The Mole Concept and Stoichiometry

Why moles?

Atoms are too small to count or weigh singly, so chemists count them in groups called moles. One mole contains 6.02 x 10^23 particles (Avogadro's constant, L). This is the basis of quantitative chemistry in the KICD Senior School Chemistry syllabus, used in industries such as Magadi soda and sugar processing.

Relative masses

• Relative atomic mass (Ar): e.g. H = 1, C = 12, O = 16, Na = 23, Ca = 40.
• Relative formula mass: sum of Ar values in a formula.
• Molar mass: relative formula mass expressed in g/mol.

Example: Na2CO3 = 2(23) + 12 + 3(16) = 106 g/mol.

Key formulae

• Moles = mass / molar mass
• Number of particles = moles x 6.02 x 10^23
• Gas volume = moles x molar volume (24 dm3 at r.t.p.)
• Solution: moles = molarity (mol/dm3) x volume (dm3)

Worked example 1: mass to moles

How many moles are in 5.3 g of sodium carbonate? Moles = 5.3 / 106 = 0.05 mol.

Worked example 2: particles

How many molecules are in 0.5 mol of water? 0.5 x 6.02 x 10^23 = 3.01 x 10^23 molecules.

Empirical formula

Steps: (1) write percentages or masses, (2) divide by Ar, (3) divide by the smallest, (4) round to whole numbers.

A compound contains 40% C, 6.7% H and 53.3% O.

• C: 40 / 12 = 3.33; H: 6.7 / 1 = 6.7; O: 53.3 / 16 = 3.33
• Divide by 3.33: C 1, H 2, O 1
• Empirical formula CH2O.

Stoichiometry (reacting quantities)

A balanced equation gives the mole ratio of reactants and products.

Worked example 3: Magnesium burns: 2Mg + O2 -> 2MgO. What mass of MgO forms from 12 g of Mg? (Mg = 24, O = 16)

• Moles Mg = 12 / 24 = 0.5
• Ratio Mg : MgO = 1 : 1, so 0.5 mol MgO
• Mass = 0.5 x 40 = 20 g

Worked example 4 (gas volume): Zinc reacts with dilute acid: Zn + 2HCl -> ZnCl2 + H2. What volume of hydrogen at r.t.p. forms from 6.5 g of Zn? (Zn = 65)

• Moles Zn = 0.1; moles H2 = 0.1
• Volume = 0.1 x 24 = 2.4 dm3

Percentage yield

Percentage yield = (actual yield / theoretical yield) x 100.

Common errors

• Using an unbalanced equation.
• Forgetting to multiply subscripts in molar mass.
• Mixing up g and kg or cm3 and dm3 (1 dm3 = 1000 cm3).

Try it

• Calculate the molar mass of calcium hydroxide, Ca(OH)2.
• Find the number of moles in 90 g of water.
• What volume of CO2 at r.t.p. forms when 10 g of CaCO3 reacts with excess acid?
• Find the empirical formula of a compound with 2.4 g C and 0.6 g H.