Int Chem Chap 3

3 CHEMICAL FORMULAE AND EQUATIONS

Relative atomic mass, Ar 1. The mass of an atom is very small and therefore difficult to be determined accurately. 2. However, we can determine the mass of an atom relative to a standard atom. 3. The atomic mass of an atom found by comparing it with a standard atom, is called relative atomic mass, or Ar, for short.

Relative atomic mass, Ar Using Hydrogen 4. This method had several weaknesses. (i) Relative masses of certain elements cannot be determined because these elements do not combine readily with hydrogen or are, unable to displace hydrogen; (ii) It was found that the relative masses of some elements were not accurate. (iii) Since hydrogen exists as a gas at room temperature, it is difficult to determine its mass accurately.

Relative atomic mass, Ar 5. (a) Oxygen was then used as a standard to compare the masses of atoms, (b) However, problems arose when the existence of three isotopes of oxygen were discovered.

Relative atomic mass based on carbon-12 scale 1. Carbon-12 was chosen for a few reasons. (a) Carbon-12 is used as a reference standard in the mass spectrometer. (b) Many elements can combine with carbon-12. (c) Carbon-12 exists as a solid at room temperature and thus can be handled easily. (d) Carbon-12 is the most abundant carbon isotope, occurring about 98.89%. (e) The table of relative atomic masses based on carbon-12 is in close agreement with the tables based on oxygen.

Relative atomic mass based on carbon-12 scale 2. Carbon-12 isotope is assigned a mass of exactly 12 units.  Relative atomic mass of an element = average mass of oneatomof theelement the

1 of the mass of an atomof carbon − 12 12

Relative atomic mass based on carbon-12 scale 3. For example, the average mass of one nitrogen atom is 14 times larger than 1/12 of the mass of carbon-12 atom. Therefore, the relative atomic mass of nitrogen is 14. [Relative atomic mass is not the actual mass of the atom. It is only a comparison value. Therefore, it has no unit.]

Relative molecular mass, Mr

=

1.The relative molecular mass of a substance is the average mass of a molecule of the substance when compared with of the mass of an atom of carbon-12. Relative molecular mass of a thesubstance average mass of one moleculeof the subs tan ce

1 of the mass of an atomof carbon − 12 12

Relative molecular mass, Mr 2. The relative molecular mass of a substance can be calculated by adding up the relative atomic masses of all the atoms present in a molecule of the substance. 3. For this reason, it is essential to know the molecular formula of the substance first.

Relative molecular mass, Mr 

4. Table 3.1 shows how to calculate the relative molecular masses of some substances.

[Relative atomic mass: H,1; C,12; N,14; O,16]

Substance

Molecular formula

Hydrogen gas

H2

Relative molecular mass,Mr 2(1)=2

Oxygen gas

O2

2(16)=32

Water

H2 O

2(1)+16=18

Ammonia

NH3

14+3(1)=17

Propane

C3H8

3(12)+8(1)=44

Ethanol

C2H5OH

2(12)+5(1)+16+1 =46

Relative molecular mass, Mr 

5. The term `relative molecular mass' can only be used for substances that are made up of molecules. For ionic compounds, the term `relative formula mass' is used instead. 

Relative molecular mass, Mr 

6. Table 3.2 shows the calculation of the relative formula masses of some ionic compounds.

[Relative atomic mass: H,1; O,16; Na,23; Al,27; S,32; Cl,35.5; K,39; Cu,64]

Substance

Molecular formula

Relative formula mass, Mr

Sodium chloride

NaCl

23+35.5=58.5

Potassium oxide

K2O

2(39)+16=94

Aluminium sulphate

Al2(SO4)3

2(27)+3[32+4(16 )]=342

Hydrated copper(II) sulphate

CuSO4.5H O 2

64+32+4(16)+5[ 2(1)+16]=250

3.2 The Mole and the Number of Particles 

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1. In daily life, we count objects in units such as dozen and pair. In chemistry, we measure substances in, `mole'. The symbol of mole is mol. 2. One mole is an amount of substance that contains as many particles as the number of atoms in exactly 12 g of carbon-12, which is 6.02 x 1023 particles. 3. The number of particles per mole (6.02 x 1023 mol-') is determined experimentally and is known as the Avogadro constant or Avogadro number. The Avogadro constant (NA) is defined as the number of particles in one mole of a substance.

3.2 The Mole and the Number of Particles  

 4. In simple words, 1 mole of substances contains 6.02 x 1023 particles, 2 moles of substance contain 2 x 6.02 x 1023 particles, and soon.

Atom, molecules, ions

3.2 The Mole and the Number of Particles 5. We can convert the number of moles of any substance to the number of particles in it and vice versa using the following relationship. xN A  → Numberof particles Numberof mol ←  

÷ NA

3.3 The Mole and the Mass of Substances 

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1. The molar mass of a substance is the mass of one mole of the substance. It has a unit of grams per mole (g mol-1). 2. In other words, the molar mass of a substance is the mass of 6.02 x 1023 particles of the substance.

3.3 The Mole and the Mass of Substances 3. Based on the Table 3.4, we an measure 1 mole or magnesium simply by weighing 24g of magnesium. This amount of magnesium contains 6.02 x 1023 magnesium atoms. [relative atomic mass: 

H,1;He,4; C,12;Na,23;Mg,24;Cl,35.5; Zn,65;Br,80]

Substance

Relative mass

Molar mass (g mol-1 )

Magnesium, Mg

Ar= 24

24

Helium, He

Ar= 24

4

Hydrogen gas, H2

Mr= 2(1)=2

2

Methane, CH4

Mr= 12+4(1)=16

16

Sodium Chloride, NaCl

Fr = 23+35.5=58. 5 Fr = 65+2(80)=22 5

58.5

Zinc bromide, ZnBr2

225

3.3 The Mole and the Mass of Substances 

4. We can easily calculate the mass of any number of moles of substance or vice versa using the following relationship. × molar mass

 → mass in grams number of moles ←  ÷ molar mass

3.3 The Mole and the Mass of Substances 

5. Equal numbers of moles of substances always contain the same number of particles.

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6. For this, reason, we can compare the number of .particles in substances just by comparing the number of moles of the substances.

3.4 The Mole and the Volume of Gas 

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1. We can determine the number of moles of any gas by measuring its volume. 2. This, however, cannot be done for solids and liquids. 3. The molar volume of a gas is defined as the volume of one mole of the gas. 4. This means that the molar volume of a gas is also the volume occupied by 6.02 x 1023 particles of the gas.

3.4 The Mole and the Volume of Gas 

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5. The molar volume of any gas is 22.4 dm3 mol-1 at STP or 24 dm3 mol-1 at room conditions. For example • 1 mole of oxygen gas occupies 22.4 dm3 at STP. • 1 mole of hydrogen gas occupies 24 dm3 at room temperature.

3.4 The Mole and the Volume of Gas 

6. The following relationship shows the cone of the volume of a gas to the number of moles and vice versa.

× molar volume  → volume of gas number of moles ← ÷ molar volume

3.4 The Mole and the Volume of Gas 

7. The following shows the relationships between the number of moles, number of particles, mass and volume of gases.

3.4 The Mole and the Volume of Gas 

8. In most calculations, we first convert other quantities such as the number of particles, mass or volume to the number of moles.

3.5Chemical Formulae 

1. Each chemical substance it given a name and a chemical formula.

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2. A chemical formula is a representation of a chemical substance using letters for atoms and subscript numbers to show the numbers of each type o f atoms that are present in the substance. 3. For example, the chemical formula of water is H2O. The chemical formula of carbon dioxide is CO2.

3.5Chemical Formulae 

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4. Based on the chemical formula of a substance, we can gather the following information. (a) The elements that make up the substance. (b) The ratio or number of atoms of each element in the substance.

Chemical formulae of elements 

1. For elements that exist as atoms, their chemical formulae represent their atoms. For example, the formulae for carbon and copper are C and Cu respectively.

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2. For elements that exist as molecules, their chemical formulae represent their molecules. The subscript number at the bottom right shows the number of atoms in each molecule.

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3. For example, the chemical formula of oxygen gas is as follows:

Chemical formulae of compounds 

(a) Glucose consists of three elements namely, carbon, hydrogen and oxygen. Six carbon atoms combine with twelve hydrogen atom and six oxygen atoms.

Chemical formulae of compounds 

(b) Magnesium nitrate consists of the elements magnesium, nitrogen and oxygen. Each magnesium atom combines with two nitrogen atoms and six oxygen atoms.

Empirical formulae and molecular formulae 

1. There are two types of chemical formulae - empirical formulae and molecular formulae

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2. The empirical formula of a compound gives the simplest whole number ratio of atoms of each element present in the compound.

Empirical formulae and molecular formulae 

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3. The molecular formula of a compound gives the actual number of atoms of each element that are present in one molecule of the compound. (a) For example, C6H12 O6 is the molecular formula for glucose. From the molecular formula, we know that each molecule of glucose consists of 6 carbon atoms, 12 hydrogen atoms and 6 oxygen atoms. (b) The simplest ratio of carbon to hydrogen to oxygen atoms in the molecule is 1:2:1. Therefore, the empirical formula of glucose is CH2O.

Determining empirical formulae 

1. The empirical formula of a compound can be determined experimentally by finding the simplest ratio of moles of atoms of each element in the compound.

Determining empirical formulae 2.The following shows the steps in determining the empirical formula of a compound.  (a) Find the mass of each element in the compound experimentally.  (b) Convert the masses to the numbers of moles of atoms.  (c) Find the simplest ratio of moles of the elements

Determining empirical formulae Example   2.24 g of iron combines chemically with 0.96 g of oxygen to form an oxide. What is the empirical formula of the oxide? [Relative atomic mass: 0, 16; Fe, 56]  

Determining empirical formulae  Solution: Since 2 moles of iron atoms combine with 3 moles of oxygen atoms, the empirical formula of the oxide is Fe2O3.  

Element

Iron, Fe

Oxygen, O

Mass(g)

2.24

0.96

No. of moles of atoms

2.24 = 0.04 56

Ratio of moles

0.04 =1 0.04

0.06 = 1.5 0.04

2

3

Simplest ratio of moles

0.96 = 0.06 16

Determining empirical formulae Example A potassium compound has a percentage composition as the following. K, 31.84%; Cl, 28.98%; O, 39.18% What is the empirical formula of the potassium compound? [Relative atomic mass: 0, 16; Cl, 35.5; K, 39]

Determining empirical formulae Example 3.23 Solution: From the percentage composition, we know that every 100 g of the compound contains 31.84 g of potassium, 28.98 g of chlorine and 39.18 g of oxygen. So, by taking 100 g of the compound:  

Determining empirical formulae Example Solution: 1 mole of potassium atoms combines with 1 mole of chlorine atoms and 3 moles of oxygen atoms. Therefore, the, empirical formula. of the potassium compound is KClO3

Element

Potassium ,K

Chlorine , Cl

Oxygen, O

Mass(g)

31.84

28.98

39.18

31.84 = 0.816 39

28.98 = 0.816 35.5

No. of moles of atoms Ratio of moles Simplest ratio of moles

0.816 =1 0.816

1

0.816 =1 0.816

1

39.18 = 2.449 16

2.449 =3 0.816

3

Activity: Determining the empirical formula of copper (II) oxide 

Apparatus: Combustion tube with a small hole at the end, Bunsen burner, stoppers, glass tube, retort stand; and clamp, balance, U tube, spatula, porcelain dish.

Activity: determining the empirical formula of copper (II) oxide 

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Apparatus: Combustion tube with a small hole at the end, Bunsen burner, stoppers, glass tube, retort stand; and clamp, balance, U tube, spatula, porcelain dish. Materials: Hydrogen gas, copper(II) oxide, anhydrous calcium chloride, wooden splinter

Activity: determining the empirical formula of copper (II) oxide

Activity: determining the empirical formula of copper (II) oxide 

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1. This method can also be used to determine the empirical formula of oxides of low reactivity metals such as tin(II) oxide and lead(II) oxide. 2. The empirical formula of copper(II) oxide cannot be determined by heating copper(II) oxide with reactive metals such as magnesium or calcium.

Activity: Determining the empirical formula of magnesium oxide

Activity: Determining the empirical formula of magnesium oxide 1. The following are the precautions taken in this activity.  (a) The crucible is covered with its lid to prevent the white fumes of magnesium oxide from escaping. This would affect the accuracy coif the mass obtained.  (b) The lid is removed at intervals to allow oxygen to enter the crucible and reacts with the magnesium ribbon.  (c) Heating, cooling and weighing are repeated until a constant mass is obtained to ensure that, the magnesium ribbon reacts completely to form magnesium oxide.

Finding molecular formulae 

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1. Actually, the molecular formula of a compound is a multiple of its empirical Molecular formula=(Empirical formula) x n whereby n is a positive integer.

Finding molecular formulae 

2. Table 3.5 shows the molecular and empirical formulae of some compounds.

Compound

Empirical formula

Molecular Formula

n

Water

H2O

(H2O)1= H2O

1

Ethane

CH3

(CH3)2= C2H6

2

Propene

CH2

(CH2)3= C3H6

3

Finding molecular formulae  

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Example A hydrocarbon compound has an empirical formula of CH2 and a relative molecular mass of 70. Find the molecular formula of the compound. [Relative atomic mass: H, 1; C, 12]

Finding molecular formulae      

Solution: Let the molecular formula of the compound to be (CH2) n, Based on the formula (CH2) n, the relative molecular mass = n[12 + 2(1)] = 14n However, it is given that the relative molecular mass = 70 Therefore, 14n = 70 n = 70 14

n =5 

So, the molecular formula of the compound is (CH2)5 ,which is C5H10 .

Chemical formulae of ionic compounds 

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1. Ionic compounds are compounds consisting of cations and anions. Cations are positively charged ions whereas anions are negatively charged ions. 2. It is important that you know the formulae of cations and anions before constructing the chemical formulae of ionic compounds.

Chemical formulae of ionic compounds 

Cation (positive ion)

3. Table 3.6 and Table 3.7 show the formulae of some common cations and anions Formula of cation Charge of cation

Sodium ion

Na+

+1

Potassium ion

K+

+1

Silver ion

Ag+

+1

Hydrogen ion

H+

+1

Chemical formulae of ionic compounds Cation (positive ion)

Formula of cation

Charge of cation

Ammonium ion

NH4+

+1

Copper(I) ion

Cu+

+1

Calcium ion

Ca2+

+2

Magnesium ion

Mg2+

+2

Zinc ion

Zn2+

+2

Chemical formulae of ionic compounds Cation (positive ion)

Formula of cation

Charge of cation

Barium ion

Ba2+

+2

Iron(II) ion

Fe2+

+2

Copper(II) ion

Cu2+

+2

Tin(II) ion

Sn2+

+2

Chemical formulae of ionic compounds Cation (positive ion)

Formula of cation

Charge of cation

Lead(II) ion

Pb2+

+2

Aluminium ion

Al3+

+3

Iron(III) ion

Fe3+

+3

Chromium(III) ion

Cr3+

+3

Tin(IV) ion

Sn4+

+4

Lead(IV) ion

Pb4+

+4

Chemical formulae of ionic compounds Anion (negative ion)

Formula of anion

Charge of anion

Flouride ion

F-

-1

Chloride ion

Cl-

-1

Bromide ion

Br-

-1

I-

-1

Hydroxide ion

OH-

-1

Nitrate ion

NO3-

-1

Nitrite ion

NO2-

-1

H-

-1

Iodide ion

Hydride ion

Chemical formulae of ionic compounds Anion (negative ion)

Formula of anion

Charge of anion

Ethanoate ion

CH3COO-

-1

MnO4-

-1

O2-

-2

Carbonate ion

CO32-

-2

Sulphide ion

S2-

-2

Manganate(IV) ion Oxide ion

Chemical formulae of ionic compounds 

3. Table 3.6 and Table 3.7 show the formulae of some common cations and anions

Anion (negative ion)

Formula of anion

Charge of anion

Sulphate ion

SO42-

-2

Sulphite ion

SO32-

-2

Thiosulphate ion

S2O32-

-2

Chromate(VI) ion

CrO42-

-2

Dichromate(VI) ion

Cr2O72-

-2

PO43-

-3

phosphate ion

Chemical formulae of ionic compounds 

4. Even though ionic compounds contain charged particles, their chemical formulae are electrically neutral. This is because the total of positive charges are equal to the total of negative.

Chemical formulae of ionic compounds 

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5. The chemical formula of an ionic compound can be constructed as the following. (a) From its name, identify and write down the formula of its cation and anion. (b) Determine the number of cations and anions by balancing the positive and negative charges. (c) Write the formula of the compound. The number of cations and anions are written as subscript numbers.

Chemical formulae of ionic compounds 

Example

Chemical formulae of ionic compounds 

Example

Naming of chemical compounds 

1. Chemical compounds are named systematic according to the guidelines given by International Union of Pure and App Chemistry (IUPAC).

Naming of chemical compounds 

2. For ionic compounds, the name of the cation comes first, followed by the name of the anion

Cation

Anion

Name of the compound

Sodium ion

Chloride ion

Sodium chloride

Magnesium ion

Oxide ion

Magnesium oxide

Zinc ion

Sulphate ion

Zinc sulphate

Naming of chemical compounds 

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3. Certain elements such as transition metals, can form more than one type of ions. Roman numerals (such as I, II and III) are used in their naming to differentiate the ions. (a) For example, iron can form two cation: Fe 2+ - named as iron(II) ion Fe 3+ - named as iron(III) ion

Naming of chemical compounds 

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3. Certain elements such as transition metals, can form more than one type of ions. Roman numerals (such as I, II and III) are used in their naming to differentiate the ions. (a) For example, iron can form two cation: Fe 2+ - named as iron(II) ion Fe 3+ - named as iron(III) ion (b) Therefore, the names of oxides with iron are, iron(II) oxide and iron(III) oxide respectively.

Naming of chemical compounds 

4. For simple molecular compounds, the name of the first element is maintained as it is. However, the name of the second element is added with an –ide. For example, as molecular; compound consisting of hydrogen and chlorine is given the name hydrogen chloride.

Naming of chemical compounds 

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5. Greek prefixes are used to show the number of atoms of each element in a compound. (a) Here are some examples. CO- Carbon monoxide CO2 – Carbon dioxide

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SO3 – Sulphur trioxide

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CCl4 – Carbon tetrachloride

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PCl5 – Phosphorus pentachloride

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N2O4 – Dinitrogen tetroxide

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Cl2O7 – Dichlorine heptoxide

3 CHEMICAL FORMULAE AND EQUATIONS 3.6 Chemical Equations

3.6Chemical Equations 

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1. A chemical reaction is said to occur when a few starting substances react to produce new substances. 2. The starting substances are called reactants. The new substances formed are called products.

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reactants   → products producing

3.6Chemical Equations 

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3. A chemical equation is a precise description of a chemical reaction. 4. It can be written in words, but it is usually more convenient and quicker to use chemical formulae.

3.6Chemical Equations 

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5. The reactants are written at the lefthand side of the equation whereas the products are written at the right-hand side of the equation. Reactant(s) Product(s) C(s) + O2(g)  CO2(g)

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Zn(s) + Cl2(g)

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HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

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ZnCl2(s)