaction. It tell us how much product we have obtained after a chemical reaction. It can be obtained by comparing theoretical and actual yield in the form of the percentage yield. Actual Yield % yield = × 100 Theoretical yield (19) Define mass spectrometer. (2016) Mass spectrometer is an instrument which is used to measure the exact masses of different isotopes of an element. (20) How many molecules of water are there in 10 gram of ice? (2016) Mass in gram Number of molecules (Ice) = × NA Muhammad Shahid Number of molecules (Ice) =

Molar mass 10 18

× 6.02 × 1023

A-one Professors Academy of sciences Jauharabad

Number of molecules (Ice) = 3.34 × 1023 molecules (21) What is the justification of two peaks in the mass spectrum of bromine, while for iodine only one peak at 127 a.m.u is indicated? (2017) (1) Bromine has two naturally occurring isotopes with almost equal relative abundance therefore; its mass spectrum shows two strong peaks. (2) While iodine has only one naturally occurring isotopes, therefore, its spectrum shows only one peak. (22) Why theoretical yield of a chemical reaction is greater than the actual yield? (2017) Theoretical yield of a chemical reaction is greater than the actual yield because it is calculated from a balanced chemical equation where we supposed that all the reactants are converted into products and no side reaction occurred. (23) Why the actual yield is always less than the theoretical yield? (2017) (2018) The actual yield is always less than theoretical yield due to following reasons.  Inexperience worker wastes a significant amount of product.  The reversible reactions are never completed.  Side reaction may occur.  All the reactant may not be converted into products.  Mechanical loss may occur like during e.g. Filtration, evaporation, crystallization, distillation etc. (24) Write only four names of any four methods employed for the separation of isotopes. (2018) The separation of isotopes can be done by using following methods.  Gaseous diffusion  Thermal diffusion  Distillation  Ultracentrifuge  Electromagnetic separation  Laser separation (25) Define gram atom by giving an example. (2018) The atomic mass of an element expressed in grams is called gram atom. For example 1 gram atom of Hydrogen = 1.008 g (26) Write down limitations of a chemical equation. (2018) The chemical equation has following limitations They don’t tell about the conditions of reactions.

CHAPTER # 1 BASIC CONCEPTS

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They don’t tell us rate of reaction. These can also be written for such reactions that does not occur. These cannot tell about the time to complete the reaction. (27) What is the function of ionization chamber in mass spectrometer? (2018) In mass spectrometer the function of ionization chamber is to ionize the vapours of the substance. For this purpose vapours are passed into ionization chamber. In this fast moving electrons are thrown upon them. As a result gaseous atoms are ionized and positive ions are produced. (28) Calculate the mass of 10-3 moles of water? No of moles of water = 10-3 moles Molar mass of water = 18 g/mol Mass of water = No. of moles × Molar mass of water = 10-3 × 18 = 1.8 × 10-2 gram = 0.018 g. (29) What are isotopes? Atoms of the same element, having same atomic number but different mass number are called isotopes. The phenomenon of isotopy was discovered by Soddy. Examples 12 13 14 Carbon has three isotopes written as 6C, 6C, 6C and expressed as C-12, C-13, C-14, having 6, 7 and 8 neutrons in their nuclei respectively. Hydrogen has three isotopes: 1 1H

2

3

Protium 1H Deuterium 1H Tritium (30) Which information we can get from mass spectrum? In mass spectrometer, the result is recorded in the form of graph containing peaks. The relative numbers of ions are present along y-axis (ordinate) and m/e values of ions are present along x-axis (abscissa). This is called spectrum. It gives following information (1) The number of peaks gives the number of isotopes of an element. (2) The height of peaks gives the relative abundance of isotopes. (3) The position of peak gives the mass of the isotopes. (31) Justify the following statement:  23 gram of sodium and 238 gram of uranium have equal number of atoms in them. 1 mole of Na = 23g 1 mole of U = 238g Since one mole of each element contains Avogadro’s number of atoms and there is one mole of each of Na and U. hence 23g of Na and 238g of U contains equal number of atoms i.e. 6.02 × 1023.  Mg atom is twice heavier than that of carbon? One atom of carbon contains 6 protons and 6 neutrons in its nucleus and its atomic mass on atomic mass unit scale is 12 amu. While one atom of Mg contains 12 protons and 12 neutrons in its nucleus and its atomic mass on 24𝑎𝑚𝑢 atomic mass unit scale is 24amu. Thus, = 2. Hence, one atom of Mg is twice heavier than 12𝑎𝑚𝑢 that of one atom of Carbon.  180 gram of glucose and 342 gram of sucrose have same number of molecules but different number of atoms. 180g of glucose = 1 mole 342g of sucrose = 1 mole 1 mole of each compound contains Avogadro’s number of molecules. Hence, 180 g of glucose (1mole) and 342 g (1mole) of sucrose contain equal number of molecules i.e. 6.02 × 1023 Since one molecule of glucose (C6H12O6) contains 24 atoms. Whereas, one molecule of sucrose (C12H22O11) contains 45 atoms, therefore, equal number of molecule of glucose and sucrose will have different number of atoms.  4.9 g of H2SO4 when completely ionized in water have equal number of positive and negative charges but the number of positively charge ions is twice the number of negatively charge ions.

CHAPTER # 1 BASIC CONCEPTS

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(32) 

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PROFESSOR EASY NOTES

H2SO4 ionizes in solution as H2SO4 2H+ + SO42Muhammad Shahid This balance equation shows that A-one Professors Academy of sciences Jauharabad One mole of H2SO4 produces Number of positively charged ions (H+ ions) = 2 Number of negatively charged ions (SO42-) = 1 Number of positive charge = 2 (due to two H+ ions) Number of negative charge = 2 (due to two negative charge on SO42-) Hence whatever the amount of H2SO4, it will always produce equal number of positive and negative charges but number of positively charged ions will be twice the number of the negative charged ions. One mg of K2Cr2O7 has thrice the number of ions than the number of molecules when ionized in water. K2Cr2O7 ionizes as K2Cr2O7 2K+ + Cr2O72This equation shows that 1 formula unit of K2Cr2O7 produces two K+ ions and one Cr2O72- ion in solution. Thus a total of three ions are produced by the ionization of 1 formula unit of K2Cr2O7. Hence whatever be the amount of K2Cr2O7. The number of ions in its solution will always be thrice than the number of its formula units (molecules). Two grams of H2, 16 g of CH4 and 44 g of CO2 occupy separately the volume of 22.414 dm3, although the size and masses of molecules of these gases are very different from each other. 2 grams of H2 = 1 mole = 6.023 × 1023 molecules 16 grams of CH4 = 1 mole = 6.023 × 1023 molecules 44 grams of CO2 = 1 mole = 6.023 × 1023 molecules Since 2 grams of H2, 16 grams of CH4 and 44 grams of CO2 contains equal numbers of molecules and according to Avogadro’s law, equal molecules of all the gases at S.T.P. occupy the same volume i.e. 22.414 dm3. In gases distance between two molecules is approximately 300 times its own diameter. Thus, volume occupied by gas molecule does not depends upon the size or mass of molecules and it only depends upon the numbers of molecules. Hence equal molecules of H2, CH4 and CO2 at STP will occupy same volume i.e. 22.414 dm3. Explain the following with reasons Many chemical reactions taking place in our surrounding involve the limiting reactant. The reactant which controls (limits) the amount of product formed during a chemical reaction is called limiting reactant. In our surrounding there are many chemical reactions take place which involve limiting reactants some of these reactions are: (i) Burning of coal to form CO2. Coal is limiting reactant. (ii) Burning of natural gas to form CO2 and H2O. Here natural gas is our limiting reactant. (iii) Rusting of iron, iron is limiting reactant. In above reactions oxygen is always in excess, while other reactants are consumed earlier. So other reactants are limiting reactants. One mole of H2SO4 should completely react with two moles of NaOH. How does Avogadro’s number help to explain it? H2SO4 + 2NaOH  Na2SO4 + H2O In the above chemical reaction 1 molecule of H2SO4 requires = 2 formula units of NaOH. 23 1 × 6.02 × 10 molecules of H2SO4 requires = 2 × 6.02 × 1023 formula units of NaOH or 1 mole of H2SO4 requires = 2 moles of NaOH Hence, One mole of H2SO4 should completely react with two moles of NaOH. One mole of H2O has two moles of bond, three moles of atoms, ten moles of electrons and twenty eight moles of total fundamental particles present in it. One molecule of H2O has two bonds therefore one mole of H2O has 2 moles of bonds. One molecule of H2O has three atoms (Two Hydrogen atoms and one Oxygen atom) therefore one mole of H2O has three moles of atoms.

CHAPTER # 1 BASIC CONCEPTS

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One molecule of H2O has 10 electrons (2 electrons of 2 H and 8 electrons of 1 Oxygen atom) therefore one mole of H2O has 10 moles of electrons. One molecule of H2O has twenty eight fundamental particles i.e.10 electrons (2 e- of H + 8 e- of O) + 10 protons (2 protons of H + 8 protons of O) & 8 neutrons. Therefore one mole of H2O has 28 moles of total fundamental particles present in it.  N2 and CO has same number of electrons, protons and neutrons. Both N2 and CO have same number of electrons, protons and neutrons as it is clear from the following explanation. For N2 Muhammad Shahid No. of electrons in N2 = 7 + 7 = 14 A-one Professors Academy of No. of protons in N2 = 7 + 7 = 14 sciences Jauharabad No. of neutrons = 7 + 7 = 14 For CO Number of electrons in C = 6 No. of electrons in O =8 Total no. of electrons = 6 + 8 = 14 No. of protons in C =6 No. of protons in O =8 Total no. of protons = 6 + 8 = 14 No. of neutrons in C =6 No. of neutrons in O =8 Total no. of neutrons = 6 + 8 = 14 Hence N2 and CO has same number of electrons, protons and neutrons. (33) Write down briefly the principle of mass spectrometry. In mass spectrometry a substance is first volatilized and then ionized with the help of high energy beam of electrons. The gaseous positive ions, thus formed, are separated on the basis of their mass to charge ratio (m/e) and then recorded in the form of peaks.

Important Long questions of previous board papers (1) What is limiting reactant? How can you determine it? Explain with a chemical reaction. (2008) OR (2) Define limiting reactant. How it controls the amount of products? Give example. (2009) (3) What is mass spectrometer? How it is used to determine the relative atomic masses? (2009) (4) Calculate the number of grams of K2SO4 and water produced when 14 g of KOH are reacted with excess of H2SO4, also calculate the number of molecules of water produced. (2010) (5) Write steps involved in determination of empirical formula of a compound? (2012) (6) NH3 gas can be prepared by heating together two solids NH4Cl and Ca(OH)2. If a mixture containing 100 gram of each solid is heated then, calculate the number of grams of NH3 produced. (2011) (2015) 2NH4Cl + Ca(OH)2  CaCl2 + 2NH3 + 2H2O (7) What is empirical formula? Discuss how empirical formula is determined by combustion analysis? (2014) (8) A sample of liquid consisting of carbon, hydrogen and oxygen was subjected to combustion analysis, 0.5439 g of the compound gave 1.039 g of CO2, 0.6369 g of H2O. Determine the empirical formula of the compound. (2016) (9) A well-known gas is enclosed in a container having volume 500 cm3 at S.T.P. its mass comes out to be 0.72 g. What is the molar mass of this gas? (2017) (10) When lime (CaCO3) is roasted, quick lime (CaO) is produced according to the following equation. The actual yield of CaO is 2.5 kg, when 4.5 kg of lime stone is roasted. What is the percentage yield of this reaction? (2017) CaCO3  CaO + CO2 (11) Ethylene glycol is used as automobile antifreeze. It has 38.7% carbon, 9.7% hydrogen and 51.6% oxygen. Determine its empirical formula. (2018) (2018)

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 02

“experimental techniques in chemistry” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 2 EXPERIMENTAL TECHNIQUES IN CHEMISTRY

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PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) (1) What are the major steps involved in crystallization? (2008) (2017) There are following steps which are involved in crystallization process. I. Choice of a solvent II. Preparation of saturated solution Muhammad Shahid III. Filtration A-One Professors Academy of IV. Cooling Sciences Jauharabad V. Collecting the crystals VI. Drying of the crystallized substance VII. Decolourization of undesirable colours (2) Give the main characteristics of the solvent used for crystallization. (2008) OR Write qualities of a good solvent. (2012) OR Give the salient features of an ideal solvent. (2013) An ideal solvent should have the following characteristics. i. The solvent should dissolve a large amount of the substance at its boiling point and only a small amount at room temperature. ii. It should not react chemically with the solute. iii. It should not dissolve the impurities. iv. It should be cheap. v. On cooling it should deposit well-formed crystals of the pure compound. vi. It should be safe to use and should be easily removable. (3) How in Gooch crucible the rate of filtration can be increases? (2009) The rate of filtration can be increased by placing Gooch crucible in a suction filtering apparatus. (4) Define sublimation with an example? (2012) OR What is sublimation? Name any two sublimates. (2009) (2014) The process in which a solid, when heated, vapourizes directly without passing through the liquid phase and these vapours can be condensed to form the solid again is called sublimation. For example Naphthalene, NH4Cl, benzoic acid, Iodine and camphor undergo sublimation. (5) Define filtration. How fluted filter paper is prepared from ordinary filter paper? (2010)

It can be defined as “The process of separation of insoluble particles from liquids by passing the mixture through a filter media is called filtration. Fluted filter paper can be prepared from ordinary filter paper by folding it in such a way that a fan like arrangement with alternate elevations and depressions at various folds is obtained.

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(6) In solvent extraction technique, why repeated extraction using small portions of solvent are more efficient than using a single extraction but larger volume of solvent? (2010) This is because more amount of solute is obtained by repeated extractions using small portions of volume of solvent than using a single extraction using larger volume of solvent. (7) What is solvent extraction? Muhammad Shahid OR A-one Professors Academy of sciences Jauharabad State solvent extraction and give its importance. (2011) OR Define solvent extraction and partition law. (2016) A technique in chemical analysis in which a solute can be separated from a solution by shaking the solution with another solvent in which the solute is more soluble and the added solvent does not mix with the solution is called solvent extraction. (8) Describe method to collect crystals from mother liquor. (2011) The crystals can be collected from mother liquor in the following way.  Mixture of mother liquor and crystals is filtered through Gooch crucible using a vacuum pump.  Full suction is applied to drain the mother liquor from the crystals.  When the filter cake is rigid enough it is pressed firmly with a cork to drain left over liquid.  The crystals are then washed with cold solvent and the process is repeated many times.  The mother liquor is concentrated by evaporation and cooled to obtain the fresh crop of crystals (9) How can HCl and KMnO4 solutions can be filtered by Gooch crucible? (2012) HCl and KMnO4 solutions can be filtered by using Gooch crucible if its perforations are covered with asbestos mat then these solutions will not react paper and can be filtered easily. (10) Define crystallization? (2012) The process of obtaining crystals of a substance by cooling its hot saturated solution is called crystallization. OR The process in which a dissolved solute comes out of solution and forms a crystalline solid is called crystallization. (11) Define distribution law or partition law? (2013) OR Define solvent extraction and partition law. (2016) OR State the law of distribution constant. (2017) This law states that a solute distributes itself between two immiscible liquids in a constant ratio of concentrations irrespective of the amount of solute added. (12) How vacuum desiccator is used to dry the crystals? (2013) (2015) (2017) (2018) It is safe and reliable method of drying the crystals. In this process the crystals are spread over a watch glass and kept in a vacuum desiccator for several hours. The drying agents used in desiccator are CaCl2, silica gel or phosphorus pentaoxide (P2O5).

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(13) Differentiate between adsorption chromatography and partition chromatography. (2013) (2014)  Chromatography, in which the stationary phase is solid, is called adsorption chromatography, while in which the stationary phase is a liquid is called partition chromatography.  In adsorption chromatography, a substance leaves the mobile phase to become adsorbed on the surface of solid phase, e.g. column chromatography, while in partition chromatography the substances being separated are distributed throughout both the stationary and mobile phases, e.g. paper chromatography (14) How the colouring impurities are removed from a crude substance? (2015) The colouring impurities are removed from a crude substance by boiling the substance in a solvent with the sufficient quantity of finely powdered animal charcoal and filtering the hot solution. The coloured impurities are adsorbed by animal charcoal and the pure decolourized substance crystallizes out from the filtrate on cooling. (15) How filter paper is folded? (2015) OR Which points should be kept in mind while folding of filter paper? The following points should be kept in mind while folding of filter paper:  The filter paper should be folded twice. The first fold should be along the diameter of the paper. The second fold should be such that edges do not quite match.  The paper should be opened on the slightly larger section in order to get a cone with three fold-thickness halfway around and one of thickness the other halfway around with an apex angle slightly greater than 600. (16) Name four important solvents chosen for crystallization. (2015) The solvents which are mostly used for crystallization are water, rectified spirit (95% Ethanol), absolute ethanol, chloroform, acetone etc. (17) How crystals are dried using filter paper? Give its two disadvantages. (2016) Crystals are dried by using filter paper. In this method the crystals are pressed between the several folds of filter paper and the process is repeated several times. Disadvantages:  Crystals can be crushed into fine powder.  Fiber of filter paper contaminate the product. (18) What are uses of chromatography? (2017) (2018) This technique is useful in organic synthesis for separation, isolation and purification of the products. They are equally important in qualitative and quantitative analysis. They are also used to determine the purity of substances. (19) What is difference between Gooch’s crucible and Sintered glass crucible? (2017) 1

Gooch’s crucible It is made of porcelain.

2

It has a porous base.

3

Its base is covered with filter paper or asbestos mat.

4

It is not convenient to use because preparation is needed.

Sr.No

Sintered glass crucible It is made of glass. It has a sintered glass disc sealed into its bottom. Filter paper or asbestos mat is not required. It is convenient to use because no preparation is needed as with Gooch’s crucible.

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(20) Why is there a need to crystallize the crude product? (2017) When a chemical compound is prepared then it is mostly a crude (impure) product. It cannot be used in crude form because impurities will reduce it efficiency. Therefore, there is a need to purify the crude product through the process of crystallization. (21)What is Rf value? (2018) Rf stands for retardation factor. Each component in the mixture has a specific retardation factor called Rf value the Rf value is related to the distribution coefficient and is given by: 𝐃𝐢𝐬𝐭𝐚𝐧𝐜𝐞 𝐭𝐫𝐚𝐯𝐞𝐥𝐥𝐞𝐝 𝐛𝐲 𝐚 𝐜𝐨𝐦𝐩𝐨𝐧𝐞𝐧𝐭 𝐟𝐫𝐨𝐦 𝐭𝐡𝐞 𝐨𝐫𝐢𝐠𝐢𝐧𝐚𝐥 𝐬𝐩𝐨𝐭 Rf = 𝐃𝐢𝐬𝐭𝐚𝐧𝐜𝐞 𝐭𝐫𝐚𝐯𝐞𝐥𝐥𝐞𝐝 𝐛𝐲 𝐬𝐨𝐥𝐯𝐞𝐧𝐭 𝐟𝐫𝐨𝐦 𝐭𝐡𝐞 𝐨𝐫𝐢𝐠𝐢𝐧𝐚𝐥 𝐬𝐩𝐨𝐭

(22)Differentiate between stationary phase and mobile phase. (2018) Sr.No

1 2

Stationary phase It may be solid or liquid supported on an inert solid. For example silica gel, water adsorbed in paper etc.

Mobile phase It may be liquid or gas. It flows over the stationary phase. For example acetone alcohol etc.

(23)A solid organic compound is soluble in water as well as in chloroform. During its preparation, it remains in aqueous layer. Describe a method to obtain it from this layer. The organic compound can be obtained by solvent extraction technique. Here organic compound is present in aqueous layer. Since organic compound is also soluble in chloroform, so this can be separated by shaking aqueous layer with chloroform in a separating funnel as chloroform is almost immiscible with water. The compound goes into the chloroform layer, which is separated from aqueous layer. Then chloroform is evaporated to get pure organic compound. (24) A water insoluble organic compound aspirin is prepared by the reaction of salicylic acid with a mixture of acetic acid and acetic anhydride. How will you separate the product from the reaction mixture? During preparation of aspirin, the reaction mixture is poured into the cold water to precipitate aspirin. Then it is filtered by a suitable filter media. The crude product is then crystallized from the mixed solvent i.e. equal volume of water and acetic acid. (25) What is qualitative analysis? The analysis which is concerned with the detection or identification of elements present in a compound is called qualitative analysis. (26) Define Quantitative analysis. The analysis which involves the determination of the relative amounts of the substances present in a sample is called quantitative analysis. (27) What are the four major steps in quantitative analysis? The four major steps in quantitative analysis are: I. Obtaining a sample for analysis II. Separation of desired constituents III. Measurement, and calculation of results IV. Drawing conclusion from the analysis (28) What is meant by filtration? The process of separation of insoluble particles from liquids by passing the mixture through a filter media is called filtration.

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PROFESSOR EASY NOTES

(29)What is the basic principle of crystallization? The basic principle of crystallization is the fact that the solute should be soluble in a suitable solvent at high temperature and the excess amount of the solute is given out as crystals when it is cooled. (30)What is Gooch crucible? It is made of porcelain having a perforated bottom which is covered with paper pulp. For quick filtration, Gooch crucible is placed in a suction filtering apparatus. (31)What is sintered glass crucible? Sintered glass crucible is a glass crucible with a porous glass disk sealed into the bottom. It is very convenient to use. (32)What is ether extraction? The ether extraction is used to separate the products of organic synthesis from water. The aqueous solution containing the organic product is shaken with ether in a separating funnel and allowed to separate the two layers. The ether layer containing the organic compound is separated and the organic product is obtained by evaporating the ether. (33)How crystals are dried by heating? Write its disadvantage. Crystals can be dried by heating them in an oven at 100 ºC. Disadvantages: Muhammad Shahid  Crystals can melt. A-one Professors Academy of  Crystals can decompose. sciences Jauharabad

Things to remember:

CHAPTER # 2 EXPERIMENTAL TECHNIQUES IN CHEMISTRY

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PROFESSOR EASY NOTES

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 3

“Gases” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) 1) Why high pressure and low temperature make gas non-ideal? (2007) (2010) OR Gases deviate from ideal behaviour more significantly at high pressure. Why? (2016) OR Why do real gases deviate from the ideal behavior at low temperature and high pressure? When the temperature of a gas is lowered, the average kinetic energy decreases and high pressure bring gas molecules close to each other, thus the attractive forces between the molecules become significant so real gases deviate from the ideality at low temperature and high pressure. 2) Lighter gases diffuse rapidly than heavier gases. Give reason. (2008) Although the average kinetic energies of different gases are same at the same temperature, but their molecular masses are different, so their velocities will also be different at the same temperature. The lighter gas molecules would have greater velocities so they diffuse rapidly than heavier gases. 3) What is plasma state? Where plasma is found? (2008) Plasma is often called the “Fourth state of matter”. It occurs only in lightning discharges and in artificial devices like fluorescent lights, neon signs, etc. The ionized gas mixture, consisting of ions, electron and neutral atoms is called plasma. Plasma is a distinct state of matter containing sufficient number of electrical charges to affect its behaviour, electrical and magnetic properties. Entire universe is almost of plasma. Plasma is found in everything from the sun to quarks. It is the stuff of stars. Our sun is a 1.5 million kilometer ball of plasma, heated by the nuclear fusion. On earth, it occurs in lightning bolts, flames, auroras and fluorescent lights. 4) Justify that the volume of a given gas becomes theoretically zero at -273 0C. (2008) (2018) The quantitative definition of Charles’s law is “At constant temperature, the volume of a given mass of a gas increases or decreases by 1/273 of its original volume at 00C for every 1 0C rise or fall in temperature”. According to this definition, if the volume of the gas at 0 0C is 273cm3. and temperature is decreased by 10C i.e. Temperature is -1 0C, then volume of the gas decreases and new volume becomes 272 cm3 as follows Student information: 𝑡 273 + 𝑡 V0 = volume at 0 °C Vt = V0 (1 + ) = V0 ( ) V-1 °C = 273 (

273 273−1 273

273 272

) = 273 (

273

Vt = volume at given temperature t = Given temperature

3

) = 273 × 0.9963369963 = 272 cm .

Now if the temperature decreases to -273 0C then new volume will become zero. 273 + 𝑡 Vt = V0 ( ) 273 Muhammad Shahid 273−273 V-273 °C = 273 ( ) = 273 (0) = 0 cm3. A-One Professors Academy of 273

PV = K

sciences Jauharabad The volume of a gas would become theoretically zero. 273 When T = -2730C; then V-2730C = V0 (1 )=0 273 Thus the volume of a given mass of a gas becomes theoretically zero at -2730C. 5) The plot of PV versus P is a straight line at constant temperature and with a fixed number of moles of an ideal gas. Justify it. OR The product of pressure and volume of a gas at constant temperature and number of moles is a constant quantity. Why? (2009) According to Boyle’s law, when the temperature and number of moles of a gas are constant, then the increase in pressure will decrease the volume in proportion to the increase in pressure, so that the product of pressure P

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

and volume remains constant (PV = k) By doubling the pressure the volume becomes half. Thus P1V1 = P2V2 = P3V3 = k (T and n are constant) 6) What is absolute Zero? What happens to real gas while approaching it? (2009) (2012) (2015) The hypothetical temperature at which the volume of a gas would become zero is called absolute zero, and is taken as the zero point on the Kelvin scale temperature (Zero Kelvin = -273.16 0C). This temperature, -273.16 is never achieved because all gases liquefy or solidify before reaching this temperature. Thus real gases will liquefy or solidify while approaching absolute zero. For routine calculations, the value of absolute zero is taken as -2730C. 7) At higher altitudes, the pilots feel uncomfortable breathings. Why? (2010) (2011) (2012) (2017) At higher altitudes, the pilots feel uncomfortable breathings because the partial pressure of oxygen in the unpressurized cabin is low as compare to 159 torr, where one feel comfortable breathing. 8) Define critical temperature and critical pressure. (2010) Critical temperature: “The highest temperature at which a substance can exist as a liquid, is called its critical temperature”. It is denoted by TC Critical pressure: The minimum pressure which is required to liquefy a gas at its critical temperature is called its critical pressure. It is denoted by PC. 9) Give two causes for deviation of gases from ideality. (2011) (2012) (2014) The deviation of real gases from ideality is due to the following two faulty assumptions. (1) The actual volume of gas molecules is negligible as compared to the volume of the vessel. (2) There is no forces of attraction among the molecules of a gas. However real gas molecules have finite volumes, and they attract one another. 10) State Dalton’s law of partial pressure. (2012) (2018) It can be defined as “The total pressure exerted by a mixture of non-reacting gases is equal to sum of their individual partial pressures”. Let we have three gases as 1, 2 , 3 and their partial pressures are p1, p2, p3. The total pressure (P) of the mixtures of gases is given by Pt = p 1 + p2 + p3 11) 280cm3 of H2 gas and 280 cm3 of CO2 gas at STP contain equal number of molecules. Why? (2012) (2018) According to Avogadro’s law “Equal volume of all the ideal gases at same temperature and pressure contain equal number of molecules”. Since both H2 and CO2 have equal volume at STP, so they will also have equal number of molecules. 12) Derive the units for general gas constant (R) in general gas equation, (a) When the pressure is in atmosphere and volume in dm3. When the pressure is in N/m2 and volume in m3. (c) When energy is expressed in ergs. OR Calculate the SI units of R? (2012) (2013) (2017) The unit of R can be calculated by Avogadro’s principle easily. Its value depends upon the units chosen for pressure volume and temperature. So there will be different units of R:  When the pressure is in atmosphere and volume in dm3  When the pressure is in N/m2 and volume in m3.  When energy is expressed in ergs.  When the pressure is in atmosphere and volume in dm3 𝑃𝑉 Since 𝑅 = 𝑛𝑇 𝑎𝑡𝑚× 𝑑𝑚3

𝑅= = dm3 atm K-1 mol-1 𝑚𝑜𝑙 × 𝐾  When the pressure is in N/m2 and volume in m3. 𝑃𝑉 Since 𝑅 = 𝑅=

𝑛𝑇 𝑁𝑚−2 × 𝑚3 𝑚𝑜𝑙𝑒 × 𝐾

= Nm K-1 mol-1 = J K-1 mol-1

(Since Nm = J)

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PROFESSOR EASY NOTES

 When energy is expressed in ergs. In CGS system “erg” is used to express the energy units. 1 J = 107 erg So unit of R will be R = erg K-1 mol-1 But remember that R will have different numerical value than SI unit i.e. R = 8.314 J K-1 mol-1 = 8.314 × 107 erg K-1 mol-1 13) What are isotherms? What happens to the positions of isotherms when they are plotted at high temperature for a particular gas? OR Why the graph plotted between pressure and volume moves away from pressure axis at higher temperature. (2013) (2018) Isotherms are the graphs plotted between pressure and volume when temperature and number of moles are constant. When the isotherms are plotted at high temperature, then they go away from the axis. The reason is that, the volumes of the gases increase at high temperature. When temperature is increased then volume has increased and new curve is obtained away from the axes. 1

T=250C T=00C

Volume (dm3)

Volume (dm3)

vα P Volume increases by changing T

Please follow Book Diagram,1st one is given to make the idea clear

T=250C T=00C

Pressure (atm) Pressure (atm) 14) Differentiate between diffusion and effusion of gases? (2013) The spontaneous intermixing of the molecules of different gases by random motion and collisions to form homogeneous mixture is called diffusion. It takes place in all directions. For example spreading of fragrance of rose. The escape of gas molecules one by one without collisions through a hole of molecular size into a region of low pressure is called effusion. It takes place form a hole. For example escape of gas molecule from punctured tyre. 15) Explain that the process of respiration obeys the Dalton’s law of partial pressure. (2013) (2015) The process of respiration depends upon the difference in partial pressures. When animals inhale air then oxygen moves into lungs as the partial pressure of oxygen in air is 159 torr. While the partial pressure of oxygen in lungs is 116 torr. CO2 produced during respiration moves out from the lungs into air, as the partial pressure of CO2 is more in the lungs than in air. Thus, process of respiration obeys the Dalton’s law of partial pressure. 16) What is Avogadro’s law of gases? (2014) This law states, “Equal volumes of all the ideal gases at the same temperature and pressure contain equal number of molecules”. Since equal number of moles of different ideal gases at the same temperature and pressure contain equal number of molecules, therefore the number of moles “n” of any ideal gas is directly proportional to its volume (V). V∝n

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

One mole of an ideal gas at 273.16K and 1atm pressure has a volume of 22.414 dm3. One mole of a gas has Avogadro’s number of particles, so 22.414 dm3 of various ideal gases will have Avogadro’s number of molecules, i.e. 6.02 × 1023. 17) How cooling is produced in joule Thomson effect? (2015) According to joule Thomson effect; “When a compressed gas is allowed to expand into a region of low pressure it gets cooled”. Actually the molecules of compressed gas are very close to each other and appreciable forces of attraction are present among them. When a gas is allowed to undergo sudden expansion through the nozzle jet, the gas molecules move apart. In this way energy is needed to overcome the intermolecular forces. This energy is taken from the gas itself, which is cooled. In this way cooling is produced in joule Thomson effect. 18) State Graham’s law of diffusion, give its mathematical expression. (2016) OR What is Graham’s law of diffusion? This law states that the rate of diffusion or effusion of a gas is inversely proportional to the square root of its density at constant temperature and pressure. 1 Rate of diffusion ∝ (at constant T and P) √𝑑 Muhammad Shahid 𝑘 Rate of diffusion = A-One Professors Academy of √𝑑

Or rate × √𝑑 = k 19) Derive Boyle’s law from kinetic gas equation. (2016) (2017) Kinetic equation for an ideal gas is: 1 PV= mN𝑐 2 ------ (1) 3

sciences Jauharabad

1

According to kinetic molecular theory of gases, the kinetic energy of N gas molecule ( mN𝑐 2 ) is 2 directly proportional to the absoulte temperature (T) 1 2 1

mN𝑐 2 ∝ T

mN𝑐 2 = KT ----- (2) Multiply and divide equation 1 by 2, we will get 2 1 PV= × mN𝑐 2 2

2

3

2

1

PV= × [ mN𝑐 2 ] ----- (3) 3 2 By putting equation 2 in equation 3. 2 PV= KT 3 If “T” = constant Then PV = K 𝐾 V= 𝑃 1

V∝ 𝑃 This is Boyles’s law. It shows that at constant temperature, the volume of a given mass of gas is inversly proprotional to the pressure exerted on it. 20) Water vapours don’t behave ideally at 273 K. (2016) Water vapours at 273 K have sufficient attractive forces because 273 K is the freezing point of water. Moreover 273 K is below than the critical temperature of water so Water vapours don’t behave ideally at 273 K. 21) SO2 is non-ideal at 273 K but behaves ideally at 327 K. (2017) SO2 is non-ideal at 273 K because this is low temperature and gas molecules have appreciable forces of attraction while 327 K is very high temperature, gas molecules have negligible forces of attraction hence SO2 behaves ideally at 327 K.

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

22) Write four applications of plasma. (2017) OR Write four applications of plasma. (2018) Plasma has following applications:  It is used to make corrosion resistant tools.  It can be used to pasteurize food.  It light up our offices, homes, helps in working of electronic equipments.  It drives laser and particle accelerators. 23) What is difference between centigrade and Fahrenheit scale and which relationship is used for their interconversion? (2017) Sr.NO

Centigrade scale Fahrenheit scale The distance between two points is The distance between two points is divided 1 divided into 100 equal parts. into 180 equal parts. On this scale the freezing point of On this scale the freezing point of water is water is 0 0C and boiling point is 100 2 32 0F and boiling point is 212 0F 0 C. The following relationships are used for the inter-conversion of centigrade and Fahrenheit scale. 5 Fahrenheit to Centigrade: 0C = (0F – 32) 9 9 0 (C 5

Centigrade to Fahrenheit: 0F = + 32) 24) Hydrogen and helium are ideal at room temperature, but SO2 and Cl2 are non-ideal, how do you explain it? (2018) H2 and He have very low boiling points which are far below the room temperature, so the room temperature is very high for H2 and He. Moreover, the molecules of H2 and He have very small masses and sizes. Due to very high temperature, small masses and sizes of H2 and He, the intermolecular forces are negligible and they behave ideally. While the boiling points of SO2 and Cl2 are close to room temperature and their molecular masses and sizes are relatively high. Thus significant attractive forces are present among their molecules, therefore SO2 and Cl2 are non-ideal at room temperature. 25) How does K.M.T explains Avogadro’s law? (2018) Consider two gases 1 & 2at the same P and having the same V. their number of molecules are N1 & N2, masses of molecules are m1 & m2 and mean square velocities are 𝑐12 & 𝑐12 respectively. Their kinetic equations can be written as follows: Kinetic equation for an ideal gas is: 1 PV= m1N1 𝑐12 for gas (1) 3 1

PV= m2N2 𝑐22 for gas (2) 3 By comparing equation 1 & 2. 1 1 m1N1 𝑐12 = m2N2 𝑐22 3

3

m1N1 𝑐12

= m2N2 𝑐22 ---------(3) When the temperature of both gases is the same, their mean kinetic energies per molecule will also be same, so 1 1 m1 𝑐12 = m2 𝑐22 2

2

m1 𝑐12

Or = m2 𝑐22 -------------- (4) Divide equation 3 by equation 4, we will get. N1 = N2 Hence equal volumes of all the gases at the same temperature and pressure contain equal number of moleculoes, which is Avogadro’s law.

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

26) Why the volume correction is done by van der Waal. The real gas molecules have definite volume. When high pressure is applied this volume is not negligible. So van der Waals did a correction in volume factor in order to make it applicable to the real gases. Thus, the volume available to gas molecules is the volume of the vessel minus the volume occupied by gas molecules. Vfree = Vvessel – b The factor b is called the excluded volume which is constant and characteristic of a gas. 27) The Van der Waal’s constant ‘b’ of a gas is four times the molar volume of gas. The gas molecules are incompressible spheres. When they are packed together like spheres, there remains some empty spaces in between the spheres which are also incompressible. Hence, the incompressible volume “b” is actually more than the actual volume of gas molecules. Actually it is four times than their molar volume. b = 4Vm 28) What is Kelvin scale of temperature? A scale of temperature at which -273.16oC is taken as zero Kelvin is called Kelvin scale temperature. Zero on the Kelvin scale corresponds to -273 0C. Therefore, K = 0 0C + 273.16oC. 29) What are various scales of thermometry? There are three scales of thermometry; i.e., centigrade, Fahrenheit and absolute or Kelvin scale. 5 9 0 0 C = (0F – 32), F = (0C + 32) and K = 0C + 273 9 5 30) How will you explain that the value of the constant K in PV=K (Boyle’s law) depends upon (a) Temperature of the gas. (b) Quantity of the gas  Temperature of the gas: According to Boyle’s law at constant temperature and number of moles the PV = K. When temperature is changes, the volume of gas is also changes i.e. increases or decreases. Therefore the product PV will also change. So we can conclude (‫ )ہجیتناکنانل‬that value of constant in PV = K will also change.  Quantity of the gas: According to Boyle’s law at constant temperature and the number of moles the PV = K. When number of moles (i.e. quantity of gas) are changes, the volume of gas is also changes i.e. increases or decreases. Therefore the product PV will also change. So we can conclude (‫ )ہجیتناکنانل‬that value of constant in PV = K will also change. 31) Why do we feel comfortable in expressing the densities of gases in the units of g/dm3 rather than g/cm3? In gases, Molecules are widely separated from each other. Small amount of gas is present in a large volume. Therefore values of densities will be much smaller if we expressed in cm3. So a bigger unit dm3 is used because appreciable amount of gas will be present per dm3 of gas and values of densities will be larger. So we feel comfortable in expressing the densities of gases in the units of g/dm3 rather than g/cm3. 32) Do you think that 1 mole of H2 and 1 mole of NH3 at 0 ºC and 1 atm pressure will have Avogadro’s number of particles? If not, why? One (1) mole of every substance has Avogadro’s number of particles. So 1 mole of H2 and 1 mole of NH3 at 0 ºC and 1 atm pressure will have Avogadro’s number of particles. 33) Justify that 1 cm3 of H2 and 1 cm3 of CH4 at STP will have same number of molecules, when one mole of CH4 is 8 times heavier than that of H2. In gases distance between two molecules is approximately 300 times their own diameter. So volume occupied by gas molecules does not depends upon the size or mass of molecules. The volume of gas only depends upon the number of molecules. Hence according to Avogadro’s law equal volume of H2 and CH4 at STP will have same number of molecules, although CH4 is 8 times heavier than that of H2.

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

P

34) Dalton’s law of partial pressure is only obeyed by those gases, which don’t have attractive forces among themselves, explain it. Dalton’s law of partial pressure is an ideal gas law which assumes that there are no attractive forces among the gas molecules. Hence every gas molecule moves independently. If gas molecules have attractive forces then every molecule will not move independently because when a molecule is about to colloid with the wall of container, it is attracted away from the wall by attractive forces of nearby molecules. As a result pressure exerted by the gas would be less than that of an ideal gas so Dalton’s law of partial pressure is only obeyed by those gases, which don’t have attractive forces among themselves. 35) How will you calculate the partial pressure of dry gas which is collected over water? When a gas is collected over water, it becomes moist. The pressure exerted by this moist gas is, therefore sum of the partial pressure of the dry gas and that of water vapours. Mathematically, Pmoist = Pdry + Pwater vapour Pmoist = Pdry + aqueous tension Pdry = Pmoist - aqueous tension The partial pressure exerted by the water vapours is called aqueous tension. 36) Why do we get a straight line when pressure exerted on a gas are High temeprature plotted against inverse of volumes? This straight line changes its T2 position in the graph by varying the temperature. 1 T1 When the pressure exerted on a gas is plotted against , we get a straight V line because pressure and inverse of volume are directly proportional to Low temperature each other at constant temperature. This straight line changes its position by varying the temperature, because this straight line at higher temperature will be close to zero; which means when P is very close to 1/v zero, then the volume is so high that 1/V is very close to zero. 37) What is Charles’s law? Which scale of temperature is used to verify V/T = k (P and n are constant)? Charles’s law gives the relationship between gas volume and temperature. This law states that the volume of a given mass of a gas is directly proportional to the absolute temperature when the pressure is kept constant. Mathematically V ∝ T or V=kT (when P and n are constant) 𝑉 =k (when P and n are constant) 𝑇 Thus, doubling the absolute temperature, causes the gas volume to double. If the temperature is changed from T1 to T2, the volume changes from V1 to V2, then Muhammad Shahid 𝑉1 𝑉2 =k and =k A-One Professors Academy of 𝑇1

Or

𝑉1 𝑇1

=

𝑉2 𝑇2

𝑇2

sciences Jauharabad

The value of the constant ‘k’ depends on the pressure and amount of the gas. Kelvin scale is used to verify that V/T = k 38) Do you think that the volume of any quantity of the gas become zero at -273 ºC. Is it not against the law of conservation of mass? How do you deduce the idea of absolute zero from this information? No, volume of a gas cannot be zero at -273 0C. Because mass of gas will be destroyed, which is against the law of conservation of mass (i.e. mass can neither be created nor destroyed). Since -273 0 C is unattainable lowest temperature. Therefore it is taken as absolute zero of Kelvin scale, so absolute zero is defined as “The hypothetical temperature at which the volume of all gases become zero is called absolute zero”. 39) Throw some light on the factor 1/273 in Charles’s law. This factor helped us in the development of Kelvin scale. The factor 1/273 in Charles’s law shows that at constant pressure the volume of a given mass of a gas increases or decreases by 1/273 of its original volume at 00C for every 10C rise or fall in temperature respectively.

CHAPTER # 3 GASES

24 1

PROFESSOR EASY NOTES 1

If we have 273 cm3 of a gas at 0 0C, then its part = × 273cm3 = 1 cm3 273 273 So for every 10C rise or fall in temperature, volume will increase or decrease by 1 cm3. The general equation to know the volumes of the gases at various temperature is; 𝑡 Vt = V0 (1 + ) 273 Where Vt is the volume of a gas at t 0C, V0 is the volume of the gas at 0 0C and t is the temperature on the centigrade scale. 40) The straight line in (a) is parallel to x-axis and goes away from the Pressure axis at higher pressure. At higher pressure the molecules of the gases come close to each other and attractive forces are created among gas molecules and gases do not remain ideal. Therefore, volume does not decrease in a regular way and hence the value of product PV increases and straight line goes away form the pressure axis. The nature of the curve depends upon the nature of gas. Deviation from ideal behaviour

PV

Ideal

Non-ideal

(Low=P)

(High = P)

PV = k

PV = k

Muhammad Shahid A-One Professors Academy of sciences Jauharabad

P

41) Pressure of NH3 gas at a given condition (say 20 atm pressure and room temperature) is less as calculated by Van der Waal’s equation than that calculated by general gas equation. The general gas equation is applicable for only ideal gases. Ideal gases have no forces of attraction and molecules hit the walls of container with greater force. Hence, pressure calculated by the general gas equation is always higher than calculated by van der Waals equation for non-ideal gases. NH3 is a polar gas and attractive forces exist among its molecules and it behaves non-ideally at room temperature and 20 atm pressure. These attractive forces decrease the force with which the molecules hit the wall. As a result, the pressure is less than that of an ideal gas. Therefore, pressure of NH3 gas is less as calculated by van der Waal's equation than that calculated by general gas equation. 42) Why gases do not settle? According to the kinetic molecular theory of gases, the gas molecules are in a constant motion, they collide with one another and their collisions are perfectly elastic therefore gases do not settle. 43) Briefly describe the significance of the constants ‘a’ and ‘b’ in the van der Waals equation. Significance of ‘a’ In van der Waals equation “a” is called co-efficient of attraction forces per unit volume. It indicates the strength of intermolecular forces in gases. Stronger the intermolecular forces, higher the value of “a” and vice versa. Unit of ‘a’ The factor ‘a’ is coefficient of attraction or attraction per unit volume of the gas its unit can be calculated form lessened pressure of the gas. P’ =

𝑛2 a

𝑉2 𝑃′𝑉 2

a= 2 𝑛 a = atm x (dm3)2/mol2 a = atm.dm6.mol-2 In SI units, pressure is in Nm-2 and volume is in m3. a = Nm-2 x (m3)2/mol2 a = Nm+4 mol-2

CHAPTER # 3 GASES

25

PROFESSOR EASY NOTES

Significance of ‘b’ In van der Waals equation “b” is known as effective volume occupied by gas molecules in a highly compressed state but not in liquid state. It depends upon the size of gas molecules. Greater the size of molecules greater will be the value of “b” and vice versa. Unit of ‘b’ The factor ‘b’ is excluded volume occupied by one mole of the gas. Hence, its unit should be 3 dm mol-1 or SI unit should be m3mol-1. 44) Gases deviate more from the general gas equation at 00C and deviate to less extent at 1000C. Why? At 00C the attractive forces between the gas molecules are significant thus gases become non-ideal. At 1000C the forces of attraction are negligible and they behave ideally. 45) What is meant by partial pressure of gas? The partial pressure of a gas in a mixture of gases is the pressure that it would exert on the walls of the container, if it were present all alone in the same volume under the same temperature. 46) How do you justify that from general gas equation that increase in temperature or decrease of pressure decreases the density of the gas? OR How will you calculate the density of an ideal gas from the general gas equation? The general gas equation for a gas can be written as PV = nRT 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑔𝑎𝑠 𝑚 Muhammad Shahid No. of moles of gas = n = = 𝑀𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑔𝑎𝑠

𝑀

A-One Professors Academy of sciences Jauharabad

Putting the value of n in the above equation gives 𝑚 PV = RT 𝑀 Rearranging the equation 𝑚 𝑚 PM = RT ∴ 𝑑= 𝑉 𝑉 PM = dRT 𝑃𝑀 Therefore, the density of the gas will be d = 𝑅𝑇 This formula shows that the density of gas is inversely proportional to the absolute temperature. If we increase the temperature then gas will expand and molecules will move apart and density will decrease. Similarly, formula shows that density is directly proportional to pressure. If we decrease the pressure then again molecules will move apart and density will decrease. 47) Can we determine the molecular mass of an unknown gas if we know the pressure, temperature and volume along with the mass of that gas? The general gas equation for a gas can be written as PV = nRT 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑔𝑎𝑠 𝑚 No. of moles of gas = n = = 𝑀𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑔𝑎𝑠

𝑀

Putting the value of n in the above equation gives 𝑚 PV= RT 𝑀 For the molecular mass of gas we can arrange this equation as 𝑅𝑇 M= 𝑚 𝑃𝑉 If we know the values of P, T, V and mass of gas then using this relation we can determine the molecular mass of an unknown gas. 48) How the density of an ideal gas doubles by doubling the pressure or decreasing the temperature on Kelvin scale by ½? We know that 𝑃𝑀 d= 𝑅𝑇

CHAPTER # 3 GASES

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PROFESSOR EASY NOTES

The density of an ideal gas is directly proportional to the pressure on the gas and is inversely proportional to the absolute temperature. Thus by doubling the pressure on the gas, the density becomes double. Similarly, when the temperature becomes one half, the density becomes double; 2𝑃𝑀 𝑃𝑀 2𝑃𝑀 d= ; d= 𝑇 = 𝑅𝑇

𝑅( ) 2

𝑅𝑇

Important Long questions of previous board papers 1. Calculate the mass of 11.2 dm3 of H2 at 0 0C and 2 atm pressure considering hydrogen to be ideal under these conditions. (2008) 2. What is plasma? Give three application of plasma. (2010) 3. Why real gases deviate from ideal behaviour? Discuss its causes. (2011) o 4. Calculate the density of methane at 0 C and 1.0 atm pressure. What will happen to density if (a) temperature is raised to 27 oC. (b) Pressure increases to 2 atm at 0 oC. (2012) 5. Derive van der Wall’s equation for real gases and give the physical significance of Vander Wall’s constants “a” and “b”. (2013) 6. What are applications of Dalton’s law of partial pressure? (2014) (2015) 7. Explain Boyle’s law and Avogadro’s law from kinetic molecular theory of gases. (2016) 8. What is kinetic molecular theory of gases? Write its postulates. (2015) (2016) 9. Derive general gas equation also calculate the value of “R” in S.I units. (2017) 10. What is meant by liquefaction of gases? Describe Linde’s method of liquefaction of gases. (2017) 11. Give explanation of applications of Dalton’s law of partial pressure of gases. (2018) 12. What is joule Thomson effect? Explain Linde’s method of liquefaction of gases. (2018) Things to remember:

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 4

“liquids and solids” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

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Important Short questions (Exercise is also added) 1. Explain London dispersion forces? (2007) (2009) (2010) (2015) “The momentary force of attraction created between instantaneous dipole and induced dipole is called dipole-induced dipole interaction or London dispersion forces”. These forces can exist between non-polar atoms or molecules like H2, Cl2 and noble gases (He, Ne etc.). When the electrons of one atom come close to the electron of other atom, they are pushed away from each other. In this way, a temporary dipole is created in one atom. This instantaneous (temporary) dipole can induce a similar dipole on an adjacent atom, causing the atoms to be attracted to each other. This interaction is called the London dispersion force. Muhammad Shahid 2. Justify that evaporation causes cooling. (2007) (2014) (2016) A-one Professors Academy of OR sciences Jauharabad Give the reason that evaporation causes cooling. (2017) When high energy molecules leave the surface of liquid, low energy molecules are left behind and average K.E of remaining molecules decreases so the temperature of the liquid falls. As the evaporation is a continuous process and liquid absorb more and more heat from surrounding, heat moves from the surrounding to the liquid and the temperature of the surrounding also falls. Hence evaporation causes cooling. 3. In a very cold winter the fish in the garden ponds owe their lives to hydrogen bonding. OR How sea life survives under frozen sea? (2008) Water shows strong hydrogen bonding due to more electronegative partial negative oxygen atom and partial positive small hydrogen atoms. When temperature falls below 4 0C then water expands and water molecules start arranging themselves in a regular pattern, the distance between them increases and water expands. As temperature falls to 0 0C, the water on the surface of pond or lake changes into ice and its volume increases almost 9%. Hence, density of solid ice is less than water and it floats on surface. Ice layer acts as an insulator and below the surface of this ice the temperature of water is 4 0C and fish can survive easily. 4. Why the melting and boiling points of alkanes increases with increase in molar mass? (2009) The melting and boiling points of alkanes increases with increase in molar mass because with increase in molar mass the number of atoms increases. Larger molecules have more places for attraction and greater polarizability than smaller molecules so larger molecules have stronger forces and higher melting and boiling points. 5. Why does ice float on water? Explain shortly. (2009) (2011) Water shows strong hydrogen bonding due to more electronegative partial negative oxygen atom and partial positive small hydrogen atoms. When temperature falls below 4 0C then water expands and water molecules start arranging themselves in a regular pattern and distance between them increases and water expands. As temperature falls to 0 0C, the water on the surface of pond or lake changes to ice and its volume is almost 9% more than liquid water. Thus, when water freezes, it occupies more space and density decreases. The result is that the ice has low density compared to that of water and ice floats on water. 6. Define molar heat of vaporization, fusion and sublimation? (2010) (in 2017 molar heat of vapourization )  The amount of heat absorbed by one mole of a solid when it melts into liquids at its melting point at one atmospheric pressure is called molar heat of fusion. It is denoted ∆Hf.  The amount of heat absorbed when one mole of a liquid is changed into vapours at its boiling point at one atmospheric pressure is called molar heat of vapourization. It is denoted ∆Hv.  The amount of heat absorbed when one mole of a solid sublimes to give one mole of vapours at a particular temperature and one atmospheric pressure is called molar heat of sublimation. It is denoted ∆Hs. 7. Food can be cooked quickly in pressure cooker. Explain with reason. (2011) (2014) Boling point of a liquid depends upon the external atmospheric pressure.

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b.p. ∝ external pressure External pressure can be increased artificially in pressure cooker. It is a closed container. The vapours of water formed are not allowed to escape. In this way, they develop more pressure in the cooker and the boiling temperature increases. As more heat is absorbed in water, so food is cooked quickly under increased pressure. 8. I2 is a solid but Cl2 is a gas at room temperature. Give reason. (2012) (2018) Iodine has larger size than Cl2. The larger molecules tend to have greater polarizability. In halogens the electronic cloud size increases down the group in the periodic table due to increase in atomic size. Hence they are easily polarized and develops strong London dispersion forces. Therefore I2 is a solid but Cl2 is a gas at room temperature. 9. Water and ethanol can mix easily in all proportions. (2012) Both water and ethanol (Ethyl alcohol) are highly miscible with each other because H2O and ethyl alcohol can form hydrogen bonding with each other, therefore they are miscible with each other in all proportions. 10. Why boiling point of H2O is greater than HF? (2013) In HF the F atom can make only one Hydrogen bond per molecule due to presence of only one hydrogen atom. While H2O can form two H-bond per molecule because it has two hydrogen atoms and two lone pair of electrons. Hence, due to presence of strong and extensive hydrogen bonding in H2O, the boiling point of water is greater than HF. 11. The boiling point of water is different at Murree hills and at Mount Everest. (2013) Boling point of a liquid depends upon the external atmospheric pressure. b.p. ∝ external pressure On Murree hills the external pressure is 700 torr which is lower than normal pressure, Hence water boils at 98 0C. While at Mount Everest the external atmospheric pressure is 323 torr. Hence water boils at 69 0C. So boiling point of water is different at Murree hills and at Mount Everest. 12. Give four uses of liquid crystals. (2013) (2016) Uses of liquid crystals: (1) Liquids crystals can be used as temperature sensors. (2) Liquid crystals are used in the display of electrical devices such as digital watches, calculators etc. (3) In chromatographic separations, liquid crystals are used as solvent. (4) Oscillographic and TV displays use liquid crystal screen. 13. Vacuum distillation can be used to avoid decomposition of a sensitive liquid. (2015) The distillation carried out under reduced pressure is known as vacuum distillation. It has many advantages because less fuel is required. The decomposition of many sensitive compounds can be avoided e.g. glycerin boils at 290 oC at 760 torr but decomposes at this temperature. Hence glycerin cannot be distilled at 290 oC. Under vacuum, the boiling temperature of glycerin decreases to 210 oC at 50 torr. It is distilled at this temperature without decomposition and can be purified easily so vacuum distillation can be used to avoid decomposition of a sensitive liquid. 14. Why evaporation take place at all temperatures? (2015) Molecules of liquids have K.E. at all temperature. They are continuously moving and colloid with one another. As a result of collision some molecules may get higher K.E. than others. Such molecules escape from the surface of liquid. Hence evaporation takes place at all temperatures. At low temperature molecules have low K.E. so the rate of evaporation is also low. However rate of evaporation increases with increase in temperature. 15. What are intramolecular forces of attractions? Give example. (2017) These can be defined as “The forces of attraction between atoms with in a molecule are called intramolecular forces of attraction”. Example: All type of chemical bonds i.e. ionic bond, covalent bond, coordinate covalent bond.

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16. The boiling point of a liquid remains constant although heat is continuously supplied to the liquid at its boiling point. Explain. OR Boiling needs constant supply of heat. Justify it. (2018) At the boiling point, the kinetic energy of the molecules becomes maximum and any further heating at this stage will not increase the temperature, rather the supplied heat will only be utilized to break the intermolecular forces and convert the liquid into its vapours. Therefore, the boiling point of a liquid remains constant although heat is continuously supplied to the liquid at its boiling point. 17. Why does the heat of sublimation of a substance is greater than that of heat of vapourization? In sublimation, solid directly changes into vapour phase, while in vapourization a liquid changes into vapour. The values of ∆Hs are greater than ∆Hv because attractive forces in solids are stronger than liquids. Therefore, heat of sublimation of a substance is greater than that of heat of vapourization. 18. Why heat of sublimation of iodine is very high? (2018) Iodine has the largest size among halogens and exists in solid state due to strong London dispersion forces. The London dispersion forces are created due to the large size of iodine molecules which are polarized by the neighbouring molecules. Hence, its heat of sublimation is very high. 19. HCl is a stronger acid than HF, why? (2018) HCl is a stronger acid than HF because it can ionize its hydrogen ion (H+) easily while HF is a weaker acid as compared to HCl due to strong hydrogen bonding in HF. Because the partially positive hydrogen is entrapped between two partially negative highly electronegative F atoms and cannot be ionized easily. F-δ F-δ F-δ +δ +δ +δ +δ +δ H H H H H -δ -δ -δ F F F 20. Why vapour pressure of CCl4 is 87 torr while isopentane is 580 torr at 20 oC. (2018) Vapour pressure is inversely proportional to intermolecular forces. 1 Vapour pressure ∝ Intermolecular forces Greater the strength of intermolecular forces lower will be the vapour pressure. CCl4 has stronger intermolecular forces as compare to isopentane so the vapour pressure of CCl4 is 87 torr while isopentane is 580 torr at 20 oC. 21. HF is a weaker acid, while HI is strong acid, explain it. HF is a weaker acid as compared to HI is due to strong hydrogen bonding in HF. Because the partially positive hydrogen is entrapped between two partially negative highly electronegative F atoms and cannot be ionized easily. F-δ F-δ F-δ Muhammad Shahid +δ +δ +δ +δ +δ H H H H H A-one Professors academy of F-δ F-δ F-δ sciences Jauharabad 22. Earthenware vessels keep water cool. Explain. Earthenware vessels are porous, when water is kept in earthenware vessels it evaporates through these pores. Due to continuous evaporation of water from the pores the high energy molecules escape and average K.E of remaining molecules decreases so temperature decreases. Therefore, earthenware vessel keeps water cool. 23. One feels sense of cooling under the fan after bath. Under fan the rate of evaporation from the surface of body is increased due to speed of air. When high energy molecules escape from the surface of the body by taking heat from body the temperature of the body falls. Hence one feels sense of cooling under the fan after bath. 24. Dynamic equilibrium is established during evaporation of a liquid in a closed vessel at constant temperature. Explain. When a liquid is allowed to evaporate inside a closed vessel, high-energy molecules leave the liquid and start gathering above the surface of the liquid. These molecules collide with the walls of the vessels as well as with the surface of the liquid. They are recaptured by the surface of liquid. This is called condensation. Two opposing processes, evaporation and condensation continue till a stage is

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reached when rate of evaporation becomes equal to the rate of condensation. The condition in which two opposing processes are occurring simultaneously at equal rates is called a dynamic equilibrium. 25. Why the melting point and boiling point of halogens increases down the group? In general larger molecules tend to have greater polarizability. In halogens the electronic cloud size increases down the group in the periodic table due to increase in atomic size. Hence they are easily polarized and develops strong London dispersion forces. Therefore melting and boiling points of the halogens increase with increasing molecular mass down the group. 26. What is meant by polarizability? Polarizability of a molecule is the quantitative measurement of the extent to which its electronic cloud can be distorted or polarized. Greater the number of atoms in a molecule, greater its polarizability. The greater the polarizability of a molecule, the more easily its electron cloud can be distorted to give a momentary dipole. 27. Water is a liquid at room temperature while H2S is a gas? H2O has strong hydrogen bonding. Therefore, H2O is a liquid because of hydrogen bonding, while H2S has no hydrogen bonding thus it is a gas at room temperature. 28. Define boiling point. Is it related with the external pressure? The temperature at which the vapour pressure of a liquid becomes equal to the atmospheric pressure is called the boiling point of the liquid. OR The temperature at which the vapour pressure of a liquid becomes to any other external pressure is called boiling point of the liquid. The boiling point is related with the external pressure. b.p. ∝ external pressure When the external pressure is changed, the boiling point of the liquid is also changed. A liquid can be made to boil at any temperature by changing the external pressure. The boiling point of the liquid is increased by increasing the external pressure and vice versa. 29. Steam can cause severe burns. Why? When water is heated then it boils at 100 0C. At the boiling point temperature remains constant and heat provided is used to change liquid water to steam (gas). The extra heat absorbed in steam is called latent heat of vapourization. Steam has 40.6 kJ/mol extra heat. When steam comes in contact with skin, it condenses, releasing considerable heat (latent heat), which causes severe burns. 30. How does the soaps and detergents perform the cleansing action? Soaps and detergents perform the cleansing action because the polar part of their molecules are water soluble due to hydrogen-bonding and the non-polar parts are alkyl or benzyl remains outside and are insoluble in water. 31. What is hydrogen bonding? Hydrogen bonding is the electrostatic force of attraction between a highly electronegative atom and partial positively charge hydrogen atom. The electronegative ion or atom ( usually on F,O, or N atom on another molecule). For example, a hydrogen bond exits between the H atom in an H-F molecule and the F atom of an adjacent HF molecule, where the dots represent the hydrogen bond between the molecules. Muhammad Shahid F-δ F-δ +δ +δ +δ +δ A-one Professors academy of H H H H sciences Jauharabad F-δ F-δ F-δ 32. a) What types of intermolecular forces will dominate in the following liquids. (i) Ammonia (NH3) (ii) Octane (C8H18) (iii) Argon (Ar) (iv) Propanone (CH3-CO-CH3) (v) Methanol (CH3-OH) (i) Why ammonia forms a single hydrogen bond? In ammonia, nitrogen is the more electronegative atom having one lone pair of electrons. Nitrogen is also covalently bonded to three hydrogen atoms.

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lone pair

H

H

N

H ------ N

H

H

PROFESSOR EASY NOTES Covalent Bonding

H

Hydrogen Bonding

Due to more electronegativity, the nitrogen atom attracts bonded pair of electrons towards itself and hence develops partial negative charge (-δ) and hydrogen atoms develop partial positive charge (+δ). Therefore, more electronegative atom attracts the partial positive hydrogen of the neighboring molecule forming Hydrogen Bonding. Ammonia molecule forms only one hydrogen bond per molecule as it has only one lone pair of electrons. (ii) Octane (C8H18) The octane is a non-polar hydrocarbon but due to large number of atoms the electronic cloud of the molecule is easily distorted by the neighbouring molecule. Hence, instantaneous dipole is created (molecule becomes polarized) and it cause London dispersion forces or instantaneous dipole induced dipole forces among the octane molecules. (iii) Argon (Ar) Argon is a noble gas and does not form bonds with any element and exist as monoatomic molecule. In the bulk of gaseous atoms, the electronic cloud of an Argon atom can distort the electronic cloud of neighbouring atom. This creates London dispersion forces among the Argon atoms. (iv) Propanone (CH3-CO-CH3) In propanone, the oxygen atom is more electronegative and hence develops partial negative charge (-δ) and carbon atom attached to oxygen develops partial positive charge (+δ). Due to these opposite ends (dipole), the molecules of propanone attract each other by dipole-dipole forces. (v) Methanol (CH3-OH) In methanol, the presence of –OH group creates hydrogen bonding among methanol molecules. (b) Propanone (CH3-CO-CH3), propanol (CH3-CH2-CH2-OH) and butane (CH3-CH2-CH2-CH3) have very similar relative molecular masses. List them in the expected order of increasing boiling points. Explain your answer. (i) Butane (CH3-CH2-CH2-CH3) Butane is non-polar hydrocarbon. But due to its large size the electronic cloud is polarized or distorted by the neighbouring molecules and hence, London dispersion forces are developed among butane molecules. Due to weak London dispersion forces, its boiling point is lower than propanone and propanol. (ii) Propanone (CH3-CO-CH3) In propanone, the carbonyl group is polar in nature and hence molecules of propanone develop dipoledipole forces. Due to dipole-dipole forces its boiling point is higher than butane but lower than propanol. (iii) Propanol (CH3-CH2-CH2-OH) In propanol, the presence of –OH group causes hydrogen bonding among the molecules of propanol. Therefore, its boiling point is higher than butane and propanone. The overall order of increasing boiling points of these compounds is as follow: Butane < Propanone < Propanol 33. Explain the following with reasons. (i) In the hydrogen bonded structure of HF, which is the stronger bond: the shorter covalent bond or the longer hydrogen bond between different molecules. When a covalent bond is formed between H and F atom and the bond energy released in very high. H + F  H-F ΔH = - 567 kJ/mol This shows that a very strong and shorter bond is formed between H and F atoms. The molecule of HF is polar in nature due to more electronegativity of fluorine atom. A partial negative charge develops on fluorine atom and a partial positive charge is created on hydrogen atom. The molecules of HF attract each other due to hydrogen bonding and molecules are arranged in a

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zig-zag manner. But the hydrogen bond is longer and weaker bond than shorter covalent bond of H and F atom. Moreover hydrogen bond is only 20% effective in strength as compared with a covalent bond. F-δ F-δ F-δ +δ +δ +δ +δ +δ H H H H H F-δ F-δ F-δ (ii) The origin of the intermolecular forces in water. In water molecule, the oxygen atom is more electronegative and covalently bonded to hydrogen atoms. The more electronegative atom of oxygen pulls the electron pairs towards itself and develops partial negative charge (-δ) and hydrogen atoms develop partial positive charges (+δ). There are two lone pairs of electrons on oxygen atom. Therefore, oxygen atom attracts the partial positive hydrogen atom of the neighbouring molecule and a strong hydrogen bond is formed as compared to simple dipole-dipole interaction. 29. (a) Briefly consider some of the effects on our lives if water has only a very weak hydrogen bonding present among its molecules? The density of ice as compared to water greatly affects our lives. If water has only very weak hydrogen bonding among its molecules, then ice would be denser than water. If ice were to be denser than water, ice forming at the top of a lake would sink to the bottom, and the whole lake could freeze to solid. Most aquatic life could not survive under these conditions. The expansion of water upon freezing causes water pipes to break in freezing weather. So, the pattern of our lives would have been totally different in the presence of very weak hydrogen bonding in water. (b) The critical temperatures of carbon dioxide and methane are 31.14 0C and – 81.9 0C, respectively. Which gas has the stronger intermolecular forces? Briefly explain your answer. The critical temperature is the highest temperature at which a gas can be liquefied by applying sufficient pressure called critical pressure. The gases which have appreciable forces of attraction among its molecules, usually their critical temperatures are very high. But gases which have negligible forces of attraction have very low critical temperatures. The reason is that at low temperature the kinetic energy of molecules decreases. When pressure is applied at the critical temperature then molecules come close to each other and they develop more forces of attraction and can be liquefied. CO2 is a non-polar gas but its molecule is larger in size than methane. Hence, its molecules have appreciable van der waal’s forces of attraction. Therefore, when sufficient pressure is applied then it can be easily liquefied at 31.14 0C. Methane is a non-polar gas but its molecules are smaller in size than CO2 and have negligible forces of attraction among its molecules at room temperature. When temperature is decreased to – 81.9 0C then molecules develop forces of attraction by applying pressure and hence can be liquefied at this lower temperature. 30. Three liquids have the properties mentioned against their names. Sr.No.

Properties

Water

Propanone

Pentane

1

Molecular Formula Relative Molecular mass (amu) Enthalpy Change of Vapourization (kJ/mol) Boiling Point (0C )

H2O 18 41.1 100

CH3-CO-CH3 58 31.9 56

CH3(CH2)3CH3 72 27.7 36

2 3 4

(a) What type of intermolecular forces predominate in each liquid? (i) Water (ii) Propanone (iii) Pentane  In water, the intermolecular forces are called hydrogen bonding.  In propanone, the intermolecular forces are called dipole-dipole interactions.  In pentane, the intermolecular forces are weak London dispersion forces.

Muhammad Shahid A-one Professors academy of sciences Jauharabad

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(b) What do you deduce about the relative strength of these forces in the liquids? Justify your conclusions. 1. In water, the hydrogen bonding is very strong type of intermolecular force. Due to strong hydrogen bonding, heat required to vapourize one mole of liquid water is 41.1 kJ called molar heat of vapourization at its boiling point i.e. 100 0C. 2. In propanone, the dipole-dipole interactions are not strong like hydrogen bond of water. The heat required to vapourize one mole of liquid propanone is 31.9 kJ called molar heat of vapourization at its boiling point i.e. 56 0C. The enthalpy change of vapourization of propanone is less than water and it proves that propanone has weaker intermolecular forces than water. 3. In pentane, the London dispersion forces are much weaker than dipole-dipole forces of propanone and hydrogen bonding of water. The heat required to vapourize one mole of liquid pentane is 27.7 kJ at its boiling point i.e., 360C. Therefore, the enthalpy change of vapourization of pentane is less than propanone and water. (c) If the liquids are shaken together in pairs, (i) Which pair would be unlikely to mix? Water is polar compound but pentane is non-polar. So these liquids will not mix up. (ii) Explain this immiscibility in terms of the forces between the molecules. There is a general principal ‘like dissolves like’. Water is polar and has hydrogen bonding among its molecules but pentane is non-polar and has very weak London dispersion forces among its molecules. Therefore, both will not mix up due to different nature of intermolecular forces. (iii) Choose one of the pairs that mix and say whether the enthalpy change on mixing would be positive or negative. Water and propanone are both polar in nature. So water can dissolve propanone due to formation of hydrogen bond. When some new attractive forces are created then energy is released. Therefore, enthalpy change on mixing of water and propanone would be negative.

Things to remember:

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SOLIDS

PROFESSOR EASY NOTES

Muhammad Shahid A-One Professors Academy of sciences Jauharabad

1) Why the metals are malleable and ductile? (2007) Malleable: The property of metal capable of forming sheets without being broken is called malleable. OR It is the property by virtue of which a metal can be rolled into sheets. Ductile: The property of metal capable of forming wires without being broken is called ductile. OR It is the property by virtue of which a metal can be drawn into wires. When a stress is applied on the surface of a metal, its layers slip pass each other, changing the structure of the metal without fracturing. Due to this reason metals are malleable and ductile. 2) Define crystal lattice and lattice points. (2008) Crystal lattice is an array of points representing atoms, ions or molecules of a crystal, arranged in different sites in 3-D space. In crystalline solids, these atoms, ions or molecules are located at definite positions in space. These points are called lattice points or lattice sites. 3) What is symmetry? Give elements of symmetry. (2009) The repetition of faces, angles or edges when a crystal is rotated by 360° along its axis is called symmetry. It is an important property of the crystal. There are various types of symmetry elements found in crystals. The more important symmetry elements are center of symmetry, axis of symmetry and plane of symmetry.

4) Ionic crystals are highly brittle. Why? (2010) (2017) (2017) In ionic crystalline solids, opposite charged ions are held together by strong electrostatic forces of attraction but in alternate fashion (+ve ion is surrounded by –ve ions and –ve ion by +ve ions). These ions occupy lattice points in three dimensional plane forming a strong crystalline lattice. When a force is applied then the layers of ions slip upon each other in result similar charged ions in layers face each other, then repulsive forces are created and ionic crystal is broken down easily. Therefore, ionic crystals are brittle and show the phenomenon of brittleness. Applied Force Repulsive forces

+ve ion -ve ion Alternate layers of opposite charged ions

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5) The electrical conductivity of metals decreases by increasing temperature. Explain with reason. (2010) The electrical conductance of metals is due to presence of free electrons. Sometimes, the electrical conductivity of metals decreases with the increase in temperature. The reason is that with the increase in temperature the positive metal ions also begin to oscillate (vibrate) and their motion hinders the free movement of mobile electrons. This hindrance decreases the electrical conductivity. 6) Differentiate b/w amorphous solids & crystalline solids? (2011)  A solid in which atoms, ions or molecules are arranged in a definite three dimensional pattern is called a crystalline solid. These solids have sharp melting points, for example diamond, SiO2, sucrose.  A solid whose atoms, ions or molecules are not regularly arranged or their regular shapes are destroyed are called an amorphous solids (amorphous means shapeless). These solids lack welldefined faces and shapes. They do not have sharp melting points, e.g. glass, rubber. 7) Define lattice energy with example. (2012) (2013) The energy released when one mole of the ionic compound is formed from the gaseous ions under standard conditions is called lattice energy. It can be defined as the energy required to break one mole of the solid into isolated ions in the gas phase. It is expressed in kJ mol-1. Na+ (g) + Cl- (g)  NaCl (s) ∆H = -787 kJ mol-1 + NaCl (s)  Na (g) + Cl (g) ∆H = +787 kJ mol-1 Lattice energy gives us some idea of the force of attraction between opposite ions in crystalline solid. 8) Discuss cubic crystal system. (2012) (2018) In this system all the three axes are of equal length and all are at right angle to a one another. a 9) Why Graphite is a good conductor along the layers. (2013) a Cubic In Graphite, the carbon atoms are arranged in layers so graphite has a layered structure. Each carbon atom is bonded to three other carbon atoms. In the layer it have one free electron. These electrons are delocalized along the layers which makes graphite a good conductor of electricity along the layers. While Graphite is not a good conductor perpendicular to the layers. 10) Differentiate between isomorphism and polymorphism. (2013) Muhammad Shahid OR A-one Professors Academy of Define isomorphism and polymorphism with examples. (2015) sciences Jauharabad OR Define polymorphism with example. (2016) (2018)  Polymorphism is a phenomenon in which a compound exists in more than one crystalline forms. The compound is called polymorphic and its different forms are called polymorph of each other. For example, AgNO3 exists as rhombohedral and orthorhombic crystalline forms. CaCO3 exists as Trigonal and orthorhombic crystalline forms.  Isomorphism is the phenomenon in which two different compounds exist in the same crystalline form. These different substances are called isomorph of each other. Mostly the ratio of atoms in various compounds are same and hence they show phenomenon of isomorphism. For example, Cu and Ag have same atomic ration i.e., 1:1 and exist in cubic crystalline form. 11) Sodium is softer than copper but both are very good electrical conductors. (2014) (2018) The forces in metallic solids depend upon number of valence electrons. Greater the number of electrons, stronger the forces. Sodium has lesser number of electrons in its valence shell than copper, therefore it has weak forces than copper. Due to weaker forces sodium is softer than copper. Both are equally very good conductor of electricity because both have free electrons. However, due to more free electrons in Cu, it is better conductor than sodium and used in electrical wires. 12) Ionic crystals do not conduct electricity in the solid state. Why? (2015) (2018) Ionic crystals or ionic solids do not conduct electricity in solid state because their ions have fixed positions due to strong electrostatic forces of attraction in a three dimensional plane. 90o

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However, in aqueous solution or in molten state, the ions are free to move hence ionic compounds are conductor of electricity in molten or aqueous state. 13) Explain why Iodine dissolves readily in tetra chloromethane (CCl4). (2015) (2017) (2018) There is a general principle that ‘like dissolves like’. It means non-polar solutes are soluble in nonpolar solvents and polar solutes are soluble in polar solvents. Iodine is a non-polar substance hence it readily dissolves in CCl4. It is more soluble in CCl4 (carbon tetrachloride) than in water. 14) Give four properties of covalent crystals. (2016) Properties of covalent crystals:  They have high melting and boiling points.  They have low volatility.  They have no free electrons. So bad conductor of electricity.  Their bond is directional in nature. 15) What is transition temperature? Give example. (2017) The temperature at which the two crystalline forms of the same substance can coexist in equilibrium with each other is called transition temperature. At this temperature one crystalline form of a substance changes to another. Above and below this temperature, only one form exists. 13.2 0C Grey tin (Cubic) White tin (Tetragonal) Muhammad Shahid 0 128 C A-one Professors Academy of KNO3 (Orthorhombic) KNO3 (Rhombohedral) sciences Jauharabad 16) Define allotropy? (2018) It can be defined as “The existence of an element in more than one crystalline forms is known as allotropy and these forms of the element are called allotropes or allotropic forms.” 13.2 0C Grey tin (Cubic) White tin (Tetragonal)gg 17) Justify that molecular crystals are softer than ionic crystals. In ionic crystals, opposite charged ions are held together by strong electrostatic forces of attraction. These ions occupy lattice points in three dimensional plane forming a strong crystal lattice. Hence, ionic solids are very hard e.g. NaCl has a melting point of 801oC. Molecular solids are composed of atoms or molecules held by weak intermolecular forces i.e. dipole-dipole forces or London dispersion forces. Due to weak intermolecular forces, molecular crystals are very soft as compared to ionic crystals e.g. solid ice melts at 00C. 18) Crystals of salts fracture easily, but metals are deformed under stress without fracturing. Explain the difference. In ionic crystalline solids, opposite charged ions are held together by strong electrostatic forces of attraction but in alternate fashion (+ve ions are surrounded by –ve ions and –ve ion by +ve ions). These ions occupy lattice points in three dimensional plane forming a strong crystal lattice. When a force is applied then the layers of ions slip over each other and similar charged ions in layers face each other and repulsive forces are created and crystal of salts fracture easily. Applied Force Repulsive forces

+ve ion -ve ion Alternate layers of opposite charged ions

Metals do not consist of oppositely charged ions. In metals strong metallic bond is present due to mobile electrons which hold positively charged metal ions together. When force is applied their layers slip over each other so the structure of the metal changes without fracturing. Therefore, metals are malleable and ductile.

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PROFESSOR EASY NOTES

19) Diamond is hard and an electrical insulator. In diamond, each carbon atom is covalently bonded to four other carbon atoms due to overlap of sp3-sp3 hybrid orbitals. This bonding gives rise to a tetrahedral structure with a mutual bond angle of 109.5o and bond length 154 pm. When many unit cells link with each other through covalent bonds in three dimensions a huge structure of crystal lattice is formed. Due to extensive bonding of carbon atoms in a network arrangement, the diamond is the hardest substance. It is an insulator due to no free electrons. 20) Sodium chloride and Cesium chloride have different structures. The structure of ionic crystals depends upon the radius ratio of cations and anions. If the radius ratio of ions in different compounds is same then their crystals will have same geometry and if radius ratio is different than their crystalline structure will be different. The ionic radius of sodium ion is 95 pm and chloride ion is 181 pm. The radius ratio of cation ion and anion will be 0.52. Sodium ion is surrounded by six chloride ions and its coordination number is 6. The ionic radius of cesium ion is 169 pm and chloride is 181pm. The radius ratio of cation and anion will be 0.93. Cesium ion is larger than sodium ion and is surrounded by 8 chloride ions and its coordination number is 8. 21) The vapour pressures of solids are far less than those of liquids. In solids the intermolecular forces are stronger than liquids, hence it is very difficult to vapourize solids. So vapour pressures of solids are far less than those of liquids. 22) Amorphous solid like glass is also called super cooled liquid. True solids have well ordered, three dimensional arrangement of atoms, molecules and ions in their crystal lattice. In liquids, the molecules are not arranged in ordered form and have random structure. The amorphous solid also have random arrangement of particles and their molecular structure is more like liquids. Hence, amorphous solid like glass is also called super cooled liquid e.g. glass rubber etc. 23) How polymorphism and allotropy are related to each other? Give examples. Polymorphism is a phenomenon in which a compound exists in more than one crystalline forms. The compound is called polymorphic and its different forms are called polymorph of each other. For example, 1. AgNO3 exists as rhombohedral and orthorhombic crystalline forms. 2. CaCO3 exists as Trigonal and orthorhombic crystalline forms. Allotropy is a phenomenon in which an element exists in more than one crystalline forms. These different forms are called allotropes or allotropic forms. For example, 1. Sulphur exists as rhombic and monoclinic crystalline forms. 2. Tin exists as grey tin (cubic) and white tin (tetragonal) All polymorphic forms are chemically identical and have same chemical properties but physical properties are different due to different structural arrangement of atoms, molecules and ions. Similarly, all allotropic forms of an element have same chemical properties but different physical properties due to structural arrangement of atoms. 24) Cleavage of the crystals is itself anisotropic behaviour. Some of the crystals show variation in physical properties depending upon the direction. Such properties are called anisotropic property and the phenomenon is called anisotropy. For example, refractive index, coefficient of thermal expansion, electrical and thermal conductivity, etc. Cleavage is the breakage of a crystal along definite planes. These planes are called cleavage planes. Since cleavage of a crystal can take place only in a particular direction, so it is an anisotropic behaviour. 25) The crystals showing isomorphism mostly have the same atomic ratio. Isomorphism is the phenomenon in which two different compounds exist in the same crystalline form. These different substances are called isomorph of each other. A crystalline form is independent of the chemical nature of the atoms and depends only on the number of atoms, sizes of atoms and their way of combination.

CHAPTER # 04

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PROFESSOR EASY NOTES

Mostly the ratio of atoms in various compounds are same and hence they show phenomenon of isomorphism. For example, Cu and Ag have same atomic ration i.e. 1:1 and exist in cubic crystalline form. 26) The transition temperature is shown by elements having allotropic forms and by compounds showing polymorphism. The different crystalline forms of an element are called allotropes and phenomenon is called allotropy but different crystalline forms of a compound are called polymorphs and phenomenon is called polymorphism. The temperature at which the two crystalline forms of the same substance can coexist in equilibrium with each other is called transition temperature. Therefore, one crystalline form of a substance can be changed to another form at transition temperature. Above and below this temperature, only one form exists. 13.2 oC Grey tin (Cubic) White tin (Tetragonal) 128 oC KNO3 (Orthorhombic) KNO3 (Rhombohedral) The concept of transition temperature cannot be used for substances which have only one form. Thus only elements with allotropic forms and compounds with polymorphic forms show transition temperature. 27) One of the unit cell angles of hexagonal crystal is 1200. In hexagonal crystal system, two axes are of equal length and are in one plane making an angle of 1200C with each other. The third axis which is different in length than other two is at right angle to these two axes and it is present in hexagon.

Muhammad Shahid

0 120

Hexagonal crystal

A-one Professors Academy of sciences Jauharabad

0

90

28) In the closest packing of atoms of metals, only 74% space is occupied. In metals, the metal atoms like hard balls (spheres) are present very close in the form of layers upon each other. This is called close packing of atoms in metals. Crevices or voids or interstices

In this arrangement, spaces are left among the metal atoms called crevices or void or interstices. Due to these spaces, volume occupied by atoms of metal is less than the total volume of metal. Thus it has been found that in closely packed structure only 74% of total volume is occupied by the metal atoms. Therefore, 26% space is left as voids. 29) The number of positive ions surrounding the negative ion in the ionic crystal lattice depends upon the sizes of the two ions. The structure of ionic crystals depends upon the radius ratio of cations and anions. If radius ratio of an ion is greater than it will be surrounded by greater number of ions and vice versa. Radius of cation Radius ratio = Radius of anion The number of ions surrounding the opposite charged ion in a crystal is called its coordination number. Coordination number depends on radius ratio of ions. If radius ratio is greater than coordination number will be greater and ion will be surrounded by larger number of ions. The ionic radius of sodium ion is 95 pm and chloride ion is 181 pm. The radius ratio of cation ion and anion will be 0.53. Hence, sodium ion is surrounded by six chloride ions and its coordination number is 6.

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PROFESSOR EASY NOTES

The ionic radius of cesium ion is 169 pm and chloride ion is 181pm. The radius ratio of cation and anion will be 0.93. Hence, cesium ion is larger than sodium ion and is surrounded by 8 chloride ions and its coordination number is 8. Therefore, the number of positive ions surrounding the negative ion in the ionic crystal lattice depends upon the sizes of the two ions. 30) State the biological applications of liquid crystals. Liquid crystalline substances are used to locate the veins, arteries, infections and tumors, since these parts of the body are warmer than the surrounding tissues. Specialists use the techniques of skin thermography to detect blockages in veins and arteries. When a layer of liquid crystal is pointed on the surface of the breast, a tumor shows up as a hot area, which is, coloured blue. This technique is useful in the early diagnosis of breast cancer. 31) What is meant by habit of a crystal? The shape of a crystal in which it usually grows is called habit of a crystal. Crystals are usually obtained by cooling the saturated solution or by slow cooling of the liquid substances. Crystals grow in various directions. If the conditions for growing a crystal are maintained then the shape of the crystal always remains the same. If the conditions for growing of crystal are changed, the shape of the crystal may change. For example, a cubic crystal of NaCl becomes needle like when 10% urea is present in its solutions as an impurity. 32) How a crystal system is identified. Write the names of crystal systems. A crystal system is identified by the dimensions of its unit cell along its three edges or axes, a, b, c and three angles between the axes α, β, and γ. These six parameters of the unit cell are called unit cell dimensions or crystallographic elements. The seven crystal systems are: cubic system, tetragonal system, orthorhombic system, monoclinic system, hexagonal system, trigonal system and triclinic system. 33) Why the ionic crystalline solids have high melting and boiling points? In ionic crystals, the cations and anions are held together by strong electrostatic forces of attraction. Very high energy is required to separate the cations and anions from each other. That is why ionic crystals have high melting and boiling points. 34) Why do most of the metals when freshly cut show metallic luster? The shining appearance of a metal, means its luster, is caused by the mobile electrons. Whenever metals are freshly cut they possess luster. When light falls on the metallic surface the incident light collides with the mobile electrons and they are excited. These electrons when de-excited give off some energy in the form of light. This light is reflected from the metal surface at all angles giving metal its peculiar luster.

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PROFESSOR EASY NOTES

Important Long questions of previous board papers 1. 2. 3. 4. 5.

Define ionic solids. Discuss properties of ionic solids in detail. (2007) What are covalent solids? Give their properties? (2008) Write four uses of liquid crystals. (2009) Briefly explain covalent solids, write their four uses. (2010) What is hydrogen bonding? Explain the role of hydrogen bonding in food and biological material. (2011) 6. Define vapour pressure of a liquid write Manometric method for determination of the pressure vapour pressure. (2012) 7. Write a note on factors affecting London forces. (2013) 8. Define vapour pressure. Write a method to measure vapour pressure of a liquid. (2013) (2014) 9. What are ionic solids? Give their properties. (2015) 10. Define vapour pressure. How vapour pressure is measured in Manometric method? (2016) 11. What is meant by the term hydrogen bonding? How does hydrogen bonding explain the properties of proteins? (2017) 12. What are molecular solids? Give their characteristics. (2017) 13. Explain the term molecular solids. Give three properties of molecular solids. (2018) 14. Briefly explain the properties of metallic solids. (2018)

Things to remember:

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 05

“atomic structure” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 5 STRUCTURE OF ATOMS

45

PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) (1) What particles are formed by the decay of free Neutrons? Write balance equations. (2007) 1 Free neutron decays into a proton ( 1P) with the emission of an electron ( -1oe) and a neutrino o ( on). 1

0n



1

1P

+

o

-1e

o

+ on

(2) What are X-rays? How they are produced? (2007) X-rays are electromagnetic radiations having high energy and frequency. These are produced when fast moving electrons collide with heavy metal anode in the discharge tube. (3) Whichever gas is used in the discharge tube, the nature of the cathode rays remains the same. Why? OR Why e/m value of cathode rays is same for all gases? (2007) OR Why e/m value of cathode rays is just equal to that of electron? (2009) (2017) The cathode rays are actually electrons. Since electrons are present in all atoms and their nature is same. Thus, nature of cathode rays remains same, no matter which gas is used in the discharge tube. Therefore, these are considered as fundamental particle of atom. It was proved experimentally by J.J. Thomson. He calculated e/m ratio of cathode rays by taking different gases in the discharge tube. But he always found the same e/m ratio. It shows that cathode rays obtained by different gases have same nature. The e/m ratio of cathode rays is always equal to that of electron because cathode rays are basically electrons. (4) Describe the shape of s and p orbital. (2007) (2014) Shapes of s orbitals: s-orbital has a spherical shape and is represented by a circle which in turn represents a cut of sphere. With the increase in value of principle quantum number (n), the size of s-orbital increases as shown in the diagram.

Muhammad Shahid Node

1s

2s

A-One Professors Academy of sciences Jauharabad

3s

Shapes of p orbitals: There are three values of magnetic quantum number for p-subshell. So p-subshell has three orientations in space i.e. along x, y and z-axes. z z z

y px

y py

y pz

(5) Define atomic orbital what is about the probability of finding electron between two orbitals? (2008) It can be defined as “The volume of space in which there is 95% chances of finding electron is called atomic orbital”. The probability of finding electron between two orbitals is zero.

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PROFESSOR EASY NOTES

(6) Justify that the distance gaps between different orbits go on increasing from the lower to the higher orbits. (2009) According to Bohr’s theory the equation of orbit is r = 0.529 × n2. This equation shows that radius of the orbit is directly proportional to the square of orbit number (n), so higher orbits have more radii and vice versa hence radius of orbits goes on increasing with increasing the orbit numbers. (7) What is the origin of line spectrum? (2009) According to Bohr, the electron could “jump” from one allowed energy state to another by absorbing, or emitting photons of radiant energy of certain specific frequencies. When an element, such as hydrogen gas is heated, its electron “jumps” from lower energy state (E1) to higher energy state (E2) by absorbing energy. When it comes back, the same energy is released. This difference in energy (E2 – E1) is emitted as a radiation of a definite frequency in the form of spectral line. ∆E = E2 – E1 = h𝜈 The spectral lines of Lyman series are produced when the electron jumps from n2 = 2, 3, 4, 5 ……. to n1 = 1 Similarly, spectral lines of Balmer series originated when electron jumps from n2 = 3, 4, 5, 6 …… to n1 = 2 orbit. Paschen, Brackett and Pfund series are produced as a result of electronic transition from higher orbits to 3rd, 4th and 5th orbit respectively. (8) Why is it necessary to decrease the pressure in the discharge tube to get the cathode rays? (2010) (2016) (2017) The current does not flow through the gas at ordinary pressure even at high voltage about 5000 volts. However, when the pressure inside the tube is decreased and a high voltage of 5000 to 10000 volts is applied, then an electric discharge takes place through the gas producing a uniform glow inside the tube. When pressure is reduced to 0.01 torr the original glow disappears, some rays are produced which creates a greenish glow. These rays are called cathode rays. Therefore, it is necessary to decreases the pressure in the discharge tube to get the cathode rays. (9) What are defects in Rutherford atomic model? (2010) (2018) I. Rutherford atomic model is based on the law of motion and gravitation which can be applied to the neutral bodies like planets but not to the charged bodies like protons and electrons. II. Electron being a charged a body, will emit energy continuously and the radius of the orbiting electrons should become smaller and smaller and the electron should fall into the nucleus and finally the atom would collapse. Since this does not happen, the Rutherford planet-like structure was defective and unsatisfactory. III. If the electron, radiate energy continuously, the atom should give continuous spectra, but actually atoms form line spectra. (10) Why positive rays are also called canal rays? (2010) Since positive rays produced in the discharge tube passed through the canals or holes of the cathode, therefore, positive rays are also called canal rays. (11) State Moseley’s law and write down it’s equation. (2010) (2013) OR What is Moseley’s law? Moseley’s law states the frequency of a spectral line in X-ray spectrum varies as the square of atomic number of an element emitting it. This law convinces us that it is the atomic number of the element, which determines its characteristic properties, both physical and chemical. The relationship between the frequency ‘v’ of a particular of X-rays and the atomic number Z of the element emitting it is given below √𝜈 = a(Z-b)

CHAPTER # 5 STRUCTURE OF ATOMS

47

PROFESSOR EASY NOTES

Here ‘a’ and ‘b’ are constant characteristic of the metal under consideration. This equation is known as Moseley’s law ‘a’ is proportionality constant and ‘b’ is called screening constant of the metal. (12) Differentiate between continuous spectrum and line spectrum. (2011) The spectrum in which colours diffuse into each other without any dark space and the boundary line between the colours cannot be marked is called continuous spectrum. It is the characteristic of matter in bulk. The familiar example of continuous spectrum is rainbow. The spectrum in which different colour are not diffused into each other and coloured lines are separated by dark spaces, is called line spectrum. It is characteristic of an atom. Line spectrum of Na contains two yellow coloured lines separated by a definite distance. (13) What is the significance of Azimuthal quantum number? (2008) OR Define Azimuthal quantum number and give its importance. (2011) OR Define and explain Azimuthal quantum number. (2015) The quantum number which gives information about subshells i.e. their shape is called azimuthal quantum number. It is denoted by “l”. Its value depends upon the value of “n”. For each value of n, it can have any integer value from zero to n-l. Its value for a particular subshell is designated from 0, 1, 2 and 3 for s, p, d and f subshell respectively. Total number of electrons in a given subshell can be calculated by following formula. 2(2Ɩ + 1) (14) Electron has its dual nature. Justify. (2011) OR What is de-Broglie concept of duality of matter? (2015) According to de-Broglie all matter particles in motion have a dual character. It means that electrons, protons, neutrons atoms and molecules possess the characteristics of both the material particle and a wave. This is called wave-particle duality in matter. He gave a mathematical relation which relates the wavelength (λ) of the electron to the momentum of electron. h Muhammad Shahid λ= mv A-One Professors Academy of (15) Mention two defects of Bohr’s atomic model. (2011) (2014) (2015) sciences Jauharabad I. It cannot explain the fine structure of hydrogen atom. II. It cannot explain the Zeeman effect and Stark effect. (16) Define Hund’s rule. Give one example. (2012) (2017) (2018) It can be defined as “If degenerate orbitals are available and more than one electrons are to be placed in them, they should be placed with the same spin rather than putting them in the same orbital with opposite spins. (17) What is Lyman series? In which region it lies? (2012) When the spectrum of hydrogen is observed by a high resolving spectrometer it shows different spectral lines. These spectral lines grouped into five groups called spectral series. Layman series is one of them. It arises when excited electron come back to first shell (n1) form any higher shell (n2) where n2 may be any higher shell i.e. 2, 3, 4 … This series lies in ultraviolet (U.V) region. (18) Write down the electronic configuration of 24Cr and 29Cu. (2012) The electronic configuration of 24Cr is given following. 1

1

1

1

1

1s2, 2s2, 2p6, 3s2, 3p6 4s1, 3d5 OR [Ar] 4s1, 3dxy, 3d yz, 3d xz, 3d x2-y2, 3d z2 The electronic configuration of 29Cu is given following. 2

2

2

2

2

1s2, 2s2, 2p6, 3s2, 3p6 4s1, 3d10 OR [Ar] 4s1, 3dxy, 3dyz, 3d xz, 3d x2-y2, 3d z2 (19) What is Zeeman Effect? (2012) (2013) (2017) (2017) (2018) When the excited H-atoms are placed in a magnetic field, its spectral lines are further splits up into closely spaced lines. This type of splitting of spectral lines is called Zeeman Effect.

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(20) State Heisenberg’s uncertainty principle with mathematical expression. (2012) It is impossible to measure simultaneously the exact position and momentum of electron. h

(21)

(22)

(23)

(24)

∆x.∆P ≥ where ∆x is the position of electron and ∆P is the momentum of electron. 4π What is the function of Principal quantum number? (2013) OR What do you understand by the term principal quantum number? The principal quantum number determines the distance of electron from the nucleus, greater the value of ‘n’ greater will be the distance of electron from the nucleus. It also determines the energy of an electron in an atom.. It can have any positive integral value: 1, 2, 3 and so on. Why 4s sub-shell is filled first and 3d afterward (2013) The subshells are filled by n + l rule. According to this rule, subshells are arranged in the increasing order of n + l values. If any two subshells have same n + l values, then that subshell is placed first which have smaller value of “n”. For example 4s subshell is filled first and 3d latter because n + l value of 4s is smaller than that of 3d. For 4s n = 4, l = 0 n+l=4+0=4 For 3d n = 3, l = 2 n+l=3+2=5 Why potential energy of an electron is negative in an orbit of atom? (2013) According to Bohr’s theory, at infinity the electron is not being attracted by anything so P.E is zero. But at a point near to the nucleus, electron is attracted by the nucleus so P.E is less than zero i.e. negative. Why the electron move faster in an orbit of smaller radius? (2013) In the derivation of radius of revolving electron in nth orbit we have following relation

𝑟=

(25)

(26)

(27)

(28)

𝑍𝑒 2 4𝜋𝜀˳𝑚𝑣 2

In this relation radius (r) is inversely proportional to square of velocity (v). 1 𝑟 ∝ 2. So it becomes clear that electron will move faster in an orbit of smaller radius. 𝑣 What is atomic emission spectrum? (2014) (2017) Question were asked to state differences OR b/w atomic emission and absorption spectrum in 2017 annual paper. How atomic emission spectrum is obtained? When solids are volatilized or elements in their gaseous state are heated to high temperature or subjected to an electrical discharge, radiation of certain wavelengths are emitted. The spectrum of this radiation contained bright line against a dark background. This is called atomic emission spectrum. The lines appear bright because the corresponding wavelengths are being emitted by the element. What is atomic absorption spectrum? (2017) When a beam of white light is passed through a gaseous sample of an element, the element absorbs certain wavelengths which appear dark, while the rest of wavelengths pass through it which appear bright. The spectrum of this radiation is called atomic absorption spectrum. Write down four properties of cathode rays. OR Write down any three properties of cathode rays. (2015)  Cathode rays are negatively charged.  They cast a shadow when an opaque object is placed in their path.  Cathode rays can ionize gases.  Cathode rays produce heat when they fall on matter. Calculate the wave number of photon, when it jumps from n = 5 to n = 2. (2015) The wave number of photon can be calculated by using following equation. 1 1 Wave number (𝜈) = 1.09678 × 107 [𝑛2 − 𝑛2 ] 1

2

CHAPTER # 5 STRUCTURE OF ATOMS

49

Here n1 = 2 and n2 = 5 1 Wave number (𝜈) = 1.09678 × 107 [22 −

1 52

PROFESSOR EASY NOTES

]

Wave number (𝜈) = 23.00 × 105 m-1 (29) Define Auf-bau Principal and Pauli Exclusion Principle. (2016) (2017) (2018 Pauli Exclusion Principle) According to Auf-bau Principal Electrons should be filled in energy subshells in order of increasing energy values. The electrons are first placed in 1s, 2s, 2p and so on. While Pauli Exclusion Principle states that it is impossible for two electrons residing in the same orbital of a poly-electron atom to have the same values of four quantum numbers. OR Two electrons in the same orbital should have opposite spins (↑↓). (30) Give two postulates of plank’s quantum theory. (2016) (2017) OR What are the main points of quantum theory of radiation? The main points of this theory are:  Energy is not emitted or absorbed continuously, rather it is emitted or absorbed discontinuously in the form of wave packets of energy called quanta.  The amount of energy associated with a quantum of radiation is proportional to the frequency (𝜈 ) of the radiation. Frequency is the number of waves passing through a point per second. E∝𝜈 E = ℎ𝜈 (31) The bending of cathode rays in the electric and magnetic fields shows that they are negatively charged. In 1885, J.Perrin showed that cathode rays are negatively charged. He found that when cathode rays are passed through a magnetic field, these are curved downward by the magnetic field. Moreover, in 1897, J.J. Thomson showed that when cathode rays are passed through an electric field, these are deflected towards positive plate. These experiments show that cathode rays are negatively charged. (32) The e/m value for positive rays obtained from hydrogen gas is 1836 times less than that of cathode rays. (2018) Positive rays obtained from hydrogen gas in a discharge tube consists of protons and the cathode rays consists of electrons. A proton and an electron have equal magnitude of charge but mass of e a proton is 1836 times greater than that of an electron. Hence, value of positive ray obtained m from hydrogen gas is 1836 times less than that of cathode rays. (33) The e/m values of positive rays for different gases are different, but those for cathode rays the e/m values are same. Justify it. (2018) Positive rays are produced by the ionization of gas present in the discharge tube. The characteristic of the gas varies from gas to gas. Therefore, the e/m values of positive rays depend upon the nature of gas used in the discharge tube. Greater the number of protons and e neutrons in nucleus of an atom or molecule, smaller will be the values. While cathode rays m always consist of a beam of electrons, so e/m values remain the same. (34) Define frequency and wave number. (2018) The number of waves passing through a point per second is called frequency. Its units are cycle’s s-1 or Hz. Wave number is the number of waves per unit length, and is reciprocal of wavelength. 𝐞 (35) Calculate the mass of an electron, = 1.7588 × 1011 coulombs/kg. 𝐦

e

The value of charge on electron is 1.6012 × 10-19 coulombs. While of electron is 1.7588 × 1011 m coulombs/kg. So

CHAPTER # 5 STRUCTURE OF ATOMS e m

=

1.602 ×10−19 Mass of electron

50

PROFESSOR EASY NOTES

= 1.7588 × 1011 coulombs/kg

Rearranging the above equation. 1.602 ×10−19

Mass of electron = 1.7588 ×1011 Mass of electron = 9.109 × 10-31 kg. (36) Why does the size of He+ is much smaller than H-atom although both H-atom and He+ ion are mono-electronic systems. (2018) Both hydrogen atom and He+ have one electron in their outermost shell. However the nucleus of He+ has greater positive charge due to two protons than that of hydrogen atom due to one proton. Therefore the nucleus of He+ attracts its electron more strongly than hydrogen. Hence the size of He+ is much smaller than H-atom although both H-atom and He+ ion are mono-electronic systems. (37) What is Stark effect? (2018) When the excited atoms of hydrogen are placed in an electric field, its spectral lines are further split up into closely spaced lines. This type of splitting is called Stark effect. (38) What are positive rays or canal rays? Positively charged ions formed by the ionization of the gas molecules with the passage of cathode rays in a gas discharge tube are called positive rays or canal rays. These rays travel from anode toward cathode in a discharge tube. Muhammad Shahid (39) Total energy of bounded electron is negative. Why? A-One Professors Academy of According to Bohr’s theory the total energy of electron is En = −

(40)

(41)

(42)

(43)

(44)

mZ2 e4

sciences Jauharabad

4πε˳2 n2 h2

The negative sign indicates that the energy decreases when electron is brought from infinity to a distance “r”. At infinity the electron is not being attracted by anything so P.E is zero. But at point near to the nucleus, electron is attracted by the nucleus so total energy is less than zero i.e. negative. Justify, that cathode rays are material particles. Cathode rays drive a small paddle wheel which shows that these rays possess momentum. From this observation, it is inferred that cathode rays are not rays but particles having a definite mass and velocity. Therefore, cathode rays are material particles. Why the anode rays depend upon the nature of gas? When we determine the value of e/m of anode rays for different gases, we find that the e/m values for different gases are different due to different masses of positively charged ions, which depend upon the identity of the residual gas in the discharge tube. Hence the anode rays depend upon the nature of the gas. How positive rays are produced? The positive rays are produced, when high speed cathode rays (electrons) collide with the gas molecules enclosed in the discharge tube. They knock out the other electrons, producing positive ions, which start moving towards the cathode. M + e-  M + 2eWrite four properties of positive rays. (1) They are deflected by an electric as well as a magnetic field, showing that they are positively charged. (2) These rays travel in a straight line in a direction opposite to the cathode rays. (3) They produce flashes on ZnS plate. (4) The e/m value for the rays is always smaller than that of electron and depends upon the nature of the gas used in the discharge tube. How neutron was discovered? In 1932, Chadwick bombarded the nucleus of beryllium (94Be) with α-particles produced from a polonium source and noticed that some highly penetrating radiation were produced. These

CHAPTER # 5 STRUCTURE OF ATOMS

51

PROFESSOR EASY NOTES

radiation were called neutrons because charge detector showed them to be neutral. The nuclear reaction is as follows: 9

4Be

4

+ 2He

12



6C

1

+ 0n

Actually α-particles and the nuclei of Be are rearranged and extra neutron is emitted. (45) How is slow neutron prove to be more effective than the fast neutrons? When neutron travel with an energy 1.2 Mev (Mega electron volt), they are called fast neutrons while which have energy below 1ev are called slow neutrons. Slow neutrons are usually more effective than fast one in fission process. When slow moving 66 neutron hit the Cu metal then γ-radiations are emitted, giving a radioactive 29Cu, which is 66 converted into 30Zn by emitting an electron (β-particle). 65

1

66

Cu + hv ( γ - raditions)

Cu + 0n 29

29

66

66

Cu 29

0

Muhammad Shahid A-One Professors Academy of sciences Jauharabad

Zn + -1e 30

(46) How did Rutherford’s model of an atom provide the existence of the nucleus of the atom? Rutherford observed the effect of bombarding thin gold foil with α-particles and found that almost all the α-particles passed directly through the foil without any deflection, a few was deflected and some were reflected back on its original path. He postulated that most of the mass of the atom and all of its positive charge reside in a very small, extremely dense region, which he called nucleus. Most of the total volume of the atom is empty space in which electrons move round the nucleus. (47) Define wavelength. Wavelength is the distance between two adjacent crests or troughs and is expressed in Ao, nm or pm. (48) What are the main postulates of Bohr’s model of atom? The main postulates of Bohr’s theory are: (1) Electron revolves around the nucleus in one of the circular orbits. Each has fixed energy and a quantum number is assigned to it. (2) As long as an electron remains in a particular orbit, it does not emit or absorb energy, the energy is emitted or absorbed only when an electron jumps from one orbit to another. (3) When an electron jumps, the energy change (∆E) is given by the equation. ∆E = E2 – E1 = h𝜈 Where ∆E is the energy difference between the two orbits with energies E1 and E2. Energy is absorbed by the electron when it jumps from a lower orbit to a higher orbit and vice versa. (49) What is spectrum? A visual display of the dispersion of the components of white light, when it is passed through a prism, is called spectrum. (50) The angular momentum of moving electron is quantized. Justify it. According to Bohr, electron can revolve only in those orbits having a fixed angular momentum (mvr). The angular momentum of an orbit depends upon its quantum number and it is an

integral multiple of the factor mvr =

nh 2π

Angular momentum mvr = n For n = 1

ℎ 2𝜋 ℎ 2𝜋

. where n = 1, 2, 3, 4 ……..etc.

angular momentum of 1st orbit = 1× nd

For n = 2

angular momentum of 2 orbit = 2×

For n = 3

angular momentum of 3rd orbit = 3×

ℎ 2𝜋 ℎ 2𝜋 ℎ 2𝜋

CHAPTER # 5 STRUCTURE OF ATOMS

52

PROFESSOR EASY NOTES

The angular momentum of 2nd orbit is twice than the angular momentum of 1st orbit. Similarly the angular momentum of an orbit for n = 3 is thrice than the angular momentum for n = 1, electron is always present in these orbits, not in between hence the angular momentum is quantized. (51) What are quantum numbers? Quantum number are the sets of numerical values which give the acceptable solutions to Schrodinger wave equation for hydrogen atom. An electron in an atom is completely described by its four quantum numbers. These are: (1) Principal quantum number (n) (2) Azimuthal quantum number (3) Magnetic quantum number (m) and (4) Spin quantum number (s). (52) How will you differentiate between orbit and orbital? Orbit: A define circular path at a definite distance from the nucleus in which the electron revolves is called an orbit. An orbit indicates an exact position or location of an electron in an atom. It represents the planar motion of the electron. Orbital: It can be defined as “The volume of space in which there is 95% chances of finding electron is called atomic orbital”. The probability of finding electron between two orbitals is zero. An orbital does not specify the exact position of an electron in an atom due to the wave nature of an electron. It represents the three dimensional motion of the electron around the nucleus.

Important Long questions of previous board papers I. II. III. IV. V. VI. VII. VIII. IX. X. XI. XII. XIII. XIV. XV. XVI. XVII.

What are quantum numbers? Discuss their significance. (2007) Give the main points of Rutherford’s atomic model and also describe two defects of this model. (2008) How e/m value of electron was measured? (2009) Describe atomic emission and atomic absorption spectrum of with diagram. (2010) Write defects of Rutherford’s atomic model. How Bohr removed them? (2011) Write a note Planck’s Quantum theory. (2012) Derive radius of revolving electron in nth orbit of an atom. (2012) (2013) Write electronic configuration of elements with atomic number z = 24, z = 37. (2012) (2013) Discuss (i) Azimuthal Quantum number (ii) Magnetic quantum number (2013) (2017) What are quantum numbers? Explain Principal and Azimuthal quantum number. (2014) (2015) Write electronic configuration of elements with atomic number z = 29. (2015) Explain Heisenberg’s uncertainty principle. (2015) Write four properties of cathode rays. (2016) What are quantum numbers? Explain principle quantum number. (2017) What are quantum numbers? Explain Azimuthal quantum number and magnetic quantum number. ` (2017) Mention four defects of Bohr’s atomic model. (2018) Derive an expression for radius of nth orbit of Hydrogen atom with the help of Bohr’s atomic model. (2018)

Things to remember:

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 06

“chemical bonding” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 6 CHEMICAL BONDING

55

PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) 1) Define dipole moment. Give its various units. (2007) (2013) (2018) OR Define dipole moment. Give it’s mathematically expression. (2008) (2013) (2015) It can be defined as “The product of the electric charge (q) and the distance between the positive and negative centers (r)”. µ=q×r Units: SI unit of dipole moment is mC (meter coulomb). For a unit negative charge at a distance of 100 pm from a unit positive charge we have µ = (1.6022 × 10-19) × (100 ×10-12 m) = 1.6022 × 10-29 mC. Muhammad Shahid mC is a bigger unit. The other unit of dipole unit Debye. A-One Professors Academy of 1D = 3.336 × 10-30 mC sciences Jauharabad

1.6022 ×10−29

µ= = 4.8 D 3.366 ×10−30 2) Define bond order. Calculate the bond order of nitrogen molecule. (2007) OR Define bond order. Give an example. (2017) The number of bonds formed between two atoms after the atomic orbitals overlap, is called the bond order. It is taken half of the difference between the number of bonding electrons and anti-bonding electrons. Bond Order =

Number of electron in bonding MO orbitasl − Number of electron in anti−bonding MO 2

The bond order of nitrogen molecule can calculated as Number of electrons in bonding molecular orbitals Number of electrons in anti-bonding molecular orbitals 8−2 6 Bond Order = = =3 2 2 So nitrogen can make three bonds. N

N2

N

σ* (2px)

π* (2py) π* (2pz) 2py

A.O

2px

2px

σ (2px)

2py A.O

2pz

Energy

A.B.M.O

2pz

=8 =2

It has been made to make idea clear about bonding and anti-bonding molecular orbitals. Student can make it in final exam to ensure their full marks. Thanks

B.M.O

π(2py) π(2pz) A.B.M.O

σ* (2s) 2s A.O

2s A.O

σ (2s) B.M.O

3) The radius of anion is always larger than its parent atom. Why? (2007) (2008) (2017) The ionic radii of negative ions are always larger than the size of their parent atoms because number of electron increases but the nuclear charge on the nucleus remains same, moreover the addition of one or more electrons in the shell of a neutral atom increases the repulsion b/w the electrons causing expansion

CHAPTER # 6 CHEMICAL BONDING

56

PROFESSOR EASY NOTES

of the shell. Therefore hold of nucleus on outer electrons decreases resulting the increase in the ionic radius e.g. radius of F atom is 72 pm while that of F- ion is 136 pm. 4) Why the energy of anti-bonding molecular orbital is higher than corresponding bonding molecular orbital? (2007) In anti-bonding molecular orbital, the electron density is concentrated outside the region between two nuclei. Thus, an electron in an anti-bonding molecular orbital is actually repelled from the bonding region and have higher energy and is less stable. While in bonding molecular orbital the electron density is concentrated between two nuclei. Thus an electron in bonding molecular orbital is strongly attracted by both nuclei having lower energy and are more stable. Hence energy of A.B.M.O is greater than B.M.O. 5) Sigma bonds are stronger than pi (π) bond. Why? (2008) (2011) (2012) (2013) (2015) (2018) Sigma bond is formed by linear overlap of atomic orbitals while pi (π) bond is formed by parallel overlap of atomic orbitals. Greater the area of overlapping stronger will be the bond. The overlapping area in sigma bond is more than pi (π) bond. Hence Sigma bonds are stronger than pi (π) bond. 6) Ionization energy increases along the period of periodic table. Why? (2008) In the period of periodic table the outer shell remains the same, atomic size decreases and effective nuclear charge increases which makes the removal of an electron difficult hence ionization energy increases in the period of periodic table. 7) Why the first electron affinity for most of the element is negative while the second electron affinity for all the element is positive. (2009) The first electron affinity for most of the element is negative because 1st electron is added in neutral atom. When second electron is added in a uni-negative ion, the incoming electron is repelled by the already present negative charge and energy is absorbed therefore the first electron affinity for most of the element is negative while the second electron affinity for all the element is positive e.g. O + e-  O- E.A = -141 kJ/mol O- + e-  O2- E.A = +780 kJ/mol 8) Why atomic radius increase down the group? (2009) (2016) The increase in atomic radii from top to bottom in a group is due to increase in the number of shells and the screening effect. Thus electrostatic forces of attraction between nucleus and valence shell electrons decreases in result atomic radius increases. 9) How MOT justifies that Helium atoms cannot make He2? (2009) (2013) (2015) The electronic configuration of Helium is 1s2. The 1s orbitals of He-atoms combine to form two molecular orbitals. The two electrons enter in B.M.O and remaining two enter in A.B.M.O. The bond order for He2 is zero. Thus He2 molecule does not form. The bond order of helium molecule can calculated as Number of electrons in bonding molecular orbitals =2 Number of electrons in anti-bonding molecular orbitals =2 2−2 0 Bond Order = = =0 2

He2

He

A.B.M.O

σ* (1s) 1s A.O

Muhammad Shahid Energy

He

2

A-One Professors Academy of sciences Jauharabad

1s A.O

σ (1s) B.M.O

10) Bond distance is the compromise distance between atoms. Justify it. (2010) When two atoms come close to each other, both attractive and repulsive forces are developed between them simultaneously. However, at the time of bond formation, attractive forces dominate the repulsive forces and two atoms are bonded to each other having minimum energy. Thus it is called

CHAPTER # 6 CHEMICAL BONDING

57

PROFESSOR EASY NOTES

the bond length or bond distance or compromise distance of two atoms. So Bond distance is the compromise distance between atoms. 11) Why atomic radius cannot be measured accurately? (2010) The atomic radius cannot be measured accurately because  There is no sharp boundary of an atom.  The electronic probability distribution is affected by nearby atoms. 12) Define ionization energy. How does it vary in the periodic table? (2010) Ionization energy can be defined as “The minimum amount of energy which is required to remove an electron from the valence shell of an isolated gaseous atom to form a positive ion is called ionization energy” e.g. Na  Na+ + eI.E = 496 kJ/mol Trend in Group: Ionization energy decrease down the group because with increase in size, there is more shielding, hence the force of attraction between nucleus and valence shell decreases, so removal of an electron becomes easy with less use of energy. Trend in period: In the period of periodic table the outer shell remains the same while effective nuclear charge increases which makes the removal of an electron difficult hence ionization energy increases in the period. 13) Define coordinate covalent bond and give its two example. (2010) (2013) It can be defined as “A coordinate covalent bond is formed between two atoms when the shared pair of electrons is donated by one of the bonded atom”. The atom which donates the electron is called “donor” and the atom which accepts the electron is called an acceptor”. For example the formation of ammonium ion (NH4+) and Hydronium ion. + H H H

x

H

x

N

+ H

+

O

x

H

x

N

x

x

H

H

H

H

H

+

H

+

O H

+

H

Hydronium ion

, 14) Define polar covalent bond with example. (2011) A covalent bond between two different atoms in which the shared electron pair is more attracted towards the more electronegative atom and the bonded atoms acquire partial positive and negative charge is called polar covalent bond. For example bond in HCl (Hδ+- Clδ-) HF H2O etc. 15) Why MOT is superior to VBT? (2011) (2013) (2015) (2016) Molecular orbital theory (MOT) explains the no bond formation between noble gases. It also accounts the observed paramagnetic behaviour of O2. The excited states in molecules can easily be explained by MOT. While Valence bond theory (VBT) does not give such answers thus MOT is superior to VBT. 16) Differentiate between bonding molecular orbital and anti-bonding molecular orbital? (2012) Sr.No

1

2 3 4

Bonding molecular orbital In bonding molecular orbital the electron density is concentrated between two nuclei. An electron in bonding molecular orbital is strongly attracted by both nuclei having lower energy. Bonding molecular orbital is more stable. For example σ and π bond.

Anti-bonding molecular orbital In anti-bonding molecular orbital, the electron density is concentrated outside the region between two nuclei. An electron in an anti-bonding molecular orbital is actually repelled from the bonding region. Anti-bonding molecular orbital is less stable. For example σ* and π* bond.

CHAPTER # 6 CHEMICAL BONDING

58

PROFESSOR EASY NOTES

17) Dipole moment of CO2 is zero but that of H2O is 1.85D. Why? (2012) (2017) CO2 has linear structure. In this structure the individual bond moment cancel the effect of each other therefore dipole moment becomes zero. While H2O has angular or bent structure in which individual bond moments do not cancel the effect of each other. Hence it has dipole moment 1.85D. CO2 µ = 0 δO

δ+

C

H2O µ = 1.85 D δO

δO

δ+

Muhammad Shahid

δ+

H

H

Resultant Dipole Moment

A-One Professors Academy of sciences Jauharabad

18) Define bond length and bond energy. (2012) Bond length: The distance between the nuclei of two atoms forming a covalent bond is called bond length. For example the bond length of C-C bond in aliphatic hydrocarbons is 154 pm. Bond energy: The average energy required to break one mole of similar bonds in a substance is called bond energy. Or it can also be defined as “Bond energy is a measure of strength of a bond”. 19) Why the ionization energies values decrease down the group although nuclear charge increases. (2012)

The ionization energy decrease in a group. Because when we move down the group, the atomic size increases and more shells are added between nucleus and valance shell electrons. These additional shells decreases the electrostatic force felt by electrons present in the valance shell. As a result the removal of electron becomes easy so ionization energy decreases down the group although nuclear charge increases. 20) The bond angles of H2O and NH3 are not 109.5o like that of CH4. Although O and N-atoms are sp3 hybridized. (2012) (2014) CH4, H2O and NH3 undergo sp3 hybridization. All should have bond angle 109.5o. In CH4, there is no lone pair of electron so it have regular tetrahedral structure and bond angle is 109.5o. In NH3, there is one lone pair of electron which exert repulsion on bond pair thus bond angle is reduced 109.5o to 107.5o. In H2O, there are two lone pair of electron which exert more repulsions on bond pairs thus bond angle is reduced 109.5o to 104.5o. Hence the bond angles of H2O and NH3 are not 109.5o like that of CH4. 21) The bond energy of multiple bonds are higher than single bond. Why? (2012) The multiple bonds are stronger than single bond. For example the triple is stronger than double bond which is stronger than single bond. C

C

> C

C

>

C

C

Greater the number of bonds greater will be the bond energy. Therefore the bond energy of multiple bonds are higher than single bond. 22) Explain that ionization energy is the index of metallic character. (2013) Metals have low ionization energies because the electron of metals are loosely held due to their large size. Thus, metals can lose electron easily when they undergo chemical reaction. Elements having high values of ionization energy are non-metallic in nature. So ionization energy is the index of metallic character. 23) Why the lone pair of electrons on an atom occupy more space? (2013) (2016) A bonding pair of electron is under the influence of two nuclei of bonding atoms. It is attracted by both nuclei of atoms. While lone pair electrons are under the influence of only one atom. Because a lone pair experiences less nuclear attraction, its electronic charge is spread out more in space than for bonding pair. Therefore the lone pair of electrons on an atom occupy more space. 24) Write the Lewis structure of CS2 and CCl4? (2014) The Lewis structure of CS2 and CCl4 is given following.

CHAPTER # 6 CHEMICAL BONDING

59

PROFESSOR EASY NOTES

Cl x

Cl x

C

x

Cl

S xx C xx S

x

Cl

25) Why repulsive forces are less between bond pairs than lone pair? (2014) A bonding pair of electron is under the influence of two nuclei of bonding atoms. It is attracted by both nuclei of atoms. While lone pair electrons are under the influence of only one atom. Because a lone pair experiences less nuclear attraction, its electronic charge is spread out more in space than for bonding pair. Therefore repulsive forces are less between bond pairs than lone pair. 26) Why the atomic radius decrease along period and increase along group? (2015) The atomic size decreases in a period because i. With increase in atomic number, the effective nuclear charge increases gradually because of addition of more and more protons in the nucleus. ii. On the other hand addition of electron take place in the same valance shell i.e. shell number do not increases. While atomic radius increases along group 27) Define ionic bond. Give one example. (2015) It can be defined as “The bond formed by the complete transfer of one or more electrons form an atom with low ionization energy to another atom with high electron affinity is called ionic bond”. For example the formation of NaCl. Na  Na+ + eCl + e-  ClNa+ + Cl-  NaCl 28) What factors influence ionization energy? (2015) Ionization energy depends upon the following factors.  Atomic size  Shielding effect  Effective nuclear charge  Nature of orbital. 29) What is paramagnetic character? Give the reason for paramagnetic character of oxygen. (2015) The substances which are weakly attracted by strong magnetic field are called paramagnetic substances and this phenomenon is known as paramagnetic character. This is arises due to presence of unpaired electrons. Oxygen shows paramagnetic character because it has two unpaired electrons. O

O2

O

Unpaired electron causing paramagnetic behaviour

σ* (2px)

π* (2py) π* (2pz) 2pz

2py

A.O

2px

2px

π(2py) π(2pz)

2py A.O

B.M.O

σ (2px)

Muhammad Shahid A.B.M.O

σ (2s) *

2s A.O

2s A.O

σ (2s) B.M.O

2pz

Energy

A.B.M.O

A-One Professors Academy of sciences Jauharabad

CHAPTER # 6 CHEMICAL BONDING

60

PROFESSOR EASY NOTES

30) How the % ionic character of the polar bond can be determined? (2016) In order to determine the %age of ionic character of polar bond, we should know the actual dipole moment µobs of the molecule and actual bond length. μ %age of ionic character = obs × 100 μ ionic

µionic is obtained by multiplying the bond length with the charge of electron i.e. µionic = q × r 31) Cationic radius is smaller than parent atom. Why? (2017) The cationic radius is smaller than parent atom because the removal of one or more electron from neutral atom usually result in loss of outermost shell, it disturbs the proton-electron ratio, moreover the effective nuclear charge increases, which reduces the size of orbit by pulling the electron inwards. Hence the size decreases. 32) What is octet rule? Give one example. (2017) OR Explain the octet rule with one example. (2017) Octet rule: It can be defined as “The tendency of atoms to attain a maximum of eight electrons in the valence shell or outer most shell is known as the octet rule”. Having 8 electrons in the valance shell is a sign of stability. For example noble gases. 33) Why NH3 can form co-ordinate covalent bond with H+ but CH4 not? (2017) A coordinate covalent bond is formed between two atoms when the shared pair of electrons is donated by one of the bonded atom. In NH3, nitrogen has a lone pair of electron so it can make bond with H+ by donating pair of electron. While in CH4 there is no lone pair of electrons so it cannot make coordinate covalent bond with H+. 34) Draw the geometry of SO2 and SO3 on the basis of VSEPR theory. (2017) According to VSEPR theory SO2 and SO3 are AB3 type molecules with multiple bonding. In SO2, one corner is occupied by a lone pair and two corners each by S=O double bond. It has bent or angular structure. While in SO3, each corner is occupied by S=O double bond. It has perfectly triangular structure. O S S O, O O O 35) How electron affinity changes in a group? (2018) Electron affinity decreases in a group due to increase in number of shells and shielding effect. Because due to successive increase of electronic shells atomic radii increases in a group. This exerts a shielding effect on the force of attraction between nucleus and the valence electrons. 36) How a co-ordinate covalent bond differs from a covalent bond? (2018) Sr.No

1

2 3 4

Coordinate covalent bond A coordinate covalent bond is formed between two atoms when the shared pair of electrons is donated by one of the bonded atom.

There should be at least one lone pair of electron. One of the bonded atom is known as “Donor” and other is known “Acceptor”. For example of formation of ammonium ion etc.

Covalent bond

A covalent bond is formed by mutual sharing of electrons between two atoms. There is no need of presence of lone pair of electron. No idea of donor or acceptor. For example formation of Cl2, F2 etc.

CHAPTER # 6 CHEMICAL BONDING

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PROFESSOR EASY NOTES

37) Reactions between ionic compounds are very rapid. Give reason. (2018) The ionic compounds exist in the form of ions in an aqueous solution. The chemical reactions between ions occur rapidly. Because on mixing two solutions, no bond have to be broken, only a new bond is formed. The ionic compound have already been broken while forming their aqueous solution. For example, addition of silver nitrate solution to sodium chloride solution produces a white precipitate of silver chloride instantaneously. 38) Bond angle in water is 104.5o instead of 109.5o. give reason (2018) In H2O, there are two lone pair of electrons which exert more repulsions on bond O + + pairs thus bond angle is reduced 109.5o to 104.5o. Therefore Bond angle in water is 104.5 o o 104.5 instead of 109.5 . H water H 39) π bond is more diffused than sigma (σ) bonds. Give reason. (2018) Sigma (σ) bond is formed by head to head overlap of atomic orbitals of atoms and the electron density is present between the nuclei. While π bond is formed by sideways overlap of atomic orbitals and electron density is present above and below the line joining of two nuclei. Hence π bond is more diffused than sigma (σ) bond. 40) Define ionic radii and covalent radii. (2018) Atomic radius is defined as “The average distance between the nucleus and outermost orbit of an atom, while considering the atom spherical e.g. Na has atomic radius 157 pm. While ionic radii is defined as “The radius of the ion while considering it spherical” e.g. Na+ ion has ionic radii 95 pm. 41) In many cases, the distinction between a co-ordinate covalent bond and a covalent bond vanishes after bond formation in NH4+, H3O+ and CH3NH3+. Explain with reason. In many cases, the distinction between a co-ordinate covalent bond and a covalent bond vanishes after bond formation in NH4+, H3O+ and CH3NH3+ because in NH4+, H3O+ and CH3NH3+ ions, it has been experimentally found that all bonds are equal. Thus there is no distinction between a co-ordinate covalent bond and a covalent bond. 42) No bond in chemistry is 100% ionic. Justify it. The difference of electronegativity between the bonded atoms help us to decide the %age of ionic nature in a compound. The maximum difference of E.N between two atoms for CsF is 3.2. Calculations tell us that CsF has 92% ionic character and NaCl has 72% ionic character. No bond in chemistry has ionic character more than 92%. Hence no bond in chemistry is 100% ionic. 43) The species NH2-, NH3 NH4+ have bond angles of 105o, 107.5o, and 109.5o respectively. Justify these values by drawing their structures. o

NH2-

Sr.No -

1

NH3 3

In NH2 nitrogen is sp hybridized, it forms two covalent bonds with hydrogen atoms. There are two lone pairs on nitrogen atom. Hence according to VSEPR theory, it will have a bent structure. Since two lone pairs exerts more repulsions on bond pairs. Hence the bond angle will be reduced to 109.5o to 105o.

In NH3 nitrogen is sp3 hybridized, it forms three covalent bonds with hydrogen atoms. There is one lone pair on nitrogen atom. Hence according to VSEPR theory, it will have a triangular pyramidal structure. Since lone pair exerts more repulsions on bond pairs. Hence the bond angle will be reduced to 109.5o to 107.5o.

NH4+

In NH3+ nitrogen is sp3 hybridized, it forms three covalent bonds and one coordinate covalent bond with hydrogen atoms. There is NO lone pair on nitrogen atom. Hence according to VSEPR theory, it will have a regular tetrahedral structure. Hence the bond angle will be 109.5o.

CHAPTER # 6 CHEMICAL BONDING

62

PROFESSOR EASY NOTES

+

-

N

N

2

109.5 o

N 107.5 o

105

o

44) The abnormality of bond length and bond strength in HI is less prominent than in HCl. Give reason. Chlorine has higher electronegativity than iodine. The bond length and bond energy depend upon the electronegativity difference of bonded atoms. The E.N difference in HCl is higher than HI. Thus HCl has higher polarity and greater attractive forces for binding the atoms than HI. Therefore the abnormality of bond length and bond strength in HI is less prominent than in HCl 45) The dipole moment of CO2 and CS2 are zero, but that of SO2 is 1.61D. Give reason. CO2 and CS2 have linear structure. In these structures the individual bond moment cancel the effect of each other. Therefore dipole moment becomes zero. While SO2 has angular or bent structure in which individual bond moments do not cancel the effect of each other. Hence it has dipole moment 1.61D. CO2 µ = 0

CS2 µ = 0

SO2 µ = 1.61 D δ+

δ-

δ+

δ-

O

C

O

δ+

S

δ-

C

δ+

S

-

δ

O

S

-

δ

Muhammad Shahid A-One Professors Academy of sciences Jauharabad

O Resultant Dipole Moment

46) Dipole moment of CO2 is zero that of CO is 0.12D. Why? CO2 has linear structure. In this structure the individual bond moment cancel the effect of each other. Therefore dipole moment becomes zero. While in CO the electronegativity of oxygen is more than C. So the CO bond becomes polar and is directed from C to O. Hence the dipole moment of the CO is 0.12D. CO2 µ = 0 δ-

δ+

O

C

CO µ = 0.12 D δ-

O

δ+

C

δ-

O

47) The melting points, boiling points, heat of vaporization and heat of sublimation of electrovalent compounds are higher as compared with those of covalent bonds. Justify it In electrovalent compounds i.e. ionic compounds the ions are held together by strong electrostatic forces of attractions. A lot of energy is required to break these forces. While in covalent compounds. weak intermolecular forces are present. These forces are easy to break than ionic forces. Hence the melting points, boiling points, heat of vaporization and heat of sublimation of electrovalent compounds are higher as compared with those of covalent bonds.

CHAPTER # 6 CHEMICAL BONDING

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PROFESSOR EASY NOTES

Important Long questions of previous board papers (1) Write down the essential features of VSEPR theory also discuss the cases of water and ammonia. (2007) (2) How the bonding in the following molecules can be explained with respect to valence bond theory. (2008) (3) What is electronegativity? Discuss its variation in the periodic table. How it effects the bond strength. (2009) (4) Write the postulates of VSEPR theory and discuss the structure of NH3 with reference to this theory. (2010) (2015) (5) What is dipole moment? Give its various units also give its importance. (2011) (6) Give important postulates of VSEPR theory. (2012) (2014) (7) Define orbital hybridization. Explain the structure of CH4 on the basis of sp3 hybridization. (2012) (8) What is orbital hybridization? Explain sp3 hybridization with two examples. (2013) (9) Define bond energy and explain with the help of suitable example the effect of ionic character on its value. (2013) (10) Define dipole moment. Give its various units. How does it help to find out shapes of molecules? Give examples. (2015) (11) Write a note ionic bond by giving an example of KCl in detail. (2016) (12) Define ionization energy. How ionization energy varies along the group and across the period. (2017) (13) Explain Co-ordinate covalent bond. (2017) (14) Write down main points of valence shell electron pair repulsion (VSEPR) theory. (2018) (15) Draw shapes of following molecules according to VSEPR theory. (2018) (i) BeCl2 (ii) BF3 (iii) NH3 (iv) H2O

Things to remember:

CHAPTER # 6 CHEMICAL BONDING

64

PROFESSOR EASY NOTES

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 7

“thermochemistry” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 7 THERMOCHEMISTRY

67

PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) (1) For what purpose Bomb Calorimeter is used? (2007) Bomb Calorimeter is usually used for the accurate determination of the enthalpy of combustion for food, fuel and other substances. (2) Define exothermic and endothermic reactions. (2007) (2012) (2013) Those reactions in which heat is evolved from the system are called exothermic reactions. The sign of ∆H (enthalpy change) is negative e.g. burning of carbon in air to form CO2. Formation of ammonia in Haber’s process. While “Those reactions in which heat is absorbed from surroundings are called endothermic reactions”. The sign of ∆H is positive e.g. decomposition of water into hydrogen and oxygen. Reaction of N2 with O2 to form NO. N2 + O2  2NO (3) Define enthalpy of combustion (∆Hoc). Give one example. (2008) (2012) (2016) It can be defined as “The amount of heat evolved when one mole of a substance is completely burnt in excess of oxygen under standard conditions”. It is denoted by ∆H0c. Example: Muhammad Shahid Standard enthalpy of combustion (∆Hoc) of ethanol is -1368 kJmol-1. A-one Professors Academy of C2H5OH + 3O2  2CO2 + 3H2O ∆Hoc = -1368 kJ mol-1 sciences Jauharabad (4) Define enthalpy of solution (∆Hosol). Give two example. (2008) (2013) It can be defined as “The amount of heat absorbed or evolved when one mole of a substance is dissolved in so much solvent that further dilution results in no detectable heat change”. It is denoted by ∆Hosol. Example: Enthalpy of solution (∆Hosol) of ammonium chloride is + 16.2 kJmol-1 and that of sodium carbonate is -25.0 kJmol-1. (5) Define enthalpy of neutralization and why its value is same for dilute solution of a strong acid with a strong base. (2010) (2012) (2013) (2015) The heat of neutralization or standard enthalpy of neutralization is the amount of heat evolved when one mole of hydrogen ions (H+) from an acid, react with one mole of hydroxide ions (OH-) from a base to form one mole of water. It is denoted by ∆H0n. For example the enthalpy of neutralization of sodium hydroxide and hydrochloric acid is -57.4 kJ mol-1. H+ + Cl- + Na+ + OH-  H2O + Cl- + Na+ OR H+ + OH-  H2O The enthalpy of neutralization is merely the heat of formation of one mole of liquid water from its ionic components. The dilute solution of a strong acid and a strong base completely ionizes in aqueous solution to give H+ and OH- . The neutralization is essentially the reaction of H+ with OHand should therefore give the same value of ∆H0n per mole of water formed. (6) Explain that burning of a candle is a spontaneous process. (2010) (2017) A process is also called a spontaneous process if it need energy only to start, but once it is started then it proceeds on itself. A candle does not burn in air by itself, but the burning is initiated by a spark and once candle starts burning, then the burning of candle goes spontaneously to completion. So, burning of candle is a spontaneous process. (7) Define enthalpy of atomization (∆Hoat). Give example. (2010) (2014) It can be defined as “The amount of heat absorbed when one mole of gaseous atoms are formed from the element under standard conditions”. It is denoted by ∆Hoat. Example: ½ H2  H ∆Hoat = 218 kJ mol-1

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(8) Define state and state functions. (2011) (2017) OR What is meant by state function? Explain with two examples. (2009) The state of a system is the condition of a system. While state function is a macroscopic property of a system which has some definite values for initial and final state, and which is independent of the path adopted to bring about a change. Temperature, pressure, volume, energy and enthalpy are the examples of the state functions. By convention, capital letters are used as symbols for a state function e.g. temperature (T), pressure (P), internal energy (E). Change in a state function depends only on the initial and final state of a system. The change in volume (∆V) is given by ∆V = V2 – V1 (9) Define system and surroundings. (2013) (2018) Anything (material) under test in the laboratory or under consideration in the class room for the purpose of arguments is called a system. In other words, any real or imaginary portion of the universe, which is under study is called a system and the remaining portion of the universe is called surroundings. The real or imaginary surface separating the system from the surroundings is called the boundary. Water in a beaker is a system, the beaker and the air around it are forming its surroundings. (10) Define spontaneous and non-spontaneous processes? Give examples. (2014) (2015) (2016) (2017). A process which occurs by itself without any outside assistance and moves from a non-equilibrium state to an equilibrium state is called a spontaneous or natural process. It is unidirectional, irreversible and a real process. Once the equilibrium is established, the system will not change unless disturbed by some external aid e.g. water flows from a higher level to lower level. Neutralization of a strong acid by a strong base is a spontaneous acid-base reaction etc. While “The process which does not proceed by itself and is unnatural is called non-spontaneous process”, it can be made to occur only by supplying energy to the system from external source. It is the reverse of spontaneous process. Pumping of water uphill, transfer of heat from cold interior part of the refrigerator to the hot surroundings etc. (11) Explain that burning of coal is a spontaneous process. (2015) A process is also called a spontaneous process if it needs energy only to start, but once it is started then it proceeds on itself. Coal does not burn in air by itself, but the burning is initiated by a spark and once coal starts burning, then the burning of coal goes spontaneously to completion. So, burning of coal is a spontaneous process. (12) Is it true that a non-spontaneous process never happens in the universe? Explain It? (2018) OR Non-spontaneous process never happens in the universe. Justify it. Generally non-spontaneous process don not occur in nature by itself. A non-spontaneous process can be made to occur only by supplying energy to the system from external source, e.g. pumping of water uphill by supplying energy. There are some process which may occur in nature under specific conditions. For example nitrogen (N2) and oxygen (O2) are present in atmosphere but they do not react. This reaction can take place when energy is provided by lightning bolt. (13) What is meant by internal energy of a system? (2018) The sum of all the kinetic and potential energy of all the components of a system is called internal energy. The kinetic energy is due to the translational, rotational and vibrational movements of particles. The potential energy accounts for all the types of attractive forces present in the system. The internal energy is denoted by the symbol E. It is impossible to measure the absolute value of internal energy of a system, but it is possible to measure the value of ∆E for a change in the state of the system. (14) What is lattice energy of an ionic crystal? (2018) The lattice energy of an ionic crystal is the enthalpy of formation of one mole of the ionic compound from gaseous ions under standard conditions. Lattice energy of NaCl is -787 kJ mol-.

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Na+(g) + Cl- (g)  NaCl (s) ∆H lattice = -787 kJ mol-1 Lattice energies cannot be determined directly but values can be obtained indirectly by means of Born-Haber cycle. The lattice energy give us some idea of the force of attraction between cations and anions in crystalline ionic compounds. (15) Prove that ∆H = qp from first law of thermodynamics. (2018) The enthalpy of the system is given by H = E + PV A change in enthalpy of a system can be written as: ∆H = ∆E + ∆ (PV) ∆H = ∆E + V∆P + P∆V At constant pressure ∆P = 0, so ∆H = ∆E + P∆V------(i) According to first law of thermodynamics ∆E = q + w ---- (ii) If “w” is the pressure volume work done by the system, then w = - P∆V----(iii) By putting the value of “w” from equation (iii) in equation (ii) we will get ∆E = q - P∆V ---(iv) Now put the value of ∆E from equation (iv) in equation (i). ∆H = q - P∆V + P∆V or ∆H = q Since the pressure is constant, then ∆H = qp (16) State the first law of thermodynamics and give its mathematical form. This law states that energy can neither be created nor destroyed, but can be changed from one form to another. System can exchange energy with its surroundings in the form of heat and work. When a system undergoes any chemical or physical change, there is also change in its internal energy (∆E) which is equal to the heat (q) added or liberated from the system, plus the work done (w) “on” or “by” the system. ∆E = q + w (17) Define enthalpy? Total heat content of a system is called enthalpy (H). Enthalpy is equal to the internal energy (E) plus the product of pressure and volume (PV) H = E + PV Enthalpy is a state function and cannot be measured. However, enthalpy change (∆H) can be measured for a change in the state of system. A change in enthalpy of a system can be written as ∆H = ∆E + P∆V (at constant pressure) For liquids and solids, ∆V = 0, for such process, ∆H = ∆E The change in enthalpy (∆H) is equal to heat of reaction at constant pressure i.e. ∆H = qp (18) Define enthalpy of a reaction, ∆H. Can it be negative? How? The enthalpy change which occurs when the certain number of moles of reactants as indicated by the balanced chemical equation, react together completely to give the products under standard conditions i.e. 25 oC or 298 K and one atmosphere pressure is called standard enthalpy of a reaction (∆Ho). Yes, ∆H can be negative. In an exothermic reaction, the enthalpy of the products is less than that of the reactants. Since the system has lost heat, hence the enthalpy change, ∆H is negative. 2H2 (g) + O2 (g)  2H2O (liq) ∆Ho = -285kj mol-1 o (19) Define standard enthalpy of formation ∆H f. The standard enthalpy of formation of a compound is the amount of heat evolved or absorbed when one mole of a compound is formed from its elements. All the substances involved are in their standard physical states and the reaction is carried out under standard conditions i.e. at 250C or 298K and one atmospheric pressure. Its units are kJmol-1

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C(s) + O2 (g)  CO2 (g) ∆Hof = -393.7kj mol-1 (20) State why ∆H ≈ ∆E in case of liquids and solids. OR Is it true that ∆H and ∆E have the same values for the reactions taking place in the solution state? Yes, it is true that ∆H and ∆E have the same values for the reactions taking place in the solution state. A change in enthalpy of a system is ∆H = ∆E + P∆V The difference between ∆H and ∆E is a term PV, related to the expansion in case of liquids and solids. The change in state do not cause significant volume change i.e. ∆V ≈ 0. For such process ∆H and ∆E are approximately the same i.e. ∆H ≈ ∆E (21) Differentiate between (1) Internal energy and enthalpy and (2) Internal energy change and enthalpy change Sr. No

Internal energy

1

Internal energy is the sum of the kinetic and potential energies of the particles making of a system.

2

It is denoted by the symbol E

Sr. No 1 2

Change in Internal energy The change in internal energy (∆E) is defined as the heat (q) added to the system, plus the work (w) done on the system. ∆E = q + w

Enthalpy Total heat content of a system is called enthalpy. Enthalpy is related to the heat of reaction at constant pressure qp. It denoted by the symbol H, defined by the relationship. H = E + PV Enthalpy change The enthalpy change, ∆H, equals to heat (qp) when the process occurs at constant pressure, ∆H = qp ∆H = ∆E + P∆V

(22) Exothermic reactions are mostly spontaneous. Explain. A spontaneous change is accompanied by decrease of internal energy or enthalpy (∆H). This shows that only such reactions will occur which are exothermic. Exothermic reactions are accompanied by the loss of energy in the system. Therefore, exothermic reactions are mostly spontaneous. (23) It is essential to mention the physical states of reactants and products in a thermochemical equation. The quantity of heat evolved or absorbed during a reaction depends on the physical states of reactants and products. Therefore, it is necessary to mention the physical states of reactants and products in a thermochemical equation. Muhammad Shahid A-one Professors Academy of

H2 (g) + ½ O2 (g)  H2O (g) ΔH = -241.5 KJ mol-1 sciences Jauharabad H2 (g) + ½ O2 (g)  H2O (l) ΔH = -285.8 KJ mol-1 (24) What is the difference between heat and temperature? Write a mathematical relationship between these two parameters. Sr.No

2

Heat Heat is defined as the quantity of energy that flows into or out of a system during a change in its state. Heat is a not a state function.

Temperature Temperature is a property that determines the direction of heat flows into or out of the system. Temperature is a state function.

3

It is denoted by q.

It is denoted by T.

4

It is a path function.

It is a state function.

1

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(25) How do we determine the ∆H in the laboratory for food, fuel? A bomb calorimeter is used to determine the ∆H for food, fuel etc. A known mass of food or fuel is placed in the bomb with pressurized oxygen. The bomb calorimeter is then immersed in a known mass of water in well insulated calorimeter. Then it is allowed to attain a steady temperature. The initial temperature is noted. The food or fuel is then ignited electrically by passing the current through the ignition coil. The final temperature is measured after completion of reaction. The enthalpy of combustion of food is calculated by following formula. q = C × ∆T Where ‘c’ is the heat capacity.

Important Long questions of previous board papers 1. What is enthalpy of a reaction? Give one method for its measurement. (2008) (2009) OR 2. Describe the measurement of enthalpy of a reaction by glass calorimetric method. (2013) (2016) OR 3. Explain glass calorimetric method for the measurement of enthalpy of reaction. (2017) OR 4. Explain with diagram how enthalpy of a reaction can be measured by glass calorimeter. (2018) 5. When two moles of H2 and 1 mole of O2 at 100 oC and 1 torr pressure react to produce 2 moles of gaseous water, 484.5 kJ of energy as evolved what are values of (a) ∆H (b) ∆E for the production of one mole of H2O(g). (2012) 6. Define enthalpy of combustion. Write its measurement by Bomb Calorimeter. (2015) OR 7. Describe bomb calorimeter, for the calculation of enthalpy of a substance. (2018) 8. Define and explain Hess’s Law with example. (2007) (2010) (2011) (2014) OR 9. Explain Hess’s law of constant heat summation with examples and give its applications. (2015) 10. Define lattice energy and Born-Haber cycle. How Born Haber cycle help us to measure lattice energy. (2013) 11. What is “First law of thermodynamics”? Prove that ∆E = qv (2017) Things to remember:

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PROFESSOR SERIES

Professor easy notes for CHAPTER NO 08

“chemical equilibrium” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

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Important Short questions (Exercise is also added) 1. How the value of Kc of a reaction help to predict the direction of reversible reaction? (2007) Product The direction of a chemical reaction for a particular reaction can be predicted by means of Reactant ratio, calculated before the reaction attain equilibrium. There are three possibilities:  The ratio is less than Kc. It means that more of the product is required to attain the equilibrium, therefore the reaction will proceed in the forward direction.  The ratio is greater than Kc. It means that the reverse reaction will occur to attain the equilibrium.  When the ratio is equal to Kc, then reaction is at equilibrium. 2. How do you justify that the greater quantity of CH3COONa in acetic acid decreases the dissociating power of acetic acid and so the pH increases. (Exercise Question) OR When concentration of a salt is increased in an acidic buffer, then the pH of solution increases why? (2008) CH3COOH is weak acid and ionizes very small, while CH3COONa is a strong electrolyte and it ionizes in water to greater extent and provides acetate ions. CH3COOH + H2O CH3COO- + H3O+ CH3COONa CH3COO- + Na+ Thus, CH3COONa decrease the ionization of CH3COOH due to common ion (CH3COO-) and pH of solution increases. 3. Formation of NH3 from N2 and H2 is favoured at low temperature. Explain it. (2008) N2 + 3H2 2NH3 ∆H = -92.46 kJ/mol Since reaction is exothermic, according to Le-Chatlier’s principle, decrease in temperature favours forward reaction. The optimum temperature is 400 oC. So formation of NH3 from N2 and H2 is favoured at low temperature. 4. What is effect of presence of common ion on solubility? Give example. (2009) (2010) The suppression of ionization of a weak electrolyte by adding a common ion from outside is called common ion effect. The presence of a common ion decreases the solubility of less soluble salts. For example purification of NaCl. NaCl is completely ionized in the solution and Kc for this process can be written as NaCl Na+ + Cl[Na+ ] [Cl− ] Muhammad Shahid Kc = [NaCl]

A-One Professors Academy of

sciences Jauharabad HCl also ionizes in solution + HCl H + Cl On passing HCl gas, the concentration of Cl- ions is increased, and the solubility of NaCl will decreases, therefore NaCl crystalized out of the solution to keep the constant value of Kc. 5. Define pH and pOH. (2008) (2010) The pH of a solution can be defined as “The negative logarithm of hydrogen ion (H+) concentration”. pH = -log [H+] The pOH of a solution can be defined as “The negative logarithm of hydroxyl ion (OH-) concentration”. pOH = -log [OH-] 6. The change of temperature disturbs both the equilibrium position and equilibrium constant of reaction. Explain with reason. (2010) The enthalpy of a system can be affected by change in temperature. So, if temperature is changed, the reaction will proceed either in forward or backward direction to nullify that stress. As a result concentration of reactants and products will change without any substance being added or

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removed. Thus, the value of Kc will also be changed. Hence, both equilibrium position and equilibrium constant (Kc) will be changed. 7. HCl is used in a purification of NaCl. Explain. (2011) OR How NaCl is purified by common ion effect? (2012) The suppression of ionization of a weak electrolyte by adding a common ion from outside is called common ion effect. The presence of a common ion decreases the solubility of less soluble salts. For example purification of NaCl. NaCl is completely ionized in the solution and Kc for this process can be written as NaCl Na+ + Cl+ − Muhammad Shahid [Na ] [Cl ] Kc = A-One Professors Academy of [NaCl]

sciences Jauharabad HCl also ionizes in solution + HCl H + Cl On passing HCl gas, the concentration of Cl- ions is increased, and the ionization of NaCl will suppressed, therefore NaCl crystalized out of the solution to keep the constant value of Kc. Hence HCl is used in a purification of NaCl 8. Represent Henderson’s equation for acidic and basic buffer solution. (2011)

Henderson’s equation for acidic buffer solution:

The equation which gives the relationship between pH, pKa and concentration of an acid and its slat is called Henderson’s equation Salt pH = pKa + log Acid

Henderson’s equation for basic buffer solution:

The equation which gives the relationship between pOH, pKb and concentration of base and its slat is called Henderson’s equation. Salt pOH = pKb + log Base 9. What are buffer solutions? Why do we need buffer solutions? (2012) OR Buffer are important in many areas of chemistry. Justify it. (2013) Those solutions which resist the change in their pH when a small quantity of an acid or a base is added to them, are called buffer solutions. They have specific constant value of pH and their pH values do not change on dilution and on keeping for a long time. Need buffer solutions:  Human blood is buffered at pH 7.35. If it goes to 7 or 8, a person may die.  Some reactions required specific pH. Such reactions can be carried out in a buffer solution of specific pH.  Buffers are important in chemistry and in many other fields such as microbiology, cell biology, molecular biology, nutrition, soil sciences and clinical analysis. 10. What are buffer solutions? How an acidic buffer is prepared? (2011) Those solutions which resist the change in their pH when a small quantity of an acid or a base is added to them, are called buffer solutions. Preparation of an acidic buffer:

An acidic buffer can be prepared by mixing a weak acid and its salt with a strong base. They have pH less than 7. For example mixture of acetic acid and sodium acetate. 11. What is ionization constant of bases? (2012) Kb is the ionization constant of bases and can be calculated by the following method. Let the base is represented by B then, B + H2O BH+ + OHKc =

[BH+ ] [OH− ] [B][H2 O]

Since the concentration of H2O is constant, being in large excess so,

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Kc × [H2 O] =

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PROFESSOR EASY NOTES

[BH+ ] [OH− ] [B]

Kc × [H2 O] = Kb Kb =

[BH+ ] [OH− ] [B]

Kb value of a base is the quantitative measurement of strength of a base. 12. What are KC and Kp? Give their relationship. (2012) In reversible reactions the KC and KP are equilibrium constants. “C” shows concentration in moles dm-3, while “p” shows the partial pressure. KP = KC (RT)∆n Where ∆n is the difference of number of moles of product and number of moles of reactant. When number of moles of reactants and products are equal then KP = KC. 13. Explain the term buffer and buffer capacity. (2012) Buffers: Those solutions which resist the change in their pH when a small quantity of an acid or a base is added to them, are called buffer solutions or buffers. Buffer capacity: It can be defined as “The buffer capacity of a solution is the capability of a buffer to resist the change of pH. 14. How ammonia is synthesize by Haber process. Also give the optimum condition for reaction. (2013)

N2 + 3H2 2NH3 ∆H = -92.46 kJ/mol According Le-Chatlier’s principle the synthesis of ammonia is favoured at high pressure, low temperature. However its rate of formation is low. In order to increase the rate, catalyst is used. So the optimum conditions for the synthesis of ammonia (NH3) are given following:  Pressure = 200-300 atm  Temperature = 400 oC  Catalyst = Pieces of iron embedded in a fused mixture of MgO, Al2O3 and SiO2  Continuous removal of ammonia from the reaction mixture 15. Define pH with mathematical expression. (2013) The pH of a solution can be defined as “The negative logarithm of hydrogen ion (H+) ion concentration”. pH = -log [H+] 16. Why HCl is added before passing H2S gas in qualitative analysis 2nd group basic radical. (2013) In qualitative analysis 2nd group basic radical we required low concentration of sulphide ions (S-2). For this purpose in the solution of H2S gas we add HCl. H2S 2H+ + S-2 HCl H+ + ClDue to common H+ ion the ionization of H2S is suppressed. H2S becomes less dissociated in solution. In this way low concentration of sulphide ions (S-2) is developed. This low concentration of sulphide ions (S-2) help to do precipitation of radicals of 2nd group basic radicals during salt analysis. Therefore HCl is added before passing H2S gas in qualitative analysis 2nd group basic radical. 17. A catalyst does not affect an equilibrium constant. Comment on it. (2013) (2015) Mostly in reversible reactions, equilibrium is not reached in a short time. So a catalyst is used. It lowers the activation energy of the reaction by giving them a new path. It has no effect on equilibrium composition of a system. It simply increases the rate of forward or backward reaction. Hence it only decreases the time to reach the equilibrium state. So a catalyst does not affect an equilibrium constant. 18. Define law of mass action. (2014) (2018) Law of mass action can be defined as “The rate at which a substance reacts is directly proportional to its active mass and the rate of a reaction is directly proportional to the product of the active masses of the reacting substances”. The term active mass shows the concentration in mole dm-3.

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19. What is buffer solution? Give one example. (2015) Those solutions which resist the change in their pH when a small quantity of an acid or a base is added to them, are called buffer solutions or buffers. The mixture of acetic acid and sodium acetate is an example of acidic buffer. 20. What is common ion effect? Give example. (2015) (2018) (2018) The suppression of ionization of a weak electrolyte by adding a common ion from outside is called common ion effect. The presence of a common ion decreases the solubility of less soluble salts. For example purification of NaCl. NaCl is completely ionized in the solution and Kc for this process can be written as NaCl Na+ + ClKc =

[Na+ ] [Cl− ] [NaCl]

HCl also ionizes in solution HCl H+ + ClOn passing HCl gas, the concentration of Cl- ions is increased, and the ionization of NaCl will suppressed, therefore NaCl crystalized out of the solution to keep the constant value of Kc. Hence HCl is used in a purification of NaCl. 21. Describe the effect of temperature on the solubility of KI. (2015) KI (s) KI (aq) ∆H = 21.4 kJ/mol It is an endothermic reaction. Hence increase in temperature favours the forward reaction and vice versa. Hence more and more salt is dissolved by increasing temperature while cooling results in backward process. Thus KI(s) crystallized out. 22. How the value of equilibrium constant is related to the extent of reaction? (2012) (2015) OR How extent of a reversible chemical reaction can be indicated by equilibrium constant? (2017) The extent of a reversible reaction can be indicated by equilibrium constant in the following way:  If the equilibrium constant is very large, it indicates that the reaction is almost complete.  If the value of Kc is small, it shows that the reaction does not proceed appreciably in the forward direction.  If the value of Kc is very small, this shows a very little forward reaction. 23. What is the ionic product of water? How it changes by changing temperature? (2015) H2O H+ + OHKc =

[H+ ] [OH− ] [H2 O]

= 1.8 × 10-16 moles dm-3

Because the concentration of ions formed is very small and calculated to be 1000 g divided by 18 giving 55.5 mol dm-3. The concentration of H2O remains essentially constant so Kc [H2O] = [H+] [OH-] or 1.8 × 10-16 × 55.55 = [H+] [OH-] or 1.01 × 10-14 = [H+] [OH-] Where Kc [H2O] = Kw = ionization constant or ionic product of water. Thus, the ionic product of water (Kw) at 25 oC is Kw = [H+] [OH-] = 1.01 × 10-14 Effect of temperature: Value of Kw varies with temperature. It is because, ionization of water increases with increase in temperature. The Kw value increases almost 75 times when temperature is raised from 0 oC to 100 oC. However, increases in value of Kw is not regular. 24. Give statement of Le Chatlier’s principle. (2016) (2018) It can be defined as “If a stress is applied to a system at equilibrium, the system acts in such a way so as to nullify, as far as possible, the effect of that stress”.

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25. Calculate the pH of 10-4 mol/dm3 of Ba(OH)2. (2016) Ba(OH)2 Ba2+ + 2OH10-4 M solution of Ba(OH)2 produces 10-4 OH- ions. [OH-] = 2 × 10-4 M pOH = -log (2 × 10-4) = 3.698 pH + pOH = 14 pH = 14-3.698 = 10.302 26. What do you mean by acidic and basic buffer? Give one example of each. (2017) Acidic buffer:

An acidic buffer can be prepared by mixing a weak acid and its salt with a strong base. They have pH less than 7. The mixture of acetic acid and sodium acetate is an example of an acidic buffer. Basic buffer:

A basic buffer can be prepared by mixing a weak base and its salt with a strong acid. They have pH more than 7. The mixture of NH4OH and NH4Cl is an example of a basic buffer. 27. How solubility product can be determined from solubility? (2017) From the solubility of the compounds, we can calculate the solubility product (Ksp) by using following steps.  Solubilities are given in number of grams of solute/100 g of H2O.  Since the quantity of solute is very very small, so 100 g of water solution is considered to be 100 ml of solution. The reason is that the density of water is very close to unity.  From this amount of solute/dm3 is calculated.  The amount of solute in grams is converted into moles.  The moles of salts are used to calculate the molarity of ions.  Using balance chemical equation, Ksp can be calculated. 28. What is the solubility product? Derive solubility product expression for Ag2CrO4. (2017) OR Define solubility product giving at least one example. (2018) The solubility product can be defined as “The product of the concentration of ions raised to an exponent equal to the co-efficient of the balanced equation”. Ag2CrO4 2Ag+ + CrO42Muhammad Shahid + 2 [Ag ] [CrO2− 4 ] Kc = A-One Professors Academy of Ag2 CrO4

sciences Jauharabad

Kc × Ag 2 CrO4 = [Ag + ]2 [CrO2− 4 ] 2− + 2 Ksp = [Ag ] [CrO4 ] 29. Give the effect of pressure on the following reaction. (2017) PCl5 PCl3 + Cl2 Briefly explain the effect of pressure on the equilibrium position for the dissociation of PCl5. (2018) For the decomposition reaction of PCl5 the Kc is given as PCl5 PCl3 + Cl2

Kc =

x2 (a−x) v

The Kc expressions involves “V”, so if pressure is increased then volume will decreased according to Le-Chatlier’s principle, the reaction will proceed in backward direction to keep the constant value of Kc. Hence equilibrium position will change but equilibrium constant Kc will remain constant. 30. The change of volume disturbs the equilibrium position for some of the gaseous phase but not the equilibrium constant. Justify it. (2017) The reaction in gas phase are of two types. (i) The reactions in which number of moles of reactant and product are equal. Such reactions are not affected by change in volume. For example H2 + I2 2HI (ii) The reactions in which number of moles of reactant and products are not equal. Such reactions are affected by change in volume. For example

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PCl5

80

PCl3 + Cl2

Kc =

PROFESSOR EASY NOTES x2 (a−x) v

The Kc expressions involves “V” so change in volume will shift the equilibrium position of the gaseous system to minimize the effect of volume change. Reducing the volume of a gaseous equilibrium mixture causes the system to shift the direction that reduces the number of moles of gas and vice versa. For above reaction if pressure is increased then volume will decreased according to Le-Chatlier’s principle, the reaction will proceed in backward direction to keep the constant value of Kc. Hence equilibrium position will change but equilibrium constant Kc will remain constant. 31. Define lowry-bronsted acid base concept. (2018) (2018) An acid can be defined as “A substance which can donate a proton (H+) or have a tendency to donate a proton to another substance”. HF and H2S are acids because both gives proton (H+). Other examples include HCl HCN etc. A base can be defined as “A substance which can accept a proton (H+) or have a tendency to accept a proton from another substance”. NH3 and HCO3- are bases because both can accept proton (H+). Other examples include CO32- etc. 32. The solubility of glucose in water is increased by increasing the temperature. Justify it. Since the dissolution (‫ )لحوہان‬of glucose is endothermic process. So according to Le-Chatlier’s principle if temperature is increased, then system will move in forward direction. Thus, more glucose will dissolved. Hence solubility of glucose in water is increased by increasing the temperature.

Important Long questions of previous board papers

Muhammad Shahid A-One Professors Academy of sciences Jauharabad

1. The solubility of CaF2 in water at 25 oC is found to be 2.05 × 10-4 mol/dm3. What is the value of Ksp at this temperature? (2008) (2017) 2. What is the percentage ionization of acetic acid in a solution in which 0.1 moles of it has been dissolved per dm3 of the solution, Ka = 1.85 × 10-5 (2010) (2017) 3. Define law of mass action. Derive equilibrium constant expression for a general reaction. (2012) (2018) 4. Give the effect of temperature, pressure catalyst and concentration on the reaction. (2012) N2 + 3H2 2NH3 ∆H = -92.46 kJ/mol 5. The solubility of Ag2CrO4 is 2.6 × 10-2 at 25 oC. Calculate the solubility of the compound. (2015) Atomic masses of Ag = 107 g/mol Cr = 52 g/mol O = 16 g/mol 6. Ca(OH)2 is sparingly soluble compound. Its solubility is 6.5 × 10-6. Calculate the solubility of Ca(OH)2. (2016) 7. Define common ion effect. Give its two applications. (2018)

Things to remember:

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 09

“solutions” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 9 SOLUTIONS

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PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) 1) NaCl and KNO3 are used to lower the melting point of ice. Justify it. (2007) NaCl and KNO3 both are electrolyte and are highly soluble in water. When NaCl and KNO3 are added in water, the vapour pressure of water decreases. Due to lowering of vapour pressure, solution freezes below the freezing point of water. So it will have more cooling effect. So NaCl and KNO3 are used to lower the melting point of ice. 2) Why glucose is soluble in water but insoluble in CCl4? (2007) According to general rule of solubility “Like dissolves like” polar substances are soluble in polar solvents and non-polar are soluble in non-polar solvents. Water and glucose both are polar in nature. Glucose can make hydrogen bonding with water so it is soluble in water. While CCl4 is non-polar in nature so glucose is not soluble in CCl4. 3) What is cryoscopic constant? (2007) We know that depression in freezing point is directly proportional to molality of the solution. Tf ∝ m Muhammad Shahid Tf = Kf m A-One Professors Academy of When m = 1 Sciences Jauharabad Tf = Kf Where Kf is called cryoscopic constant and is defined as “The depression in freezing point produced when 1 mole of solute is dissolved in one kg (1000 g) of solvent. 4) The total volume of the solution by mixing 100 cm3 of water and 100 cm3 of alcohol may not be equal to 200 cm3. Justify it. (2008) (2015) There is no change in volume for ideal solutions. However, there is change in volume for non-ideal solutions. Alcohol and water form a non-ideal solution so forces of attraction between water and alcohol in pure state are greater than in solution. Thus, it shows positive deviation from Raoult’s law. During the formation of this solution heat is absorbed and volume decreases. Hence 100 cm3 of H2O and 100 cm3 of alcohol may not be equal to 200 cm3. 5) Define completely miscible and partially miscible solutions with suitable example. (2008) Completely miscible solutions: Liquids which dissolve together in all proportions are called completely miscible liquids and they form completely miscible solutions. For example solution of alcohol and water. Partially miscible solutions: The liquids which dissolve together to some extent are called partially miscible liquids and they form partially miscible solutions. For example phenol dissolved in water or ether dissolved in water. 6) Non- ideal solution do not obey Raoult’s law. Explain. (2009) (2017) A solution which does not obey Raoult’s law is called non ideal solution. Due to following reasons non-ideal solutions do not obey Raoult’s law.  The total volume of solution is not equal to the sum of volume of components of the solution.  It does not have zero enthalpy change i.e. ∆H ≠ 0.  Forces of attraction among molecules of components do not remain same in the solution just as they were in pure form. 7) Explain why CuSO4 give acidic solution when put in water? (2009) The hydrolysis of CuSO4 takes place as CuSO4 + H2O Cu(OH)2 + 2H+ + SO42This hydrolytic reaction produces weak base i.e. Cu(OH)2 and a strong acid i.e. H2SO4. Therefore CuSO4 give acidic solution when put in water. 8) Why some properties are called colligative properties? (2010)  Pure water has definite vapour pressure at a given temperature.

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 Let suppose 6 g of urea, 18 g of glucose and 34.2 g of sucrose are dissolved in 1 kg of H2O. 1 Each of this is 0.1 molal solution and contain equal number of solute particles i.e. th of 10 Avogadro’s particles i.e. 60.02 × 1022.  In all these cases, lowering of vapour pressure is same. The b.p of all the solution increases 0.0052 oC. The depression in freezing point is also same i.e. 0.186 oC.  Hence, lowering of vapour pressure, elevation in b.p and depression in freezing point are same in all cases due to same number of particles. So these are called colligative properties. 9) Define mole fraction and parts per million. (2010) (2011) Mole fraction: The mole fraction of any component in mixture is the number of mole of that component divided by total number of moles of all the components. Parts per million: It is defined as “The number of parts of a solute (by weight or volume) per million part (by weight or volume) of the solution”. 10) Differentiate between hydration and hydrolysis with one example each. (2010) Sr. No

1 2 3

4

Hydration

Hydrolysis

The process in which water molecules surround and interact with solute ions or molecules is called hydration. It does not change the pH of the solution. Physical interaction take place between solute and water molecules.

The process in which water molecules react chemically with the added solute is called hydrolysis. It changes the pH of the solution. Chemical reaction take place between solute and water molecules. The aqueous solution of CH3COONa is basic due to formation of strong base i.e. NaOH. The aqueous solution of AlCl3 is acidic in nature due to formation of a strong acid i.e. HCl.

For example CuSO4. 5H2O, Na2SO4.10H2O etc. In CuSO4. 5H2O four H2O are attached with Cu2+ and one with SO42- ion.

11) Calculate the percentage by weight of NaCl if 2 g of it is dissolved in 20 g of water. (2011) Weight of solute = 2 gram Weight of solvent = 20 gram Weight of solution = 20+ 2 = 22 gram Mass of solute (g) % by weight = × 100 Mass of solution (g) 2

% by weight = × 100 = 9.09% 22 % by weight = 9.09% 12) Give the condition of colligative properties. (2011) To observe the colligative properties following conditions should be fulfilled.  Solution should be dilute.  Solute should be non-volatile.  Solute should be non-electrolyte. 13) Define parts per million and when this concentration unit is used. (2011) (2018) It is defined as “The number of parts of a solute (by weight or volume) per million part (by weight or volume) of the solution”. This unit is used for very low concentration of solutions e.g. to express impurities of substances in water. 14) What is upper consulate temperature or critical solution temperature? Give example. (2012) (2013) The temperature at which two conjugate solutions merge into one another is called upper consulate temperature or critical solution temperature. For example for water phenol system the upper consulate temperature is 65.9 oC. 15) How will you prepare 5% W/V urea solution in water? (2013) The 5% w/v urea solution in water can be prepared by dissolving 5 gram of urea in 100 cm3 of solution. The quantity of solvent is not exactly known. In such solutions, the total volume of solution is under consideration i.e. the total volume of solution should be 100 cm3.

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16) What is the molality of the solution prepared by dissolving 5g glucose in 250g of water? ( 2013) Mass of glucose =5g Mass of water = 250 g = 0.25 kg Muhammad Shahid Molar mass of water = 18 g/mol A-One Professors Academy of Formula used: sciences Jauharabad

Molality (m) = Molality (m) =

Mass of solute (g)

Molar mass of solute(g mol−1 ) 5g 1 × 18(g mol−1 ) 0.250 kg

×

1

Mass of solvent in kg

Molality (m) = 1.11 mole kg-1 17) Na2SO4.10H2O shows discontinuous solubility curve. Reason. (2013) Sometimes, the solubility curves show sudden changes of solubilities and these curves are called discontinuous solubility curves. Na2SO4.10H2O shows discontinuous solubility curves. Actually the change of temperature changes the extent of hydration. These curves are combination of two or more solubility curves. At the break a new solid phase appears and another solubility curve begins. It is the number of molecules of water of crystallization which changes and hence solubility changes. So Na2SO4.10H2O shows discontinuous solubility curve. 18) Give two application of colligative properties. (2013) There are following two applications of colligative properties.  Use of ethylene glycol as an antifreeze in the radiator of automobile.  Use of NaCl and KNO3 to lower the melting point of ice. 19) One molal solution of urea, in water is dilute as compared to one molar solution of urea. Explain. (2015) (2016) One molal solution of urea contains 60 g of urea in 1 kg (1000 gram) of solvent. It means 60 g of urea is present in 1060 cm3 of solution, while 1 molar solution contains 60 gram of urea in 1000 cm3 of solution. Thus one molal solution of urea is more dilute than one molar solution of urea. 20) One molal solution of urea, in water is dilute as compared to one molar solution of urea. But the number of particles of the solute is same. Justify it. (2012) (2015) Since each solution contains one mole of urea, both have same number of particles i.e. 6.02 × 1023. One molal solution of urea contains 60 g of urea in 1 kg (1000 gram) of solvent. It means 60 g of urea is present in 1060 cm3 of solution, while 1 molar solution contains 60 gram of urea in 1000 cm3 of solution. Thus one molal solution of urea is more dilute than one molar solution of urea. 21) Why relative lowering of vapour pressure is independent of temperature? (2012) (2015) According to Raoult’s law: ∆P Relative lowering of vapour pressure is equal to the mole fraction of solute i.e. o = X2. P Where

∆P Po

= relative lowering of vapour pressure. While X2 =

n2

n1 + n2

= mole fraction of solute. Since

number of moles are not changed with temperature. Hence relative lowering of vapour pressure is independent of temperature. 22) Define colligative properties. Name some important colligative properties. (2014) (2015) (2016) The colligative properties are the properties of solution that depends on the number of solute and solvent molecules or ions. Important colligative properties:  Lowering of vapour pressure  Elevation of boiling point  Depression of freezing point  Osmotic pressure 23) Define molarity. Give one example. (2015) “Molarity is the number of moles of solute in 1 dm3 of the solution”. For example we want to prepare 1 molar solution of glucose. We take 180 gram of glucose and sufficient water to make the total volume 1 dm3 (1 litter) in the measuring flask.

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24) Differentiate between Zeotropic and Azeotropic solutions. (2015) Sr. No

Zeotropic solutions

Azeotropic solutions

1

Those liquids which distils with change in composition are called zeotropic mixtures.

The liquid mixtures, which boil at constant temperature and distill over without change in composition at any temperature like a pure compound, are called azeotropic mixtures.

2

For example methyl alcohol-water solution.

For example HCl-Water solution.

25) Define azeotropic mixture. Give two example. (2016) The liquid mixtures, which boil at constant temperature and distill over without change in composition at any temperature like a pure compound, are called azeotropic mixtures. For example HCl-Water solution and ethanol-water solution. 26) Why concentration in term of molality is independent of temperature while molarity depends upon temperature? (2017) (2018) Molality is defined as “The number of moles of solute dissolved in 1 kg of solvent”. Molarity is defined as “The number of moles of solute dissolved in 1dm3 of solution”. Since masses do not vary with temperature so concentration in term of molality is independent of temperature. While molarity changes with change in temperature because the expansion or contraction of the solution changes its volume. So molarity depends upon temperature. 27) Differentiate between ideal and non-ideal solution. (2017) Sr. No

Ideal solution

1

A solution which obey Raoult’s law is called an ideal solution.

2

The total volume of solution is equal to the sum of volume of components of solution.

3

It have zero enthalpy change i.e. ∆H = 0.

4

Forces of attraction among molecules of components remain same in the solution just as they were in pure form

Non-ideal solution A solution which does not obey Raoult’s law is called non ideal solution. The total volume of solution is not equal to the sum of volume of components of solution. It does not have zero enthalpy change i.e. ∆H ≠ 0. Forces of attraction among molecules of components do not remain same in the solution just as they were in pure form

28) Explain that boiling point of solvent increases due to presence of solutes. (2017) (2018) In pure solvent the whole surface of solvent is occupied by solvent molecules. However, when a nonvolatile, non-electrolyte solute is added to it, some surface is occupied by solute particles. So the escaping tendency of solvent is decreased. Therefore vapour pressure of solution becomes less than the pure solvent. Due to lowering of vapour pressure more heat is required to bring the vapour pressure of solution equal to external pressure. Hence boiling point of solvent increases due to presence of solutes. 29) What are hydrates? Give example. (2018) The crystalline substances, which contain chemically combined water in definite proportions is called a hydrate. For example CuSO4.5H2O, CaSO4.2H2O etc. 30) The sum of mole fractions of all the components is always equal to unity for any solution. Justify it. Suppose we have a solution of two components A and B. Let moles of A are nA and that of B are nB. Their mole fractions are given by: n XA = A --- (1) nA + nB

XB =

nB nA + nB

--- (2)

Their sum of mole fractions will be, XA + XB--- (3)

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Put equation (1) and (2) in equation (3) nA n + A nA + nB

nA + nB

or nA + nB nA + nB

=1

So the sum of mole fractions of all the components is always equal to unity for any solution. 31) 100 gram of 98% H2SO4 has a volume of 54.38 cm3. Justify it. 98% sulphuric acid has density = d = 1.84 g/cm3 Mass of H2SO4 = m= 100 gram Volume of H2SO4 = V? Muhammad Shahid As we know that m A-One Professors Academy of d= V=

V m

d 100

sciences Jauharabad

V= = 54.38 cm3 1.84 Hence 100 gram of 98% H2SO4 has a volume of 54.38 cm3. 32) Colligative properties are obeyed when the solute is non-volatile. Colligative properties depend upon the number of solute particles. If solute is electrolyte it will ionize and change the number of particles in the solution. Thus, value of colligative properties will be different from the expected values. So Colligative properties are obeyed when the solute is nonvolatile. 33) Colligative properties are obeyed when the solutions are dilute. Colligative properties depend upon the number of solute particles. In dilute solutions the solute particle behave independtly. While in concentrated solution solute particles do not behave independently. They have attraction for each other. Thus, normal colligative properties are only observed in dilute solutions. 34) Freezing points are depressed due to presence of solutes. Justify it. In pure solvent the whole surface of solvent is occupied by solvent molecules. However, when a nonvolatile, non-electrolyte solute is added, some surface is occupied by solute particles and its vapour pressure decreases. The solution will freeze at that temperature at which the vapour pressure of both liquid solution and solid solvent are same. Due to lowering of vapour pressure, solution freezes below the freezing point of pure solvent. 35) The boiling point of one molal urea solution 100.52 oC but the boiling point of two molal urea solution is less than 101.04 oC. Why? Colligative properties depend upon the number of solute particles. In dilute solutions the solute particle behave independtly. While in concentrated solution solute particles do not behave independently. They have attraction for each other. Hence the boiling point of one molal urea solution 100.52 oC but the boiling point of two molal urea solution is less than 101.04 oC. So the boiling point of one molal urea solution 100.52 oC but the boiling point of two molal urea solution is less than 101.04 oC. 36) Beckmann thermometer is used to note the depression in freezing point. Justify it. The depression in freezing point is very small change in temperature. Ordinary thermometer cannot read it. Beckmann thermometer can read upto 0.01K. Hence it can exactly measures the freezing point of pure solvent and solution. So Beckmann thermometer is used to note the depression in freezing point. 37) Why in summer the antifreeze solutions protect the liquid of the radiator from the boiling over. Antifreeze solution consists of solution of ethylene glycol in water. Since it is non-volatile, therefore vapour pressure of solution is lowered and boiling point increases. In summer water may boil over due to large heat in engine. However due to lower vapour pressure of solution, it requires greater heat to boil. So in summer the antifreeze solutions protect the liquid of the radiator from the boiling over.

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Important Long questions of previous board papers 1. Define the following terms with examples. (2007) (i) Molarity (ii) Molality (iii) Mole Fraction (iv) Parts per million 2. Discus the measurement of mass of solute by Beckmann’s method. (2008) 3. What is boiling point elevation? Describe a method to determine the boiling point elevation of a solution. (2009) (2010) (2012) (2016) 4. Explain the difference between ideal and non-ideal solution. (2010) 5. Discuss Raoult’s law for the solution in which both component are volatile. (2011) 6. What are colligative properties? Why are they called so? (2012) 7. Define non-ideal solutions and explain positive deviation with the help of graph. (2013) 8. Define hydrolysis. Why the aqueous solution of NH4Cl is acidic but that of CH3COONa is basic. (2013) 9. Define Raoult’s law and derive the equation ∆P = Po X2 (2014) 10. What is Raoult’s law? Give its three statements. How this law can help us to understand the ideality of a solution? (2015) 11. What are azeotropic mixtures? Explain them with the help of graph. (2015) 12. What is Raoult’s law? Give its three statements. (2017) 13. What do you mean by depression of freezing point of a solvent by a solute? Derive molecular mass of a compound by it. (2017) o 14. The boiling point of water is 99.725 C. To a sample of 600 gram of water are added 24 gram of a solute having molecular mass of 58 g mole-, to form as solution. Calculate the boiling point of the solution. (Hint: Example # 11) (2018) 15. The freezing point of Camphor is 178.4 oC. Find the freezing point of a solution containing 2.0 g of non-volatile compound, having molecular mass 140 in 40 g of camphor. The molal freezing point constant of Camphor is 37.7 oC kg mol-. (Hint: Example # 12) (2018)

Things to remember:

Muhammad Shahid A-One Professors Academy of sciences Jauharabad

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 10

“electrochemistry” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

CHAPTER # 10 ELECTROCHEMISTRY

91

PROFESSOR EASY NOTES

Important Short questions (Exercise is also added) (1) What is standard hydrogen electrode (SHE)? (2007) SHE stands for “standard hydrogen electrode” which is assigned a standard reduction potential of exactly 0 V and is based on the following half reaction. 2H+ (aq. 1M) + 2e-  H2 (g. 1 atm) Eored = 0 It consists of an electrode coated electrolytically with finely divided Pt in contact with H2 (g) at 1 atm pressure and 1M HCl solution at 25 oC. SHE is used as a reference electrode. (2) Write down the electrode reactions of silver oxide battery. (2008) The following reaction takes place in silver oxide battery. Zn + 2OH-  Zn(OH)2 + 2eAt anode Ag2O + H2O + 2e  2Ag + 2OH At cathode The overall reaction is Zn + Ag2O + H2O  Zn(OH)2 + 2Ag (3) How does electrochemical series tell us the distinction between the oxidizing and reducing agent? (2008) The value of the reduction potential of a metal or a non-metal tells us the tendency to lose electrons and act as a reducing agent. It also gives the information about the tendency of a species to gain electrons and act as an oxidizing agent. Greater the value of standard reduction potential of a given species, greater is its ability to accept electrons to undergo reduction and hence act as an oxidizing agent and vice versa. (4) Lead accumulator is a chargeable battery. Comment on it. (2008) (2013) (2015) (2017) The lead accumulator is a secondary battery. During the process of recharging, the anode and cathode of the battery are connected to the anode and cathode of the external electrical source respectively. The redox reactions takes place at the respective electrodes are reversed, regenerating Pb(s) and PbO2 (s) Muhammad Shahid At anode: A-one Professors Academy of PbSO4 + 2e-  Pb + SO42sciences Jauharabad At cathode: PbSO4(s) + 2H2O (l)  PbO2(s) + 4H+ (aq) + 2SO42-(aq) + 2eOverall reactions: 2PbSO4(s) + 2H2O (l)  Pb(s) + PbO2(s) + 4H+ (aq) + 2SO42-(aq) Recharging is possible because PbSO4 formed during discharge adheres to the electrodes. As the external source forces electrons form one electrode to another, the PbSO4 is converted to Pb at one electrode and PbO2 on the other. (5) What is the difference between a primary cell and secondary cell? (2009) (2016) (2018) A primary cell is not rechargeable for example alkaline battery, while a secondary cell can be rechargeable, for example lead accumulator and nickel cadmium battery. (6) Write anodic and cathodic reactions of galvanic cell. (2009) OR Why zinc oxidizes electrons and copper reduces in galvanic cell? (2009) Zinc has more ability to lose electron than copper. So zinc oxidizes and copper reduces. Zn  Zn2+ + 2eat anode 2+ Cu + 2e  Cu at cathode The overall reaction is the sum of these two half-cell reactions. Zn + Cu2+  Zn2+ +Cu This voltaic cell can be represented as follows. Zn | Zn2+ 1M || Cu2+ | Cu Eo = 1.1V

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(7) SHE acts as anode when connected with Cu electrode but as a cathode with Zn electrode. (2010) SHE has higher reduction potential than zinc (-0.76 volts). Thus, when both these are connected, electrons flow from Zn to SHE. Hence Zn acts as an anode and SHE as cathode. Zn  Zn+2 + 2eat anode + 2H + 2e  H2 at cathode SHE has lower reduction potential than Cu (+0.34) Thus, when connected, electrons flow from SHE to Cu. Hence SHE acts as anode and Cu as cathode. Cu+2 + 2e Cu at cathode + H2  2H + e at anode (8) What is industrial importance of electrolysis? (2010) Electrolysis has following industrial importance:  Extraction of sodium by the electrolysis of fused sodium chloride.  Aluminum is extracted by electrolyzing fused Bauxite, Al2O3.2H2O.  It is used for the purification of copper.  Copper, silver, nickel and chromium plating is done by various types of electrolytic cells. (9) Define the oxidation number with an example. (2010) It is apparent charge on an atom of an element in a molecule or an ion. It may be zero, positive or negative. For example in H2S the oxidation number of S is -2 and that of H is +1. (10) Calculate the oxidation number of Mn in KMnO4 and K2MnO4. (2010)

Oxidation number of Mn in KMnO4 Oxidation number of Mn in K2MnO4 Oxidation number of K = +1 Oxidation number of K = +1 Oxidation number of O = -2 Oxidation number of O = -2 Oxidation number of Mn = ? Oxidation number of Mn = ? [O.N of K] + [O.N of Mn] + 4[O.N of O] = 2[O.N of K] + [O.N of Mn] + 4[O.N of O] = 0 0 2 [+1] + [O.N of Mn] + 4[-2] = 0 [+1] + [O.N of Mn] + 4[-2] = 0 [O.N of Mn] -8 +2 = 0 [O.N of Mn] -8 +1 = 0 Muhammad Shahid [O.N of Mn] -6 = 0 [O.N of Mn] -7 = 0 O.N of Mn = +6 A-one Professors Academy of O.N of Mn = +7 So O.N of Mn is +6 sciences Jauharabad So O.N of Mn is +7 (11) Prove that oxidation number of some elements vary in different compounds. (2011) Sulphur has different oxidation number in its different compounds. In H2S it has -2, in SO2 +4 and in H2SO4 it have +6. Oxidation number of S in H2S. Oxidation number of S in H2SO4. (2014) Oxidation number of H = +1 Oxidation number of H = +1 Oxidation number of S =? Oxidation number of O = -2 2[O.N of H] + [O.N of S] = 0 Oxidation number of S =? 2[+1] + [O.N of S] = 0 2[O.N of H] + [O.N of S] + 4[O.N of O] = 0 [O.N of S] +2 = 0 2[+1] + [O.N of S] + 4[-2] = 0 O.N of S = -2 [O.N of S] - 8 +2 = 0 [O.N of S] - 6 = 0 So O.N of S in H2S is -2. Oxidation number of S in SO2. O.N of S = +6 Oxidation number of O = -2 So O.N of S in H2SO4 is +6. Oxidation number of S =? 2[O.N of O] + [O.N of S] = 0 2[-2] + [O.N of S] = 0 [O.N of S] - 4 = 0 O.N of S = +4 So O.N of S in SO2 is +4.

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(12) What is electrolytic conduction? (2011) The movement of ionic charges through the electrolyte brought by the application of electricity is called electrolytic conduction. In other words, electrolytic conduction is the passage of electric current through electrolytes in the fused state or in the solution form of ionic compounds. (13) Calculate oxidation number of Chromium in K2CrO4 and K2Cr2O7. (2012) Oxidation number of Cr in K2CrO4 Oxidation number of K = +1 Oxidation number of O = -2 Oxidation number of Cr = ? 2[O.N of K] + [O.N of Cr] + 4[O.N of O] = 0 2[+1] + [O.N of Cr] + 4[-2] = 0 [O.N of Cr] -8 + 2 = 0 [O.N of Cr] -6 = 0 [O.N of Cr] = +6 So O.N of Cr is +6

Oxidation number of Cr in K2Cr2O7 Oxidation number of K = +1 Oxidation number of O = -2 Oxidation number of Cr = ? 2[O.N of K] + 2[O.N of Cr] + 7[O.N of O] = 0 2[+1] + 2[O.N of Cr] + 7[-2] = 0 2[O.N of Cr] -14 + 2 = 0 2[O.N of Cr] -12 = 0 2[O.N of Cr] = +12 12 O.N of Cr = = 6 2 So O.N of Cr is +6

(14) Differentiate ionization and electrolysis. (2012) The process in which molten ionic compounds or dissolved in water split up into charged particles is called ionization. NaCl  Na+ + ClWhile when a non-spontaneous reaction takes place at the expense of electrical energy, is called electrolysis. The products are deposited at respective electrodes and electrolyte is decomposed. (15) Define oxidation and give an example. (2012) Reduction can be defined as “Addition of hydrogen or removal of oxygen during a chemical reaction is called reduction”. For example Alo  Al3+ + 3eZno  Zn2+ + 2eReduction

Removal of Oxygen

2ZnO + C  2Zn + CO2 Oxidation

Addition of Oxygen

Addition of hydrogen Reduction

H2S + Cl2  S + 2HCl Oxidation

Removal of hydrogen

(16) What is electrochemical series? (2013) When elements are arranged in the order of their increasing standard reduction potentials on the hydrogen scale, the resulting list is known as electrochemical series. The electrode potentials have been given in the reduction mode as recommended by the IUPAC. (17) What is salt bridge? What is the function of salt bridge in galvanic cell? (2014) (2017) OR (18) Describe salt bridge. (2017) A salt bridge is a U-shaped tube. It consists of saturated solution of a strong electrolyte (KCl or KNO3) supported in a jelly type material. The ends of the U-tube are sealed with a porous material like glass wool.

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A salt bridge have following function: 1. It allows electrical contact between two electrolytic solutions. 2. It prevents mixing of the two electrolytic solutions. 3. It maintains electrical neutrality in each half-cell. (19) Define electrolysis and give example. (2015) When a non-spontaneous reaction take place at the expense of electrical energy, the process is called electrolysis. The products are deposited at respective electrodes and electrolyte is decomposed. For example electrolysis of molten NaCl in Down’s cell. (20) Na, K can displace hydrogen from acids but Pt, Pd and Cu cannot. Why. (2015) (2017) Greater the value of standard reduction potential of a metal, lesser is its tendency to lose electrons to form metal ions and so weaker is its tendency to display hydrogen from acids. Pt, Pd and Cu which have sufficiently high positive values of reduction potentials, cannot displace hydrogen from acids. While, Na, K metals which have very low reduction potentials. Liberate hydrogen gas when they react with acids. 2Na +2HCl  2NaCl + H2 2K +2HCl  2KCl + H2 (21) Write down the equation involved in electrolysis of aqueous solution of NaCl. (2015) Following reaction is involved in electrolysis of aqueous solution of NaCl. NaCl Na+ + ClAt anode: 2Cl-  Cl2 + 2eAt cathode: 2H2O + 2e-  H2 + 2OHThe overall reaction is 2Na++ 2Cl- + 2H2O  Cl2 + H2 + 2Na+ + 2OH(22) Define standard electrode potential? (2016) The potential set up when an electrode is in contact with one molar solution of its own ions at 298K, is known as standard electrode potential or standard reduction potential of the element. It is represented as Eo. Standard electrode potential of hydrogen has arbitrarily been chosen as zero, while the standard electrode potentials of other elements can be determined by comparing them with standard hydrogen electrode. (23) Standard oxidation Potential of Zn is + 0.76 volts and its reduction potential is – 0.76 volts. (2016) The ability to lose electrons is called oxidation potential and ability to gain electrons is called reduction potential. According to the law of conservation of energy, energy can neither be created nor destroyed. Therefore, if standard oxidation potential of Zn is 0.76 V, then standard reduction potential will also be same but with opposite sign. Thus oxidation and reduction potential are always equal but opposite in sign. Zn  Zn+2 + 2e(Oxidation) Eo Oxidation = +0.76 +2 Zn + 2e  Zn (Reduction) Eo Reduction = -0.76 (24) Zn can displace hydrogen from dilute acid solution but “Cu” cannot? (2018) An element lies above in the electrochemical series can displace the element lies below in electrochemical series form its solution. Zinc can displace Hydrogen from dilute acid solution because zinc lies above than Hydrogen in electrochemical series. But Cu cannot displace Hydrogen from dilute acid solution because Cu lies below than Hydrogen in electrochemical series. (25) Calculate the oxidation umber of “Cr” in (i) CrCl3 (ii) K2Cr2O7 (2018) Oxidation number of Cr in CrCl3 Oxidation number of Cl = -1 Oxidation number of Cr = ?

Oxidation number of Cr in K2Cr2O7 Oxidation number of K = +1 Oxidation number of O = -2

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[O.N of Cr] + 3[O.N of Cl] = 0 [O.N of Cr] + 3[-1] = 0 [O.N of Cr] -3 = 0 [O.N of Cr] = +3 So O.N of Cr is +3

Oxidation number of Cr = ? 2[O.N of K] + 2[O.N of Cr] + 7[O.N of O] = 0 2[+1] + 2[O.N of Cr] + 7[-2] = 0 2[O.N of Cr] -14 + 2 = 0 2[O.N of Cr] -12 = 0 2[O.N of Cr] = +12 12 O.N of Cr = = 6 2 So O.N of Cr is +6 (26) How copper can be purified electrolytically? (2018) Electrolytic cell is used for the purification of copper. Impure copper is made the anode and a thin sheet of pure copper is made the cathode. These electrodes are dipped in CuSO4 solution. The atoms of Cu from impure Cu anode are converted to Cu2+ ions and migrate to cathode which is made of pure copper and deposited there. In this way Cu anode is purified. At anode: Muhammad Shahid Cu  Cu2+ +2eA-one Professors Academy of At cathode: 2+ sciences Jauharabad Cu +2e  Cu (27) How is Al anodized in an electrolytic cell? Anodized aluminum is prepared by making Al metal an anode in an electrolytic cell containing sulphuric acid (H2SO4) or Chromic acid (H2CrO4). It coats a thin layer of oxide on it. This aluminum oxide layer is an anodized Al. The aluminum oxide layer resists attack for corrosion. (28) What is electroplating? It can be defined as “The process of depositing (‫ )ہتامجان‬of one metal over the other by means of electrolysis is called as electroplating” . Principle The principle of electroplating is to establish an electrolytic cell in which anode is made of the metal to be deposited and cathode of the object on which metal is deposited. The electrolyte is an aqueous solution of a salt of the respective metal (‫)ہقلعتملٹیم‬. USES: It is used for silver plating of jewelry, steel and tableware etc. (29) A porous plate or salt bridge is not required in lead storage cells. Salt bridge maintains electrical neutrality in each half cell without physical contact between two solutions. In lead storage batteries electrolyte is same for both half cell (anode and cathode component). Further when oxidation takes place at anode, Pb2+ ions are produced and they pickup SO4-2 ions from solution and get deposit as PbSO4 at anode. The solution does not become positively charged. At the same time H+ from electrolyte (H2SO4) are changed into H2O and PbO2 into PbSO4 at cathode. Thus both half-cell remain neutral. Therefore there is no need of salt bridge. (30) An equilibrium is established when a plate is dipped in a solution of its own ions. (2018) When a plate is dipped in a solution of its own ions, the plate loses electron and get converted into positive ions. At the same time, ions from solution pick up electrons and get converted into metal. Thus an equilibrium is established between metal and its ions. Therefore no current is generated. Zn  Zn+2 + 2eEo Oxidation = +0.76 Zn+2 + 2e Zn Eo Reduction = -0.76 (31) Transition elements act as Anode in alkaline battery? Most of transition elements have negative reduction potential. It means they have high tendency to lose electrons. Thus in alkaline batteries transition metal plate (Zn, Cd) act as source of electron and oxidation takes place so transition elements act as Anode in alkaline battery.

CHAPTER # 10 ELECTROCHEMISTRY

Zn Cd

96

 

PROFESSOR EASY NOTES

Zn+2 + 2eCd+2 + 2e-

Some transition elements are highly inert like Pt, Pd. They are used as inert electrodes. An inert electrode conducts electron to and from external circuit without actually taking part in reaction. (32) A salt bridge maintains electrical neutrality in cell? (2012) (2018) Two half cells are connected by a salt bridge. Consider a Zn-Cu cell. During reactions of this cell, zinc half-cell continuously loose electrons. Thus, in this half-cell positive charge is increasing. Zn  Zn+2 + 2eWhile, Copper half-cell continuously receive electrons, thus it goes on collecting negative charge. Cu2+ +2e-  Cu Collection of positive charge in Zn half-cell and collection of negative charges in copper half-cell would stop the redox reaction. Salt bridge prevents the net accumulation of charges in either half-cell. Thus from negative Cu halfcell, the negative ions diffuse through the salt bridge into the Zn half-cell. In this way salt bridge maintains the two solutions electrically neutral. (33) What is electrochemistry? Electrochemistry is the branch of chemistry which deals with the conversion of electrical energy into chemical energy in an electrolytic cells as well as conversion of chemical energy into electrical energy in a galvanic cell or voltaic cells. (34) What is the difference between a battery and a cell? A battery is a portable, self-contained electrochemical power source that consists of one or more voltaic cells. Greater voltages can be achieved by using multiple voltaic cells in a single battery. 6 voltaic cells are connected in series produces 12 V current. An electrochemical cell is a system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current. Strictly speaking, cells are not self-contained systems e.g. fuel cell. (35) What is alkaline battery or dry alkaline cell? In dry alkaline cell, zinc rod acts as anode and manganese dioxide as cathode. The electrolyte is KOH. The voltage of cell is 1.5 volts. The reactions are: At anode: Zn + 2OH Zn (OH) 2 + 2e- (oxidation) At cathode 2MnO2 + H2O + 2e-  Mn2O3 + 2OH- (oxidation) Overall reaction: Zn + 2MnO2 + H2O  Zn (OH) 2 + Mn2O3 (36) What is fuel cell? How the fuel cell can be used as drinking water for an astronaut? A galvanic cell in which the reactants are continuously fed into the cell as the cell produces electrical energy is called a fuel cell. At the electrodes the hydrogen is oxidized to water and oxygen is reduced to OH- ions. This fuel cell is operated at high temperature so that the water formed as a product of the cell reaction evaporates and may be condensed and used as drinking water for an astronaut. At anode: H2 + 2OH-  2H2O + 2e- ] × 2 Muhammad Shahid At cathode: A-one Professors Academy of O2 + 2H2O + 4e-  4OHsciences Jauharabad Overall reaction: 2H2 + O2  2H2O (37) How is the Daniel cell represented? Galvanic cells are represented by a simple notation called cell diagram. The notation for the Daniel cell is Zn /Zn2+ 1M || Cu2+ / Cu Eo = 1.1V

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The vertical lines represent phase boundaries and a double line “||” represents a salt bridge. The anode or oxidation half-cell (Zn anode) is always written to left. The cathode (Cu) or reduction half-cell is written on the right side. (38) What are redox reactions? The chemical reactions in which oxidation and reduction processes takes place are called redox reactions. For example reaction of magnesium with oxygen. The oxidation state of Mg increases from zero to +2, hence it is oxidation, while the oxidation state of oxygen decreases from zero to -2, hence it is reduction. 0

0

+2 -2

2Mg + O2  2MgO (39) What is an anode and cathode? In an electrochemical cell, the electrode at which oxidation occurs is called the anode; the electrode at which reduction occurs is called cathode. (40) What is an electrochemical cell? An electrochemical cell is a system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current. These are of two types (1) A voltaic or galvanic current (2) Electrolytic cell (41) Zn can displace Cu from CuSO4, while Zn does not displace Mg from MgSO4 solution. Why? An element lies above in the electrochemical series can displace the element lies below in electrochemical series form its solution. Thus Zn metal which lies above in the electrochemical series have lower standard reduction potential than that of Cu. It will displace Cu from aqueous solution of CuSO4. While Zn lies below in electrochemical series have greater standard reduction potential than of Mg and therefore Zn cannot displace Mg from aqueous solution of MgSO4. Zn + CuSO4  ZnSO4 + Cu Zn + MgSO4  No reaction (42) What is electromotive force; emf? The maximum potential difference between the electrodes of a voltaic cell is called electromotive force (emf) of the cell, or E cell. Because E cell is measured in volts. It is also called the cell voltage. E cell is also called the cell potential. The potential difference between the two electrodes of a voltaic cell provides the driving force that pushes electron through the external circuit. (43) Differentiate between oxidizing agent and reducing agent? Oxidizing agent is that substance, which oxidizes other substance and reduced itself in a chemical reaction, while the substance which reduces a substance and is oxidized itself in a redox reaction is called reducing agent.

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Important Long questions of previous board papers 1. What is standard hydrogen electrode? Give its construction and working. (2008) 2. Describe the construction and working of standard hydrogen electrode. (2009) 3. Balance the following equation by ion electron method. (2010) (i) Fe+3 + Sn+2  Fe+2 + Sn+4 (ii) Cu + NO3-1  Cu+2 + 2NO2 4. Give the four industrial applications of electrolysis process. (2012) (2018) 5. Describe the electrolysis of molten NaCl and a concentrated solution of NaCl. (2013) 6. What is electrochemical series? Also explain its three applications. (2013) 7. Explain fuel cell with its construction, electrolytic reaction and diagram. (2014) 8. Define electrode potential. Describe the construction of voltaic cell and reaction occurring in the cell. (2015) 9. Write a comprehensive note lead accumulator. (2015) 10. Briefly explain any four industrial importance of electrolytic process. (2016) 11. Define electrolysis. Explain the electrolysis of very dilute solution of NaNO3. (2017) 12. Define standard electrode potential. Explain the measurement of electrode potential of copper. (2017) 13. Write a note on fuel cells. (2018)

Things to remember:

PROFESSOR SERIES

Professor easy notes for CHAPTER NO 11

“reaction kinetics” Class 11th “Solved Exercise + Old papers of SGD board"

Things to remember:

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Important Short questions (Exercise is also added) Muhammad Shahid 1) What is negative catalyst? Give an example. (2008) A-One Professors Academy of It can be defined as Sciences Jauharabad “When the rate of reaction is retarded by adding a substance, then it is said to be a negative catalyst”. It is also known as inhibitor. For example, to save petrol from pre-ignition, tetraethyl lead is added to petrol. 2) Define specific rate constant. What is the effect of temperature on it? (2008) OR What is specific rate constant or velocity constant? (2018) It can be defined as “Specific rate constant of a chemical reaction is the rate of reaction when the concentrations of the reactants are unity”. Under given conditions it remains constant. It changes with change in temperature and vice versa. 3) Define homogenous catalysis with two examples. (2009 ) (2016) (2018) OR Define heterogeneous catalysis with example. (2013) OR Differentiate between homogeneous catalysis and heterogeneous catalysis. (2014) (2017) Homogeneous catalysis: In this process, the catalyst and the reactant are in the same phase and the reacting system is homogeneous throughout. The catalyst is distributed uniformly throughout the system. Mostly liquids and gases are used as homogeneous catalysis. For example: 2SO2 (g) + O2 (g)

NO (g)

2SO3 (g)

Both the reactants and the catalyst are in same phase i.e. gas phase. Heterogeneous catalysis: In this catalyst, reactants and catalyst are in the different phases. In this the system remains heterogeneous during the reaction. Mostly the catalysts are in solid phase and reactants are in gaseous or liquid phase. For example Ni (s) H2C

CH2

(g)

+

H2 (g)

H3C

CH3

(g)

Here the Ni (Nickle) is in solid state while all the others are in gaseous state. 4) Differentiate between average and instantaneous rate of reaction. (2010-11-12-13-14-15-17) The rate any one instant during the interval is called the instantaneous rate while the rate of reaction of between two specific time intervals is called rate of reaction. The average rate and instantaneous rate are equal for only one instant in any time interval. At first the instantaneous rate is higher than the average rate but at the end the instantaneous rate becomes slower than the average rate. 5) How light affects the rate of a reaction? (2010) Light consist of photons having definite amount of energies depending upon their frequencies. When light falls on reactants, this energy becomes available to them and rate of reaction is enhanced. For example reaction of H2 and Cl2 requires light. The reaction is negligible in darkness between H2 and Cl2. Slow in day light, but explosive in sun light. Similarly reaction between CH4 and Cl2 requires light. Light plays vital role in the process of photosynthesis. 6) Give two characteristics of enzyme catalysis. (2012) The role of enzyme is like inorganic heterogeneous catalyst. They are unique, highly efficient and specific in their action. They have following characteristics.  They lower the energy of activation.  They are highly specific in their action.

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 They have maximum rate of reaction at an optimum temperature.  Their catalytic property greatly enhanced by the presence of a co-enzyme or activator. 7) Compare order of reaction and molecularity. (2012) Sr.No

Order of reaction

Molecularity It is the number of atoms, ions or molecules that must collide with one another simultaneously as a result chemical reaction takes place.

1

It is sum of the concentration terms on which the rate of reaction actually depends.

2

It may be zero, in fraction or whole number.

It is always a whole number.

3

It can be determined experimentally only and cannot be calculated by looking at reaction equation.

It can be calculated by simply adding the molecules of the slowest step.

4

It is for the overall reaction and no separate steps are written to obtain it.

The overall molecularity of a complex reaction has no significance. It is only slowest step whose molecularity has significance for the overall reaction.

8) Describe two characteristics of a catalyst. (2013) (2017)  A catalyst is more effective in finely divided form. It is because increase in surface area increases the efficiency of a catalyst and rate of reaction increases e.g. a big piece of Pt have much less catalytic activity than colloidal Pt.  Sometimes only a trace of a metal catalyst is used to affect very large amount of reactants e.g. thousands of dm3 of H2O2 can be decomposed in the presence of 1 g of colloidal platinum. 9) What is auto-catalyst? Give one example. (2014) (2017) “In some of the reactions, a product formed acts as a catalyst. It is called auto-catalyst and this phenomenon is called auto-catalysis. For example when copper is allowed to react with nitric acid, the reaction is slow in the start. It gains speed gradually and finally becomes very fast. This happens due to the formation of nitrous acid during the reaction which acts as a auto-catalyst accelerates the process. 10) How surface area affect the rate of reaction give example. (2015) With increase in surface area of reactants, the possibility of atoms and molecules of reactants to come in contact with each other is increases. So rate of reaction is increases e.g. CaCO3 in the powder form reacts with dilute H2SO4 more efficiently than its big pieces. 11) A particular catalyst is suitable for a particular reaction. Justify it. OR A catalyst is specific in its action. Explain. (2015) A catalyst is specific in its action. A particular catalyst works for one reaction, it may not work for any other reaction. If different catalysts are used for the same reaction then products may change. Example: Formic acid is decomposed to different products when we use different catalysts e.g. HCOOH HCOOH

Al2 O3 Cu

H2O + CO H2

+ CO2

Muhammad Shahid A-One Professors Academy of Sciences Jauharabad

12) Define instantaneous rate. (2016) The rate of reaction at any one instant during the interval is called instantaneous rate.

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Conc

13) Rate reaction is even Changing Parameter under the given conditions? Justify it. (2016) Rate of reaction depends upon molar concentration of reactants. As reaction proceeds, reactants are converted into products. Thus, with passage of time the molar concentration of reactants decreases. Thus instantaneous rate is greater at point “a” smaller at point "b” and least at point “c” as shown in graph.

c b a

Time

14) How higher temperature increases the rate of reaction? (2017) At low temperature molecules possess average energy and only small fraction of molecules have necessary activation energy (Ea) and rate of reaction is low. When temperature rises the fraction of high energy molecules increases and rate of reaction increases. When the temperature is raised 10K, the fraction of molecule with more energy than Ea roughly doubles and so the reaction rate is also doubled so higher temperature increases the rate of reaction. 15) Radioactive disintegration is always first order? Justify (2017) OR Why radioactive disintegration is always first order? (2018) The half-life of reaction is inversely proportional to initial concentration of reactants raised to the power one less than the order of reaction. 1 t1/2 ∝ 𝑛−1 𝑎

1

1

For first order reaction t1/2 ∝ 1−1 or t1/2 ∝ 0 𝑎 𝑎 It means that half-life of first order reaction is independent of initial concentration. It has been observed that disintegration of radioactive material is first order as its half-life is 235

235

U U independent of amount of radioactive material e.g half-life of 92 is 710 million years. If one kg of 92 disintegrates then 0.5 kg of it is converted to daughter elements in 710 million years. Out of 0.5 kg of 235

U

92

0.25 kg disintegrates in next 710 million years. 16) The units of rate constant of second order reaction is dm3 mol-1s-1 but the unit of rate of reaction is mol dm-3s-1. (2018) Rate of reaction is the change in concentration with change in time so its units are given as ∆𝒄

𝒎𝒐𝒍 𝒅𝒎−𝟑

Rate = = = mol dm-3s-1. ∆𝒕 𝒔𝒆𝒄 While in case of second order reaction Rate = k [A] [B] 𝐑𝐚𝐭𝐞

k = [𝐀][𝑩] =

𝒎𝒐𝒍 𝒅𝒎−𝟑 𝒔−𝟏 (𝒎𝒐𝒍 𝒅𝒎−𝟑 ) (𝒎𝒐𝒍 𝒅𝒎−𝟑 )

= dm3 mol-1s-1

Muhammad Shahid A-One Professors Academy of Sciences Jauharabad

17) What is zero order reaction? Give one example. (2018) When the reaction is independent of the concentration of reactants, it is known as zero order reaction. Photochemical reactions are usually zero order reactions.

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18) Define order of reaction. The order of reaction may be defined as “The number of reacting molecules, whose concentrations alter as a result of the chemical change. It is given by the sum of all the exponents to which the concentrations in the rate equation are raised. 19) Define chemical kinetics. The study of reaction rates, factor affecting the reaction rates and mechanism of chemical reactions is called chemical kinetics. 20) Define activation energy and activated complex. The minimum amount of energy which is required to start a chemical reaction is called activation energy. While activated complex is an unstable combination of all the atoms involved in the reaction for which the energy is maximum. It is short lived species and decomposes into products immediately. 21) Define rate law. The experimental relationship b/w a reaction rate and the concentration of the reactants is known as the rate law or the rate equation for that reaction. 22) Differentiate between fast step and slow step. Sr.No 1

Fast step In the reaction mechanism the rate does not depends upon this step.

Slow step It is the rate determining step in the reaction mechanism.

23) Differentiate between enthalpy change of a reaction and energy of activation of a reaction. Sr.No

1

Enthalpy change of a reaction Enthalpy change is the name given to the amount of heat evolved or absorbed in a reaction carried out at constant pressure. It is shown by the symbol ΔH, read as "delta H".

Energy of activation Energy of activation is the minimum amount of energy which is required to start a chemical reaction. It is shown by the symbol “Ea”.

24) Reaction rate decreases every movement but rate constant “K” of reaction is a constant quantity under the given conditions? Rate of reaction decreases with passage of time because reactants are converted into products with the decrease in concentration of reactants. But the rate constant is defined as rate of reaction when molar concentration of reactants are taken as unity. A+BC+D Muhammad Shahid Rate = k [A] [B] A-one Professors Academy of When the concentration of each reactant is taken as unity i.e. sciences Jauharabad [A] = 1 mol dm-3 [B] = 1 mol dm-3 so Rate = k [A] [B] = k × 1 × 1 = k It depends upon conditions especially on temperature. It remains the same throughout the reaction so it is a constant quantity. 25) 50 % of hypothetical first order reaction completes in one hours. The remaining 50 % needs more than one hour to complete? For first order reaction, the half-life is independent of initial concentration of reactants. In first hour 50 % reactants are converted into products. 50% of reactants are left behind. In next one hour 50% of remaining reactants are converted into products (i.e. 25 % of initial concentration) it means in two hours 75% of total initial concentration is converted into products. In next one hour 50 % of remaining 25 % is converted into products. It means in 3 hours 50 + 25 + 12.5 = 87.5 % of total initial amount is converted into products. It means it will take a very long time to convert all reactants into products.

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26) Unit of rate constant is different for different reactions? The unit of rate constant depends upon order of reaction. It is different for different reactions. Let us consider units of first order, second order and third order reactions. For first order reaction dx = k[A] dt dx =k dt[A]

For second order reaction dx = k[A]2 dt dx =k dt[A]2

For third order reaction dx = k[A]3 dt dx =k dt[A]3

mol dm3 mol

Sec ×

=k

dm3

1 =k Sec Sec −1 = k

mol

mol

dm3 mol Sec × ( 3 )2 dm

dm3 mol Sec × ( 3 )3 dm

1 Sec

mol

=k

=k

dm3

1 mol 2 ) dm3

Sec (

=k

=k

dm3 =k mol sec

(dm3 )2 =k sec(mol)2

dm3 mol−1 Sec −1 = k

dm6 mol−2 Sec −1 = k

So unit of “k” is Sec −1

So unit of “k” is dm3 mol−1 Sec −1

So unit of “k” is dm6 mol−2 Sec −1

27) The sum of Co-efficients of balanced equation is not always equal to order of reaction? The order of reaction is always measured experimentally. There are many reactions which do not complete in one step. In these reactions, the sum of co-efficients of balanced chemical equations is not equal to the order of reaction. The balance equation is equal to sum of all these steps. The order of reaction depends upon slow steps e.g. 2NO + 2H2

N2 + 2H2O

𝒅𝒙

= K [NO]2 [H2]1 Proposed mechanism is as follows Slow (i) 2NO + H2 N2 + H2O2 Fast (ii) H2O2 + H2 2H2O2 𝒅𝒕

Mechanism tells us that two molecules of NO and one molecule of H2 are involved in slow step so it is a third order reaction instead of fourth order. 28) The order of reaction is determined from the rate expression and the rate expression is determined from the experiments? There are many reactions which do not proceed in a single step. The overall order of reaction will depend upon slowest step. Thus order is experimentally determined. Then rate expression can be written and a mechanism is suggested for reaction e.g NO2 + CO

NO + CO2

The above expression shows that this reaction depends upon the concentration of NO2 & CO but the 𝒅𝒙 experimental rate expression is = K [NO2]2 𝒅𝒕 It is a second order reaction with respect to concentration of NO2.. It completes in two steps.

CHAPTER # 11

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(i) NO2 + NO2 (ii) NO3 + CO

106

Slow Fast

PROFESSOR EASY NOTES

NO3 + NO NO2 + CO2

29) Differentiate between rate and rate constant. Reaction rate is the change in concentration of a reactant or a product in a given amount of time. It changes with time. It depends upon the concentration of reactants. It is a variable quantity. Its units are mol dm-3s-1. ∆𝒄

𝒎𝒐𝒍 𝒅𝒎−𝟑

Rate = = = mol dm-3s-1. ∆𝒕 𝒔𝒆𝒄 The rate constant of a chemical reaction is the rate of reaction when the concentration of the reactants are unity. It does not changes with time. It is independent of the concentration of reactants. It is a constant quantity. Its units depends upon the order of reaction. Example = Consider a general reaction A + B  C + D Muhammad Shahid [A] = 1 mol dm-3 [B] = 1 mol dm-3 A-one Professors Academy of Its rate is given as sciences Jauharabad Rate = k [A][B], where ‘k’ is the rate constant. Rate = k [A] [B] = k × 1 × 1 = k 30) Briefly describe the following with examples;  Change of physical state of a catalyst at the end of a reaction. The physical state of a catalyst may be changed at the end of reaction. Example 1 = For the decomposition of KClO3, MnO2 is added in the form of granules. It is converted to fine powder at the end of reaction. Example 2 = In many cases shining surfaces of the solid catalyst become dull.  A very small amount of a catalyst may prove sufficient to carry out a reaction. Sometimes only a trace of a metal catalyst is used to affect very large amount of reactants. Example 1 = 1 mg of fine platinum powder can convert 2.5 dm3 of H2 and 1.25 dm3 of O2 to water. Example 2 = Thousands of dm3 of H2O2 can be decomposed in the presence of 1 g of colloidal platinum.  A finely divided catalyst may prove more effective. A catalyst is more effective in finely divided form. It is because increase in surface area increases the efficiency of a catalyst and rate of reaction increases. Example 1 = A big piece of Pt have much less catalyst activity than colloidal Pt. Example 2 = In the hydrogenation of vegetable oils finely divided nickel is used.  Equilibrium constant of a reversible reaction is not changed in the presence of a catalyst. A catalyst cannot affect the equilibrium constant of a reaction. It only decrease the time to reach equilibrium.

H+ Catalyst Example: CH3COOH + C2H5OH CH3COOC2H5 + H2O In this reaction, few drops of H2SO4 are used as catalyst and provides H+ ions. So equilibrium is established within hours. However, if catalyst is not used then it may take days to reach the equilibrium. In both cases, the equilibrium constant will be same at the same temperature.

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Long questions of previous board papers I. II. III. IV. V. VI. VII. VIII. IX. X. XI. XII.

Muhammad Shahid A-one Professors Academy of sciences Jauharabad

How will you find the order of reaction by half-life method? How Arrhenius equation explains the effect of temperature on the rate constant of a reaction? How Arrhenius equation explains the effect of temperature on the rate constant of a reaction? What is energy of activation? Discuss collision theory. What is enzyme catalysis? Give an example & give four characteristics of enzyme catalysis. What is catalysis? Write down its types and explain any one of them. Explain Half-life method for the Determination of order of Reaction. Give four characteristics of enzyme catalysis. Write a note on (i) Homogeneous catalysis (ii) Heterogeneous catalysis Write a detailed note on determination of rate of reaction by chemical method. How does Arrhenius equation help us to calculate the energy of activation of a reaction? What is catalyst? Write its three characteristics. Things to remember:

(2008) (2010) (2010) (2011) (2012) (2013) (2015) (2016) (2017) (2017) (2018) (2018)

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