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Chapter 5: Chemical Formulae and Equation

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Chapter 5: Chemical Formulae and Equation

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What is molecular formula & empirical formula?

The molecular formula is the formula that shows the number and type of each atom in a molecule

E.g. the molecular formula of ethanoic acid is C₂H₄O₂

The empirical formula is the simplest whole number ratio of the atoms of each element present in one molecule or formula unit of the compound

E.g. the empirical formula of ethanoic acid is CH₂O

Organic molecules often have different empirical and molecular formulae

The formula of an ionic compound is always an empirical formula

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Calculate Empirical & Molecular Formulae

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Calculate Empirical & Molecular Formulae

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Empirical formula: gives the simplest whole number ratio of atoms of each element in the compound

It is calculated from knowledge of the ratio of masses of each element in the compound

Example: A compound that contains 10 g of hydrogen and 80 g of oxygen has an empirical formula of H₂O. This can be shown by the following calculations:

Amount of hydrogen atoms = mass in grams ÷ Ar of hydrogen = (10 ÷ 1) = 10 moles
Amount of oxygen atoms = mass in grams ÷ Ar of oxygen = (80 ÷ 16) = 5 moles

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Since equal numbers of moles of atoms contain the same number of atoms, the ratio of hydrogen atoms to oxygen atoms is 2:1. Hence the empirical formula is H₂O
Molecular formula: gives the exact numbers of atoms of each element present in the formula of the compound

Divide the relative formula mass of the molecular formula by the relative formula mass of the empirical formula

Multiply the number of each element present in the empirical formula by this number to find the molecular formula

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Writing Equations

Nothing created - nothing destroyed

New substances are made during chemical reactions

However, the same atoms are always present before and after reaction

They have just joined up in different ways

Atoms cannot be created or destroyed, so if they exist in the reactants then they absolutely must be in the products!

Because of this the total mass of reactants is always equal to the total mass of products

This idea is known as the Law of Conservation of Mass

Conservation of Mass

The Law of Conservation of Mass enables us to balance chemical equations, since no atoms can be lost or created

You should be able to:

Write word equations for reactions outlined in these notes

Write formulae and balanced chemical equations for the reactions in these notes

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These show the reactants and products of a chemical reaction using their full chemical names

The reactants are those substances on the left-hand side of the arrow and can be thought of as the chemical ingredients of the reaction

They react with each other and form new substances

The products are the new substances which are on the right-hand side of the arrow

The arrow (which is spoken as “goes to” or “produces”) implies the conversion of reactants into products

Reaction conditions or the name of a catalyst (a substance added to make a reaction go faster) can be written above the arrow

An example is the reaction of sodium hydroxide (a base) and hydrochloric acid produces sodium chloride (common table salt) and water:

sodium hydroxide + hydrochloric acid ⟶ sodium chloride + water

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Chemical equations use the chemical symbols of each reactant and product

When balancing equations, there has to be the same number of atoms of each element on either side of the equation in accordance with the Law of Conservation of Mass

A symbol equation uses the formulae of the reactants and products to show what happens in a chemical reaction

A symbol equation must be balanced to give the correct ratio of reactants and products:

S + O₂ → SO₂

This equation shows that one atom of sulfur (S) reacts with one  molecule of oxygen (O₂) to make one molecule of sulfur dioxide (SO₂)

The following non-metals must be written as molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂

To balance an equation you work across the equation from left to right, checking one element after another

If there is a group of atoms, for example a nitrate group (NO₃–) that has not changed from one side to the other, then count the whole group as one entity rather than counting the individual atoms

Examples of chemical equations:

Acid-base neutralisation reaction:

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

Redox reaction:

2Fe₂O3 (aq) + 3C (s) ⟶ 4Fe (s) + 3CO₂ (g)

In each equation there are equal numbers of each atom on either side of the reaction arrow so the equations are balanced

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In aqueous solutions ionic compounds dissociate into their ions, meaning they separate into the component ions that formed them, e.g. hydrochloric acid and potassium hydroxide dissociate as follows:

HCl (aq) →  H⁺ (aq) + Cl⁻(aq)

KOH (aq)  → K+ (aq)  + OH- (aq)

It is important that you can recognise common ionic compounds and their constituent ions

These include:

Acids such as HCl and H₂SO₄

Group I and Group II hydroxides e.g. sodium hydroxide

Soluble salts e.g. potassium sulfate, sodium chloride

Follow the example below to write ionic equations