Board Exam Genchem Reviewer.docx

CHEMISTRY REVIEWER

CHAPTER 1: INTRODUCTION CHEMISTRY  Central Science  Scientific discipline treats the composition, properties and transformation of matter  Study of structure of matter and changes that matter undergoes in natural processes and planned experiments I. INTRODUCTION TO MATTER MATTER  Physical material of the universe  Anything that occupies space and has mass A. STATES OF MATTER 1. Solid 2. Liquid

3. Gas

B. PROPERTIES OF MATTER  Sets of characteristics possessed by substances which distinguish them from one another B.1 PHYSICAL PROPERTIES  Properties can be measured or observed without changing basic identity of the substances

B.1.2 EXTRINSIC PROPERTIES  Sometimes called properties”

“accidental

B.2 CHEMICAL PROPERTIES  Properties that describe the way a substance may change or react to form other substances C. CHANGES IN MATTER 1. PHYSICAL CHANGE  Change wherein a substance changes in physical appearance but not in its basic identity 2. CHEMICAL CHANGE  Sometimes called “chemical reactions”  Change which transforms a substance into a chemically different substance Evidences of a Chemical Change a) Evolution of heat and light b) Evolution of gas c) Formation of precipitate d) Production of mechanical energy e) Production of electrical energy

B.1.1 INTRINSIC PROPERTIES  Properties that are characteristics of any sample of a substance regardless of the size or shape of the sample Example: 1. Density 2. Boiling Point 3. Melting Point 4. Freezing Point

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CHEMISTRY REVIEWER II. CLASSIFICATION OF MATTER A. PURE SUBSTANCE  Matter that has a fixed composition and distinct properties  Homogeneous material consisting of only one particular kind of matter 1. ELEMENTS  Substances cannot be decomposed into simpler substances by chemical means a. Metals  Solid at room temperature  Exhibit high electrical heat conductivity  Appear lustrous b. Non-metals  Non-conductors of electricity and heat 2. COMPOUNDS a. Acid  Substance produces or donates H+ ion  Increases concentration in definite proportions by mass b. Base  Substance produces or donate OHwhen dissolve in water c. Salt  Ionic compound formed by replacing one or more H+ of an acid by other cations

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B. MIXTURES  Combination of two or more substances  Each substance retains its own identity  Hence own properties 1. HOMOGENOUS  Material no differing parts can be distinguished even with a microscope  Only one 2. HETEROGENOUS  Material with differing parts  Consists of physically distinct parts  Each with different properties  Many 1. COLLOIDAL DISPERSION  Mixtures containing particles larger than normal solutes but small enough to invisible to the eye, even with a microscope  Appear in the field of an ultra microscope 2. COARSE SUSPENSION  Consists of particles appearing as huge molecules in dispersing medium  Particles may be visible to the eye  Settle fairly rapidly to the bottom of the container

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CHEMISTRY REVIEWER III. ENERGY  Capacity to do work or transfer heat EXOTHERMIC  Energy given off to the surroundings ENDOTHERMIC  Energy absorbed by a body A. TYPES OF ENERGY 1. POTENTIAL ENERGY  Stored up energy 2. KINETIC ENERGY  Energy due to motion B. FORMS OF ENERGY 1. ELECTRICAL ENERGY  Energy passage of electricity 2. RADIANT ENERGY  Energy electromagnetic radiation 3. CHEMICAL ENERGY  Energy possessed by a substance because of chemical state 4. NUCLEAR OR ATOMIC ENERGY  Energy with the manner in which the atom is constructed 5. LIGHT ENERGY 6. SOUND ENERGY IV. SCIENTIFIC METHOD OF SOLVING PROBLEMS 1. Identifying the problem 2. Gathering data 3. Formulating hypothesis 4. Testing a hypothesis 5. Interpreting experimental data 6. Drawing generalizations 7. Testing the generalization by applying it

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V. LAWS ON MATTER AND ENERGY 1. LAW OF CONSERVATION OF MASS  Mass neither created nor destroyed  During an ordinary chemical change 2. LAW OF CONSERVATION OF ENERGY  Energy is neither created nor destroyed  In any transformation but is transformed from one form to another 3. LAW OF DEFINITE PROPORTIONS (COMPOSITION)  States every sample of a given substance always contain same proportion by weight of all its constituent elements 4. LAW OF MULTIPLE PROPORTIONS  States when two elements are combined to form two or more different compounds  If the amount of one element is constant, mass of the other element in the different compounds are in a ratio of whole numbers Example: a) S + O2 → SO2 3.2 g 3.2 g b)

S + O2 → 3.2 g 4.8 g (ratio of O 2:3)

SO3

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CHEMISTRY REVIEWER VI. BRANCHES OF CHEMISTRY 1. PHYSICAL CHEMISTRY  Deals investigation of the laws and theories if chemistry  Among primary goal is the investigation of: a) Structure and transformation of matter and the interrelationships of energy b) subject matter, experimental techniques and instruments used are common to both physics and chemistry 2. ANALYTICAL CHEMISTRY  Concerned with separation, identification and composition of all kinds of matter  Two broad classifications aof analytical chemistry are: a) QUALITATIVE ANALYSIS  Involves separation and identification of individual components of materials  Answers the question, “ What is present?”

 Deals with commonly used synthetic substances such as plastics, drugs, dyes, explosives and detergents 4. INORGANIC CHEMISTRY  Covers the chemistry of all elements and their compounds with the exception of carbon and its compounds  Comprises investigation of those substances which are not organic such as non-living matter and minerals found in the earth’s crust 5. BIOCHEMISTRY  Includes the study of materials and processes occur in living organisms  These materials are largely organic compounds ( carbon containing compounds) 6. NUCLEAR CHEMISTRY  Deals with changes in nuclei of atoms and use of these changes  Especially in study of how substances react  Radioactive nuclei, both natural and man-made, decompose spontaneously

b) QUANTITATIVE ANALYSIS  Concerned with how much of each component is present 3. ORGANIC CHEMISTRY  Study of carbon containing compounds  Term “organic” derived from original belief these type of compounds were found only in living organisms

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CHEMISTRY REVIEWER VII. CHEMICAL CALCULATIONS A. Temperature vs. Heat HEAT  Form of energy flows spontaneously from a substance of higher temperature to one of lower temperature TEMPERATURE  Measure of intensity  Intensive quantity defines the direction and rate of heat flow  Hotness or coldness of a body Methods of Temperature: 1. RELATIVE TEMPERATURE  Scale is based on freezing and boiling point of water 2. ABSOLUTE TEMPERATURE  Based on quantity of kinetic energy that molecules have B. Mass vs. Weight MASS  Amount of matter found in a substance  Measure of a substance resistance to the change in the velocity WEIGHT  Refers to the force an object is attracted to the earth **Weight = (mass)(gravity)** VIII. DERIVED UNITS 1. VOLUME Mathematical Formulas: Cube V = s3 Rectangle V=LxWxH Sphere Cylinder

Cone

V=

4πr3 3

𝑉 = 𝜋𝑟 2 ℎ 𝑉=

𝜋𝑑 2 ℎ 4

𝑉=

𝜋𝑟 2 ℎ 3

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2. DENSITY  Amount of mass in a unit volume of a substance  Symbol: 𝜌  𝜌=

𝑚𝑎𝑠𝑠 𝑣𝑜𝑙𝑢𝑚𝑒

3. RELATIVE DENSITY  Ratio of mass of a substance to the mass of water that occupies same volume  𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝜌 =

𝑚𝑎𝑠𝑠 𝑠𝑎𝑚𝑝𝑙𝑒 𝑚𝑎𝑠𝑠 𝑤𝑎𝑡𝑒𝑟

 𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝜌 =

𝑤𝑒𝑖𝑔ℎ𝑡 𝑠𝑎𝑚𝑝𝑙𝑒 𝑤𝑒𝑖𝑔ℎ𝑡 𝑤𝑎𝑡𝑒𝑟

4. HEAT  Energy transferred as a result of difference  Symbol: Q SPECIFIC HEAT  Quantity of heat required to raise the temperature of one gram of a substance by 1.0⁰C CALORIE  Quantity of heat required to raise the temperature of one gram of water by 1.0⁰C LAW OF HEAT  Amount of heat body gains or losses depend on the mass and nature of the body and change in temperature of that body LAW OF HEAT EXCHANGE  Amount of heat lost by one body is equal to the amount of heat gained by another body

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CHEMISTRY REVIEWER CHAPTER 1 EXERCISES: 1. Indicate your answer by checking the appropriate blank: CHEMICAL SYSTEM PHYSICAL CHANGE A. Corrosion of aluminum metal _______________ B. Melting of sodium chloride _______________ C. Burning gasoline _______________ D. Pulverizing of rock salt _______________ E. Evaporation of alcohol _______________ F. Plants grow _______________ G. Fruits ripen _______________ H. Leaves decay _______________ I. Rusting of iron _______________ J. Water freezes _______________

CHEMICAL CHANGE ________________ ________________ ________________ ________________ ________________ ________________ ________________ ________________ ________________ ________________

2. Indicate your answer by checking the appropriate blank: CHEMICAL SYSTEM PURE SUBSTANCE HOMOGENEOUS A. Vinegar _____________ _____________ B. Polluted Air _____________ _____________ C. Gelatin _____________ _____________ D. Pebbles and Sand _____________ _____________ E. Wine _____________ _____________ F. Salt solution _____________ _____________ G. Unpolluted air _____________ _____________ H. Black coffee _____________ _____________ J. Mayonnaise _____________ _____________

HETEROGENEOUS ______________ ______________ ______________ ______________ ______________ ______________ ______________ ______________ ______________

3. Indicate the number of significant figures in each of the following: a) 73.00 - __________ f) 6.83471 b) 0.103444 - __________ g) 9.0 c) 28.326 - __________ h) 4.0003 x 1017 d) 12 x 103 - __________ i) 0.05930 e) 1,006,155 - __________ j) 5.138493

- __________ - __________ - __________ - __________ - __________

4. Round off the following quantities to the number of significant figures indicated in the parenthesis: a) 35,670.06 (3) - _______________ d) 667,999 (4) - _______________ b) 8.997 (2) - _______________ e) 0.045050 (2) - _______________ 3 c) 0.0000032456 (3) - _______________ f) 10.9546x10 (3) - _______________

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CHEMISTRY REVIEWER 5. The annual production of sodium hydroxide, NaOH, in the United States in 1991 was 24.39 billion pounds. How many grams of NaOH were produced in that year? Hint: 453.6 grams = 1 lb

6. In a certain part of a country, there is an average of 710 people per square mile and 0.72 telephones per person. What is the average number of telephones in an area of 5.0 km2?

7. The total amount of fresh water on earth is estimated to be 3.73x10 8 km3. What is this in cubic meters and in liters?

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CHEMISTRY REVIEWER 8. The predicted high temperature for New Delhi, India on a given day is 41.0⁰C. Is this temperature higher or lower than predicted daytime high of 103⁰F in Phoenix, Arizona for the same day?

9. Imagine that the temperature scale exists in degrees Washington (⁰W) and that on this scale; water boils at 145⁰W and freezes at 15.0⁰W. a) Convert a temperature of 35.0⁰W to ⁰C b) Convert a temperature of 24.0⁰F to ⁰W

10. Find the capacity, in L, of a box 0.60 m long, 10 cm wide and 50 mm deep.

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CHEMISTRY REVIEWER 11. A certain sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of the galena?

12. A) Calculate the density of mercury, Hg, if 1.00x102 grams occupies a volume of 7.36 cc. B) Calculate the mass of 65.0 cc of Hg.

13. The density of a sample of olive oil was determined in the following way. A flask was filled with ethanol, the mass of which was 8.02 g. Then, the flask was emptied and filled with olive oil. The mass of olive oil was 9.32 g. The density of ethanol is 0.789 g/cc. Calculate the density of olive oil.

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CHEMISTRY REVIEWER 14. A steel sphere has a radius of 1.58 in. IF the steel has a density of 7.88 g/cc, what is the mass of the steel in grams?

15. You are given a bottle that contains 2.36 mL of a yellow liquid. The total mass of the bottle and the liquid is 5.26 g. The empty bottle weighs 3.02 g. What is the density of the liquid? What is its relative density?

16. When 17.6 g of a metal is placed in a graduated vessel containing 10.0 mL of water, the water level rises to the 12.2 mL mark. What is the density of the metal?

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CHEMISTRY REVIEWER 17. A 200.0 g lead sinker at 110⁰C is dipped into 500.0 g of cold water at 10.0⁰C. If the resulting temperature of the mixture is 11.19⁰C, find the specific heat of lead in cal/g⁰C.

18. A cylindrical tube 15.0 cm in length is filled with ethanol. The mass of the ethanol needed to fill the tube is found to be 9.64 g. Calculate the inner diameter of the tube in cm. The density of ethanol is 0.780 g/mL.

19. A cylindrical settling tank is 6.0 ft deep and has a radius of 15.0 ft. What is the volume of the tank in liters?

20. A mass of 350 g of copper pellets at 100⁰C was mixed with 200 g of water at 22.4⁰C. The resultant temperature of the mixture was 33.2⁰C. Calculate the specific heat of copper.

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CHEMISTRY REVIEWER

CHAPTER 2: ATOMS, IONS, MOLECULES I. UNITS OF MATTER ATOM    

Basic building block of matter Smallest unit of element Cannot be further subdivided Capable of taking part in a chemical reaction

MOLECULE  Smallest unit of element or chemical compound capable of free existence  Combination of two or more atom II. THE ATOMIC THEORY OF MATTER Credit for the first atomic theory usually given to ancient Greeks but concept may have origin even earlier civilizations. Two theories prevailed among the Greeks: 1. ARISTOTLE (4TH Century B.C.)  Believed matter is continuous  Hypothetically can be divided into smaller particles 2. LEUCIPPUS AND DEMOCRITUS (5th Century B.C.)  Held subdivisions would ultimately yield atoms which could not be further divided  Word atom derived from Greek word atomos  Atomos – means uncut or indivisible Theories of Ancient Greeks  Based on abstract though not planned experiment  Approximately two thousand years, atomic theory remained speculation JOHN DALTON  English School Teacher  Published meaningful atomic theory  published period 1803-1807  Designed this theory to explain several experimental observations  Dalton’s theory considered landmark in Chemistry

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CHEMISTRY REVIEWER The essence of Dalton’s atomic theory of matter is summarized in the following postulates. 1. Elements composed of extremely small particles called atoms. 2. All atoms of given element are identical  Atoms of different elements are different and have different properties (including different masses) 3. Atoms of an element not changed into different types of atoms by chemical reaction  Atoms neither created nor destroyed in chemical reactions 4. Compounds formed when atoms of more than one element combine  A given compound always has same relative number and kind of atoms III. EVIDENCES OF SUBATOMIC PARTICLES In Greek and Dalton’s theory, atoms were regarded as the smallest possible components of matter. End of 19th century, it began to appear the atom itself might be composed of even smaller particles. This change in viewpoint was brought about largely by evidences from experiments in electricity. 1. CATHODE RAYS AND ELECTRONS  Mid-1800’s, scientist began to study electrical discharge through partially evacuated tubes JULIUS PLUCKER  Attempted to pass electric current through a vacuum  Led discovery of the cathode rays Method/Process: i. Two electrodes are sealed in a glass tube from which air is completely removed ii. When high voltage is impressed across these electrodes, rays stream from the negative electrode called the cathode. iii. Although rays themselves could not be seen, their movements could be detected because rays cause certain materials like glass to fluoresce or give off light. This phenomenon is called fluorescence. Properties of Cathode Rays a. They travel in straight lines away from the negative electrode unless acted upon by an outside force. b. They are negatively charged. c. They consist of particles of definite mass, 1/1837 times as much as the lightest atom known. d. The nature of the cathode ray is the same irrespective of: i. the material of which the cathode is made ii. the type of residual gas present in the evacuated tube iii. the kind of metal wires used to conduct current to the cathode iv. the materials used to produce current

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CHEMISTRY REVIEWER

GEORGE JOHNSTONE STONEY  Suggested to call electrons the particles from the cathode rays that were extensively studied in latter part of the 19th century Deflections of Cathode Rays  Since unlike charges attract, the streams of electrons that constitute cathode rays are attracted toward the positive plate when two oppositely charged plates are placed on either side of them.  The rays are deflected from their usual straight path in an electric field.  The degree of deflection varies directly with the size of the charge (e) and inversely with the mass of the particle (m).  The ratio of the charge to the mass (e/m) determines the extent to which electrons are deflected from a straight line path in an electric field. JOSEPH THOMSON (1897)  Determined value of e/m for electron  by studying the deflections of the cathode rays in electric and magnetic fields  Hailed as the discoverer of the first subatomic particle e/m = - 1.7588 x 108 coulombs/g ROBERT MILIKAN (1909)  Succeeded in measuring the charge of an electron  By performing an experiment known as Milikan’s Oil Drop Experiment e/m = - 1.6022 x 10-19 coulombs −𝟏. 𝟔𝟎𝟐𝟐 × 𝟏𝟎−𝟏𝟗 𝒄𝒐𝒖𝒍𝒐𝒎𝒃𝒔 𝒎= = 𝟗. 𝟏𝟎𝟗𝟑𝟗 × 𝟏𝟎−𝟐𝟖 𝒈𝒓𝒂𝒎𝒔 −𝟏. 𝟕𝟓𝟖𝟖 × 𝟏𝟎𝟖 𝒄𝒐𝒖𝒍𝒐𝒎𝒃𝒔/𝒈

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CHEMISTRY REVIEWER 2. PROTON  Positive ions which results when one or more electrons are removed from a neutral atom

Positive Ions  Formed in an electric charge when cathode rays rip electrons from atoms or molecules of the gas present in the tube  Positive ions move toward the negative electrode  If holes have been bored in this electrode, the positive ions pass through them EUGENE GOLDSTEIN (1886)  First to observe the positive ions called positive rays WILHELM WEIN (1898) AND J.J THOMSON (1906)  Studied deflections of positive rays in electric and magnetic field e/m = 9.5791 x 104 coulombs/g e = 1.6022 x 10-19 coulombs 𝟏. 𝟔𝟎𝟐𝟐 × 𝟏𝟎−𝟏𝟗 𝒄𝒐𝒖𝒍𝒐𝒎𝒃𝒔 𝒎= = 𝟏. 𝟔𝟕𝟐𝟔 × 𝟏𝟎−𝟐𝟒 𝒄𝒐𝒖𝒍𝒐𝒎𝒃𝒔 𝟗. 𝟓𝟕𝟗𝟏 × 𝟏𝟎𝟒 𝒄𝒐𝒖𝒍𝒐𝒎𝒃𝒔/𝒈

3. NEUTRON  Serves as the key which unlocks the energy of the nucleus  Produces radioisotopes which are used in medicine, industry, agriculture and research ERNEST RUTHERFORD (1920)  Postulated the existence of an uncharged particle JAMES CHADWICK  Hailed discoverer of the neutron  Able to calculate the mass of the neutron from the data on certain nuclear reactions in which neutron produced Mass = 1.6749 x 10-24 grams

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CHEMISTRY REVIEWER IV. RADIOACTIVITY WILHELM ROENTGEN (1895)  Found X- Ray  When cathode rays struck certain materials, new type of invisible ray was emitted. HENRI BECQUEREL  Studied phenomenon phosphorescence  Discovered radioactivity PHOSPHORESCENCE  Phenomenon wherein substances become luminous after exposure to sunlight RADIOACTIVITY  Spontaneous emission of radiation SLOWDOSKA AND PIERRE CURIE  Performed experiment to isolate the radioactive components of the mineral pitchblende ERNEST RUTHERFORD  Made further studies on the nature of radioactivity  His studies revealed three types of radiation Three Types of Radiation 1. ALPHA RADIATION, ∝  Much more massive than the β − particles  Have a positive rather than a negative charge  Has charge of +2 2. BETA RADIATION, 𝜷  high speed electrons  considered the radioactive equivalent of the cathode rays  has a charge of -1 3. GAMMA RADIATION, 𝜸  High energy radiation similar to the x-rays  Does not consist of particles

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CHEMISTRY REVIEWER V. THE NUCLEAR ATOM J.J THOMSON  Proposed a model for the structure of the atoms  Named his model “Plum-pudding” model  Proposed atom consisted of a uniform positive sphere in which electrons are embedded ERNEST RUTHERFORD, HANS GEIGER, ERNEST MARSDEN (1910)  Performed the Gold Foil Experiment to study the behavior of alpha particles as they pass through a thin gold fail ERNEST RUTHERFORD (1911)  Explained observations made from his Gold Foil Experiment on the behavior of alpha particles  Discover nucleus Observations: a. The large majority of the alpha particles passed through the foil b. Some were deflected from their straight-line path c. Few were recoiled back toward their source  Postulated most of the mass of the atom and all its positive charges, reside in a very small, extremely dense region which he called nucleus VI. THE MODERN VIEW OF THE ATOMIC STRUCTURE  The electron is given a charge of -1  The proton is given a charge of +1  Neutrons are uncharged  Atoms have equal number of electrons and protons  they have no net electric charge  Protons and neutrons reside together in the nucleus of the atom  Vast majority of the atom’s volume  Space in which electrons move  Electrons are attracted to the protons in the nucleus  By force that exist between particles of opposite electrical charge  Atoms are extremely small  Most of them have diameters between 1.0x10-10 and 5.0x10-10 m  More convenient unit is used to express atomic dimensions  A⁰ - known as Angstrom (1.0 A⁰ = 1.0x10-10 m) VII. ISOTOPES, ATOMIC NUMBERS AND MASS NUMBERS

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CHEMISTRY REVIEWER An Atom is Identified by Two Numbers: 1. ATOMIC NUMBER, Z  number of protons in the nucleus 2. MASS NUMBER, A  Total number of protons and neutrons in the nucleus  Neutrons – collectively called nucleons  The number of neutrons can be determined by getting the difference between the atomic number and the mass number  An atom of an element is designated by the chemical symbol for the element with the atomic number placed at the lower left and the mass number at the upper left where X is the chemical symbol of the element. Number of neutrons = A – Z Number of electrons = number of protons = Z ISOTOPES  Atoms of the same element containing different number of neutrons and different mass  Same atomic number but different mass number MASS SPECTROGRAPH  Used to determine the types of isotopes present in an element, exact atomic masses of these isotopes and relative amount of each isotope present ATOMIC WEIGHTS  ATOMIC MASS UNIT (amu) 12  Defined as 1/12 the mass of an atom of carbon, 6 C  The assignment of 12 u to the mass of the isotope of carbon is arbitrary  Most elements occur in nature as a mixture of isotopes. With very few exceptions, these mixtures have a constant composition

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CHEMISTRY REVIEWER EXAMPLE: The data obtained by the use of a mass spectrometer show that the element chlorine consists of 35 75.53% , 17 Cl atoms (mass =34.97 u), and 24.47% 37 Cl atoms (mass = 36.95 u). The atomic weight of 17 chlorine is the weighted average of the atomic masses of the natural isotopes. Solution: 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑤𝑒𝑖𝑔ℎ𝑡 = ∑ (𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛𝑎𝑙 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 × 𝑚𝑎𝑠𝑠) 𝐴𝑡𝑜𝑚𝑖𝑐 𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑐ℎ𝑙𝑜𝑟𝑖𝑛𝑒 = (0.7553 × 34.97 𝑢) + (0.2447 × 36.95 𝑢) 𝐴𝑡𝑜𝑚𝑖𝑐 𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑐ℎ𝑙𝑜𝑟𝑖𝑛𝑒 = 35.45 𝑢  The weighted average is found by multiplying the atomic masses of each isotope by its fractional abundance and adding the values obtained.  The fractional abundance is the decimal equivalent of the percent abundance CHAPTER 2 EXERCISES: 1. Complete the table: Atomic Atomic Symbol Number, Z Bi 83 Au K+1 P-3

Mass Number, A 209 197

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31

Number of Protons 79 19

Number of Neutrons

Number of Electrons

Charge

20 16

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CHEMISTRY REVIEWER 2. Chromium, Cr has the following isotopic masses and percentage abundances: MASS NUMBER ISOTOPIC MASS, amu PERCENTAGE ABUNDANCE, % 50 49.9461 4.35 51 51.9405 83.79 52 52.9407 9.50 54 53.9389 2.36 What is the atomic weight of chromium?

3. Determine the percentage abundance for the two naturally occurring isotopes of copper. The masses of the isotopes of copper are 63Cu, 62.9298 amu; 65Cu, 64.9278, respectively. The atomic weight of copper is 63.546 amu.

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CHEMISTRY REVIEWER 4. The element C consist of 92% of atoms with a mass of 28.0 a=u each, 5.00% of atoms with a mass of 29.0 u each, and 3.00% of atoms with a mass of 30.0 u each. What is the atomic weight of element X?

5. Silver has two naturally occurring isotopes, one of mass 106.91 amu and the other mass is 108.90 amu. Find the percentage abundances for these two isotopes. The atomic weight of silver is 107.87 amu.

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CHEMISTRY REVIEWER

CHAPTER 3: IMAGES OF THE INVISIBLE Historical Background SIR JOSEPH THOMSON (1899)  British physicist  Proposed an early model of the atom, called the plum pudding model  Described the atom as a sphere of positive charge containing a few negatively charged particles, called electrons, distributed in the sphere like raisins in a plum pudding. MAX PLANCK (1900)  Proposed Quantum Theory of Radiant Energy  He suggested that radiant energy could be absorbed or given off only in definite quantities called Quanta ALBERT EINSTEIN (1905)  Proposed that Planck’s quanta are discontinuous bits of energy which were later called Photons ERNEST RUTHERFORD (1911)  British physicist  Learned from experiments that the positive charge in an atom and most of the atom's mass must be concentrated in a small, central region, called the nucleus  He proposed that electrons carrying negative charge orbited the nucleus like planets orbiting in a solar system

NEILS BOHR (1913)  Danish physicist  Proposed a theory for the electronic structure of the hydrogen atom that explained the spectrum for this element  Discovered that the electrons in atoms had certain values of energy  Proposed that the energy of an electron was related to the distance at which the electron orbited the nucleus  Electrons, therefore, orbited only at certain distances from the nucleus, distances that corresponded to these allowed energies. He called these orbits quantized orbits, because they corresponded to these allowed energies.

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CHEMISTRY REVIEWER LOUIS DE BROGLIE (1924)  Proposed that electrons and other particles have wave properties  Showed that it was possible to assign a wavelength for an electron WERNER HEISENBERG (1926)  Proposed uncertainty principle ERWIN SCHRODINGER (1926)  Austrian physicist  Used de Broglie’s regulation to develop an equation that describes the electron terms of its wave character  The Schrodinger Equation – basis of Wave Mechanics  Proposed that electrons do not orbit the nucleus but behave more like waves traveling at certain distances and with certain energies around the nucleus  This model proved to be the most accurate  Physicists no longer try to find an electron's path and position in an atom. Instead, they use equations describing the electron wave to find the region of space in which the electron is most likely to be found

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CHEMISTRY REVIEWER ATOMIC ORBITAL  Most probable place where a pair of electrons could be found  Different types of orbital have been described and studied QUANTUM NUMBERS  In wave mechanics, the electron distribution of an atom containing a number of electron is divided into shells.  The shells, in turn, are thought to consist of one or more subshells, which the electrons occupy.  Each electron of an atom is identified by a combination of four quantum numbers, which loosely indicate the following: mainshell, subshell, orbital and electron (spin) A. PRINCIPAL QUANTUM NUMBER (n)  Identifies the shell or main energy level to which the electron belongs  These shells are regions where probability of finding an electron is high  Values of n are positive integers n = 1, 2, 3, 4….  The larger the value of n, the farther the shell is from the nucleus  N-value signifies the average distance of the electron from the nucleus B. AZIMUTHAL QUANTUM NUMBER (ℓ)  Each shell consists of one or more subshells or sub energy levels  The number of subshells in the principal energy shell is equal to the value of n  Each subshell is assigned a subsidiary quantum number, ℓ  The values of ℓ for the sublevels of a shell are determined by the shell’s value of n  A letter is used to represent each value of ℓ FORMULA: Summary/Guide: n ℓ Spectroscopic notation

ℓ = (n – 1)

1 0

2 1

3 2

4 3

5 4

6 5….

s

p

d

f

g

h

 The first notations are initial letters of adjectives formerly used to identify spectral lines: sharp, principal, diffuse and fundamental EXAMPLE: 1. The subshell with n= 1 and ℓ = 0 is called the 1s subshell 2. The subshell with n= 3 and ℓ = 2 is called the 3d subshell

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CHEMISTRY REVIEWER C. MAGNETIC ORBITAL QUANTUM NUMBER (mℓ)  Each subshell consists of one or more orbital FORMULA: No. of orbital per subshell = 2ℓ + 1 Summary/Guide: n 1 2 3 4 ℓ 0 1 2 3 Spectroscopic s p d f Notation Number of 1 3 5 7 Orbital

5 4

6 5….

g

h

9

11…

 Each orbital within a given subshell is identified by a “magnetic quantum number”, mℓ  Notice that the values of ℓ and mℓ are derived from n EXAMPLE: 1. For ℓ=0, only permitted value of mℓ is 0 (1s - orbital) 2. For ℓ=2, mℓ values are +2 , +1 , 0, -1 , -2 (5d – orbital) D. MAGNETIC SPIN QUANTUM NUMBER, ms  Fourth quantum number  An electron in an orbital has properties that can be explained by assuming that it spins on its axis  Only two directions are possible, hence only two possible spin values 

1

1

ms + 2 and − 2

ORBITAL FILLING AND HUND’S RULE OF MAXIMUM MULTIPLICITY In considering the way in which the electron in an atom distribute themselves various orbital, some principles are used as guides: 1. An electron tends to go into orbital of lowest energy levels  Electrons tend to occupy orbital as close to the nucleus as possible because of the attraction between the nuclear charge and the electrons  There is a greater tendency of an electron to occupy 1s-orbital than 2sorbital and the 2s-orbital more than the 2p-orbital  Thus single electron og hydrogen will locate itself in the 1s-orbital 2. At any one energy level, electrons tend to occupy separate orbital  This is in accordance with Hund’s Rule of Maximum Multiplicity  Hund’s Rule follows Pauli’s Exclusion Principle  Thus if there are a number of electrons in an atom with parallel spins, these electrons must be in different orbital and that only two electrons with opposite spins can be stable in an orbital

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CHEMISTRY REVIEWER

HUND’S RULE OF MULTIPLICITY  States that electrons tend to avoid being in the same orbital and a pair of electrons occupying a pair of equivalent orbital have their spins in parallel in a state of lower energy PAULI’S ESCLUSION PRINCIPLE  States that “no two electrons in the same atom can have identical sets of 4 quantum numbers” ELECTRON CONFIGURATION  Electrons are arranged in an atom DERIVATION OF ELECTRON CONFIGURATION AUFBAU PRINCIPLE  A method that is first suggested by Wolfgang Paula  Aufbau means “building up” in German METHOD:  The electron that is added in going from one element to the next is called differentiating electron  It makes the configuration of an atom different from that of the atom that precedes it  The differentiating electron is added in each step to the orbital of lowest energy available to it  It is often useful to employ a mnemonic device in the derivation of the electron configuration of any given atom EXAMPLE: i. Start with hydrogen atom which has one electron in the 1s orbital ii. By adding one electron, we get the configuration of the atom of the next element, helium, which is 1s2 iii. In this manner, we go from one element until we derive the configuration of the atom we desire. DIAMAGNETIC AND PARAMAGNETIC SUBSTANCES The number of unpaired electrons in an atom, ion or molecules can be determined by magnetic measurements. PARAMAGNETIC SUBSTANCES  Drawn (or attached) into a magnetic field  Substances that contain unpaired electrons

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CHEMISTRY REVIEWER DIAMAGNETIC SUBSTANCES  Weakly repelled by a magnetic field  A material is diamagnetic if all the electrons are paired  Diamagnetism is a property of all matter but it is obscured by the stronger paramagnetic effect if unpaired electrons are also present

PNEMONICS IN WRITING ELECTRONIC CONFIGURATION NOTE: 1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

Maximum number of content every letter 2s For:

PNEMONICS IN WRITING ELECTRONIC CONFIGURATION

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f



s=2 p=6 d = 10 f = 14 gap by four

PNEMONICS IN WRITING ELECTRONIC CONFIGURATION TREND 1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

7p

3d

4d

4f

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5d

5f

6d

6f

7d

7f

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CHEMISTRY REVIEWER CHAPTER 3 FIGURES

Figure 1: Quantum Description of Electrons

Take Note the Parts

Figure 2: Example Electronic Configuration

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CHEMISTRY REVIEWER CHAPTER 3 EXERCISES: 1. Write the total number of electrons that can have the given quantum number/s pr designation in an atom. a) n = 3 b) 3d c) 4f d) n = 6 , ℓ = 1 e) 2d f) n = 3, ℓ = 2, mℓ = -2 g) n = 4, ℓ = 3, mℓ = -3, ms = +1/2 h) 1p 2. List the quantum number of each electron in a Ne atom.

3. Determine the maximum number of electrons that can be housed in the L-shell and in the N-shell.

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CHEMISTRY REVIEWER 4. How many unpaired electrons are there in a gaseous form of: a) Mg b) Br c) Li d) Ar

5. Given the following electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 40Zr : a) How many electrons are there in the outermost energy level? b) How many empty orbital does Zr have? c) How many half-filled orbital does Zr have? d) How many electrons have an ℓ=1? e) How many electrons have an mℓ= -1? f) Is Zr paramagnetic or diamagnetic?

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4p6

5s2

4d10

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CHEMISTRY REVIEWER

CHAPTER 4: CHEMICAL BONDS: THE TIES THAT BIND I. INTRODUCTION While the atomic theory was being developed, various ideas were also entertained about the combination of atoms that lead to chemical compounds. CHEMICAL BONDS  Forces that hold two bonding atoms together LEWIS STRUCTURE  Electron-dot formula wherein the valence electrons are shown together with the atomic symbol ELECTRONS  Play key role in chemical bonding A. LEWIS THEORY 1. Electrons, especially those of the outermost (valence) electronic shell, play a fundamental role in chemical bonding. 2. In some cases, chemical bonding results from the transfer of one or more electrons from one atom to another.  This leads to the formation of positive and negative ions and a bond type known as ionic. 3. In other cases, chemical bonding results from the mutual sharing of electrons between atoms.  This leads to the formation of molecules having a bond type called covalent. 4. The transfer or sharing of electrons occurs to the extent that each atom involved acquires an especially stable electron configuration. Often, this configuration is that of a noble (inert) gas, that is, involving eight outer-shell electrons, an octet. OCTET RULE  states that elements gain or lose electrons to attain an electron configuration of the nearest noble gas  Noble gases have complete outer electron shells, which make them very stable  Other elements also seek stability, which governs their reactivity and bonding behavior  Halogens are one electron away from filled energy levels, so they are very reactive Why Do Elements Follow the Octet Rule?  Atoms follow the octet rule because they always seek the most stable electron configuration  Following the octet rule results in completely filled s- and p- orbitals in an atom's outermost energy level

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CHEMISTRY REVIEWER 

Low atomic weight elements (the first twenty elements) are most likely to adhere to the octet rule

B. LEWIS SYMBOL  The Lewis Symbol of an element consists of the common chemical symbols, surrounded by a number of dots.  Leads to a more accurate prediction of chemical formulas when the number of unpaired dots is maximized before dots are shown as pairs C. LEWIS ELECTRON DOT DIAGRAM  Lewis electron dot diagrams may be drawn to help account for the electrons participating in a chemical bond between elements  A Lewis diagram counts the valence electrons  Electrons shared in a covalent bond are counted twice  For the octet rule, there should be eight electrons accounted for around each atom D. IONIC BONDING  Transfer of electrons from one atom to another  Metal and Non-metal

 Gain or lose electrons and subsequently become electrically charged  Atom that gains an electron is known as a negative ion  Atom that loses an electron is known as a positive ion Example: Na + Cl → NaCl For Sodium (Na): Step 1: Electron Configuration Atomic Number

Electronic Configuration Symbol

Atomic/Molecular Weight Na: 1s2 2s2 2p6 3s1 Step 2: Note last configuration Na: 1s2 2s2 2p6 3s1

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CHEMISTRY REVIEWER Step 3: Note number of electrons in the last shell and its similar shell (if applicable) is the number of dot/s in an element Na: 3s 1 For Chlorine (Cl): Step 1: Cl: 1s2 2s2 2p6 3s2 3p5 Step 2: Cl: 1s2 2s2 2p6 3s2 3p5 Step 3: Cl: 1s2 2s2 2p6 3s2 3p5  Account also 3s2 because same sila nung shell ng last na configuration  Review: 3p5 Shell  3s2 and 3p5 Same na “3” ang shell Step 4: Add the number of electrons sa magkaparehas ang shell 2 + 5 =7 Step 5: Conclusion Means 7 ang electron/dot ng Chlorine (Cl) PLOTTING: Na + Cl → NaCl Let dot of Na = x Let dot of Cl =

Na

+

Cl

→

[ Na ] [ Cl ]

1. An ionic bond results from the transfer of electrons between a metal and a non-metal. In this transfer, the metal atom becomes a positively charged ion (cation) and the non-metal, a negatively charged ion (anion). 2. The non-metal gains a sufficient number of electrons to produce an anion with a noble gas configuration. 3. A formula unit of an ionic compound is the smallest collection of ion electrically neutral.  The formula unit is obtained automatically when structure is written.

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CHEMISTRY REVIEWER D. COVALENT BONDING  Sharing of electrons between two bonded atoms  Non-metal and Non-metal  Covalent bonds form between atoms that have a tendency to share valence electrons to complete their outer shell TYPES OF COVALENT BONDS 1. Non-polar Covalent Bonds  Equal sharing of electron clouds between two bond atoms 2. Polar Covalent Bonds  Unequal sharing of electron clouds between two bonded atoms 3. Coordinate Covalent Bonds  Also known as Dative or Abnormal covalent bond

CHAPTER 4 EXERCISES: 1. Write the Lewis structure for the following ionic compounds: A) BaO C) K2S

B) MgCl2

2. Write the Lewis structure corresponding to the different chemical formulas: A) H2 E) N2

B) Cl2

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F) H2O

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CHEMISTRY REVIEWER C) HCl

G) NH3

D) O2

H) HCN

CHAPTER 5: CHEMICAL BONDS: THE LANGUAGE OF CHEMISTRY I. DEFINITION OF TERMS 1. CHEMICAL SYMBOL  A chemical abbreviation used to denote the elements  Consist of one or two letters  First letter capitalized  Often derived from the English name or form the foreign name of the element 2. CHEMICAL FORMULA  Uses a combination of symbols to represent a compound  Indicates the elements present and relative numbers of atoms of each element in the compound 3. FORMULA UNIT  Smallest collection of atoms in which the formula of a compound can be based 4. CHEMICAL NOMENCLATURE  From latin words “nomen” meaning name and “calare” meaning to call  System of writing formulas and naming compounds  Standards of nomenclature are established by the IUPAC IUPAC – International Union of Pure and Applied Chemistry 5. CHEMICAL EQUATION  Shorthand way of expressing a chemical reaction  Formulas of reactants are written on the left side and formulas of products are written on the right side 6. SUBSCRIPT  Indicates the relative number of atoms of an element in a compound

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CHEMISTRY REVIEWER 7. OXIDATION STATE  Relates the number of electrons an atom losses, gains or shares in combining with other atoms to form molecules or polyatomic ions  Combining capacity of an atom but it specifies its change  Assignment of oxidation states requires applying a set of rules  once assigned oxidation states are useful in naming compounds and balancing equations 8. POLYATOMIC ION OR RADICAL  Consists of two or more atoms taken as one unit with a corresponding charge 9. COEFFICIENTS OF EQUATION  Used to balance chemical equation 10. VALENCE ELECTRON  Denotes number of electrons in the outermost energy level  Indicates the combining power of an atom in a a compound

II. CHEMICAL NOMENCLATURE A. POSITIVE IONS/CATIONS 1. MONOVALENT CATIONS  Cations formed from metal atoms have the same name as the metal  Metal name + ion  Only one charge NOTE:  Without charge, atom form  With charge, ion form Example: Na+ Zn+ Al3+

Sodium ion Zinc ion Aluminum ion

2. POLYVALENT CATIONS i. ROMAN NUMERICAL METHOD  If a metal can form cations of differing charges, the positive charge is given by a roman numeral in parentheses following the name of the metal  Metal name (numerical) + ion

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CHEMISTRY REVIEWER Example: Fe2+ Fe3+ Cu+ Cu2+

Iron (II) ion Iron (III) ion Copper (I) ion Copper (II) ion

ii. CLASSICAL/OLDER METHOD  Used in distinguishing between two differently charged ions of metal is to apply the ending –ous or –ic representing lower and higher charged ions  Trivial name + ion NOTE:  Lower charge - ous  Higher charge – ic Example: Fe2+ Fe3+ Cu+ Cu2+

Ferrous ion Ferric Ion Cuprous ion Cupric ion

Pb2+ Pb4+

Plumbous ion Plumbic ion

iii. SPECIAL IONS (CATIONS THAT ARE FROM NONMETALS)  Cations formed from nonmetal atoms  Have names that end in –ium Example: NHO4+ NHO3+

ammonium ion hydronium ion

B. NEGATIVE IONS/ANIONS 1. MONOATOMIC/MONOVALENT ANIONS  Have names formed by dropping the ending of the name of the element and adding the ending –ide  Ending of the name changed to ide + ion Example: ClFS2O2-

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chloride ion fluoride ion sulfide ion oxide ion

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CHEMISTRY REVIEWER 2. OXYANIONS(RADICALS)  Polyatomic anions containing oxygen have names ending in –ate or –ite  Group 7 in the periodic table have usually has hypo- and perNOTE:    

Normal/Standard/Stable Form Less oxygen than stable form Less oxygen than –ite ion Greater than stable form Example: SO42SO32ClO3ClO22ClOClO4-

-ate ion -ite ion - hypo____ite ion - per_____ate ion

sulfate ion sulfite ion chlorate ion chlorite ion hypochlorite ion perchlorate ion

3. HYDROGEN ANIONS  Anions derived by adding H+ to an oxyanion are named by adding ad a prefix the word hydrogen or dihydrogen NOTE:  bi_____ate ion  bi_____ite ion  Hydrogen + anion name Example: HCO3-

Bicarbonate ion Hydrogen carbon ion

HSO4-

Bisulfate ion Hydrogen sulfate ion

H2PO4

dihydrogen phosphite

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CHEMISTRY REVIEWER C. IONIC COMPOUNDS  Metal + Non-metal  Cation + Anion Example: NaCl BaF2 K2CO3 Al2O3 Mg3N2

Sodium chloride Barium fluoride Potassium carbonate Aluminum oxide Magnesium nitride

D. BINARY/COVALENT COMPOUNDS  Non-metal + Non-metal  Number of atoms indicated by prefix mono, di, tri, tetra… NOTE:  Second element has the prefix Example: Cl2O5 Chlorine pentaoxide P2O5 Phosphorus pentaoxide N2O4 Nitrogen tetraoxide SO2 Sulfur dioxide CO2 Carbon dioxide CO Carbon monoxide E. ACIDS  With Hydrogen 1. -IDE ending to HYDRO____IC ACID Example: ANION __________ide

ACID hydro ______ic acid

Cl-

chloride

HCl

Hydrochloric acid

I-

iodide

HI

Hydroiodic acid

2. –ATE ending to _____IC ACID Example: ANION __________ate ClO4- Perchlorate SO42- Sulfate

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ACID ________ic acid HClO4 Perchloric acid H2SO4 Sulfuric acid Page 39

CHEMISTRY REVIEWER 3. –ITE ending to _____OUS ACID Example: ANION __________ite ClO2- chlorite SO32- Sulfite F. BASES  With OH Metal/cation name + hydroxide Example: CATION HYDROXIDE 3+ Al OHK+ OHBa2+ OH-

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ACID ________ous acid HClO2 chlorous acid H2SO3 Sulfurous acid

COMPOUND Aluminum hydroxide Potassium hydroxide Barium hydroxide

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CHEMISTRY REVIEWER

III. CHEMICAL EQUATION 𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 Interacting substances

→

𝑃𝑟𝑜𝑑𝑢𝑐𝑡𝑠 New Substances formed

 An arrow is drawn toward the product to indicate that the reactants are yielding the products  A chemical equation must be balanced IV. TYPES OF CHEMICAL REACTIONS Any reaction or chemical change may be classified as one of the following basic reaction types: 1. DIRECT COMBINATION OR SYNTHESIS  Occurs when two or more substances combine to form a more complex substance 𝐴 + 𝐵 → 𝐴𝐵 a) Metal and nonmetal to form binary compounds or metallic oxide 2𝑀𝑔 + 𝑂2 → 2𝑀𝑔𝑂 2𝑁𝑎 +

𝐶𝑙2 →

2𝑁𝑎𝐶𝑙

b) Nonmetal and oxygen form nonmetallic oxide 𝑆 + 𝑂2 → 𝑆𝑂2 𝐶 +

𝑂2

→

𝐶𝑂2

c) Metallic oxide and water form metallic hydroxide (basic anhydride) 𝑀𝑔𝑂 + 𝐻2 𝑂 → 𝑀𝑔(𝑂𝐻)2 𝐶𝑎𝑂 +

𝐻2 𝑂

→

𝐶𝑎(𝑂𝐻)2

d) Nonmetallic oxide and water form acid oxide (acid anhydride) 𝑆𝑂3 + 𝐻2 𝑂 → 𝐻2 𝑆𝑂4 𝑆𝑂2

+

𝐻2 𝑂

→

𝐻2 𝑆𝑂3

e) Metallic oxide and nonmetallic oxide form salts 𝑀𝑔𝑂 + 𝑆𝑂3 → 𝑀𝑔𝑆𝑂4 𝑁𝑎2 𝑂 +

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𝐶𝑂2

→

𝑁𝑎2 𝐶𝑂3

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CHEMISTRY REVIEWER 2. DECOMPOSITION OR ANALYSIS  Breaking down of a substance 𝐴𝐵 → 𝐴 + 𝐵 a) Heating of metallic carbonates form carbon dioxide and metal oxides 𝐶𝑎𝐶𝑂3 → 𝐶𝑂2 + 𝐶𝑎𝑂 𝑁𝑎2 𝐶𝑂3

→

𝐶𝑂2 +

𝑁𝑎2 𝑂

b) Electrolysis 2𝐻2 𝑂

→

𝑁𝑎𝐶𝑙

2𝐻2

→

2𝑁𝑎

+

𝑂2

+

𝐶𝑙2

c) Heating of metallic oxides/hydroxides form metal oxides and water 𝑀𝑔(𝑂𝐻)2 → 𝑀𝑔𝑂 + 𝐻2 𝑂 2𝐿𝑖𝑂𝐻

→

𝐿𝑖2 𝑂 +

𝐻2 𝑂

d) Heating acids form nonmetallic oxides and water 𝐻2 𝑆𝑂4 → 𝑆𝑂3 + 𝐻2 𝑂 e) Heating oxides form metals and oxygen 2𝐻𝑔𝑂 → 2𝐻𝑔 f)

+

𝑂2

Heating metallic chlorates form metallic chloride and oxygen 2𝐾𝐶𝑙𝑂3 → 𝐾𝐶𝑙 + 3𝑂2

3. SINGLE DISPLACEMENT  Reaction wherein an element or ion displaces another element or ion in a compound  Weaker metals will be replaced by active metals 𝐴𝐵 + 𝐶𝑚𝑒𝑡𝑎𝑙 → 𝐶𝐵 + 𝐴 𝐴𝐵 + 𝐶𝑛𝑜𝑛𝑚𝑒𝑡𝑎𝑙 → 𝐴𝐶 + 𝐵 Example: 𝑁𝑎𝑂𝐻 𝐴𝑔𝑁𝑂3

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+ +

𝐿𝑖 𝐾

→ →

𝐿𝑖𝑂𝐻 𝐾𝑁𝑂3

𝑁𝑎 +

𝐴𝑔

Page 42

CHEMISTRY REVIEWER 4. DOUBLE DECOMPOSITION OR METATHESIS  Reaction wherein compounds interchange ions or radicals  Replacement of the cations or substitution of elements or ions  Compounds exchange positive and negative parts  Result in the formation of different compounds 𝐴𝐵 + 𝐶𝐷

→

𝐴𝐷

+

𝐾𝑂𝐻

→

𝐾𝑁𝑂3

𝑀𝑔𝐶𝑙2

→

𝐶𝐵

Example: 𝑁𝑎𝑁𝑂3 + 𝐵𝑎𝑂

+

+

𝐵𝑎𝐶𝑙2

𝑁𝑎𝑂𝐻

+

𝑀𝑔𝑂

5. NEUTRALIZATION  Reaction between an acid and a base or an oxide to form salt and water a) An acid and a base 𝐻2 𝑆𝑂4 + 2𝑁𝑎𝑂𝐻 → 𝐻2 𝑂 + 𝑁𝑎2 𝑆𝑂4 b) Metal oxide and an acid 𝑀𝑔𝑂 +

2𝐻𝐶𝑙

→

𝑀𝑔𝐶𝑙2

c) Nonmetal oxide and a base 𝑆𝑂3 + 𝐿𝑖𝑂𝐻

→

𝐿𝑖𝑆𝑂4

+

𝐻2 𝑂

𝑀𝑔

+

𝐶𝑂3

d) Basic oxide and an acid oxide 𝑀𝑔𝑂 + 𝐶𝑂2 e) Ammonia and an acid 𝑁𝐻3

+

𝐻𝐶𝑙

→

→

+

𝐻2 𝑂

𝑁𝐻4 𝐶𝑙

CHAPTER 5 EXERCISES: 1. Write the correct formula for the following: A) Potassium hydride - _____________________________________ B) Strontium hydroxide

- _____________________________________

C) Ferric bromide

- _____________________________________

D) Ammonium dichromate

- _____________________________________

E) Iron (III) oxide

- _____________________________________

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CHEMISTRY REVIEWER F) Chromium (II) chloride

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G) Hydrosulfuric acid

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H) Zinc acetate

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I) Mercuric nitrate

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J) Bromine pentafluoride

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K) Calcium phosphate

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L) Lead (IV) nitride

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M) Potassium hydrogen sulfate

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2. Name the following compounds: A) SrO - __________________________________________________________ B) MgCl2

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C) K2CrO4

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D) Cs2SO4

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E) Mg(NO3)2

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F) Cr2O3

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G) TiCl4

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H) ZnS

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I) Ca(HCO3)2

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J) Co(ClO4)2

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K) K2HPO4

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L) NH4I

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M) Mn(OH)2

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CHEMISTRY REVIEWER

3. Balance the following equations and classify each type: A) 𝐾𝐶𝑙𝑂3 → 𝐾𝐶𝑙 + 𝑂3 B) 𝑁𝑎𝑂𝐻 + 𝐻3 𝑃𝑂4 → 𝑁𝑎3 𝑃𝑂4 + 𝐻2 𝑂 C) 𝑍𝑛 + 𝐴𝑔𝑁𝑂3 → 𝑍𝑛(𝑁𝑂3 )2 + 𝐴𝑔 D) 𝐶𝑟 + 𝐶𝑙2 → 𝐶𝑟𝐶𝑙3 + 𝐶𝑟(𝐶𝑙2 )3 E) 𝐴𝑙 + 𝐻𝐶𝑙 → 𝐴𝑙𝐶𝑙3 + 𝐻2 4. Write balanced equations for the following reactions: A) Barium reacts with hydrogen to yield barium hydride. B) Ammonium chloride reacts with lead (II) nitrate to yield ammonium nitrate and lead (II) chloride. C) Mercury (II) oxide when burned forms mercury and oxygen. D) Hydrofluoric acid reacts with silicon dioxide to yield silicon tetrafluoride and water. 5. Balance each of the following equations after predicting the products and classify each. A) Copper + Silver nitrate → B) Magnesium + water → C) Zinc + Hydrochloric acid → D) Calcium hydroxide + Sulfuric acid → E) Potassium + Water → F) Copper(II) carbonate + heat →

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CHEMISTRY REVIEWER

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