Introducing chemical reactions

What you'll learn

How to write formulae for elements, covalent compounds and ionic compounds.
How balanced chemical equations show conservation of atoms and mass.
How to use state symbols and test for common gases.
If you study Higher Tier, how moles, concentration, Avogadro’s constant and limiting reactants connect equations to real masses.

Reactions are rearrangements of atoms

A chemical reaction happens when substances change into new substances. The starting substances are called reactants. The new substances formed are called products.

Definition

Chemical reaction

A chemical reaction is a change in which atoms are rearranged to make new substances. Atoms are not created, destroyed or changed into different elements.

For example, hydrogen reacts with oxygen to form water: 2H₂(g) + O₂(g) → 2H₂O(l). The diagram shows why the equation must be balanced: the number of each type of atom is the same before and after the reaction.

Balanced reaction showing atoms rearranged in 2H2 plus O2 forming 2H2O

Key Idea

Conservation of mass

In a closed system, no substances enter or leave, so the total mass of reactants equals the total mass of products. This is the law of conservation of mass.

In a non-enclosed system — an open container, for example — the measured mass may change because a gas escapes or a gas from the air joins the reaction.

Common Mistake

Saying mass is lost

If a balance reading goes down during a reaction, do not say “mass has disappeared”. Say that a gas product has escaped from the apparatus, so it is no longer being measured.

Writing formulae

A chemical symbol represents an element, such as H for hydrogen, O for oxygen and Na for sodium. Be careful with capital letters: Co means cobalt, but CO means carbon monoxide.

A formula shows which atoms or ions are present, and in what ratio. You should be able to use a supplied Periodic Table, especially for the first 20 elements, Groups 1, 7 and 0, plus common elements used elsewhere in the course.

Simple covalent compounds

A covalent compound is made when non-metal atoms share electrons. Its formula shows the number of atoms in one molecule.

Examples include:

water: H₂O
carbon dioxide: CO₂
methane: CH₄

Ionic compounds

An ion is a charged particle. An ionic compound is made from positive ions and negative ions. The overall charge of the compound must be zero.

Common ions include Na⁺, K⁺, Mg²⁺, Ca²⁺, Al³⁺, Cl⁻, O²⁻, OH⁻, NO₃⁻, SO₄²⁻ and CO₃²⁻.

Example

Deducing an ionic formula

Work out the formula of aluminium sulfate from Al³⁺ and SO₄²⁻.

Aluminium ions have a charge of 3+, while sulfate ions have a charge of 2−.
The smallest total charge that both 3 and 2 fit into is 6, so use two Al³⁺ ions for 6+ and three SO₄²⁻ ions for 6−.
Write the formula as Al₂(SO₄)₃. The brackets show that there are three whole sulfate ions.

Tip

Do not change ion formulae

When writing ionic formulae, change the number of ions, not the formula of the ion itself. For example, sulfate stays as SO₄²⁻; it does not become SO₃.

Balanced chemical equations

A chemical equation uses formulae to show a reaction. A number in front of a formula is called a coefficient. It multiplies the whole formula.

A state symbol shows the physical state:

(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous, meaning dissolved in water

Example

Balancing aluminium oxide

Balance this equation: Al(s) + O₂(g) → Al₂O₃(s)

Oxygen atoms come in pairs on the left but threes on the right, so make 6 oxygen atoms: 3O₂(g) on the left and 2Al₂O₃(s) on the right.
Two Al₂O₃ units contain 4 aluminium atoms, so put 4Al(s) on the left.
The balanced equation is 4Al(s) + 3O₂(g) → 2Al₂O₃(s).

Common Mistake

Changing subscripts when balancing

Never change a formula to balance an equation. Change coefficients only. For example, write 2H₂O, not H₂O₂, unless the substance really is hydrogen peroxide.

Ionic equations and half equations

On Higher Tier, you also need to focus on ions that actually change during reactions.

A spectator ion is an ion present in the reaction mixture that does not change. An ionic equation leaves out spectator ions.

Example

Constructing an ionic equation

Silver nitrate solution reacts with sodium chloride solution to form silver chloride precipitate:
AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Split the aqueous ionic compounds into ions: Ag⁺(aq), NO₃⁻(aq), Na⁺(aq) and Cl⁻(aq).
Identify the ions that stay aqueous and unchanged: Na⁺(aq) and NO₃⁻(aq) are spectator ions.
Remove the spectator ions to get: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

A half equation shows electrons gained or lost by one reactant. For example:

Mg(s) → Mg²⁺(aq) + 2e⁻
Cl₂(g) + 2e⁻ → 2Cl⁻(aq)

Tests for selected gases

You need to know these gas tests accurately.

Gas	Positive test
Oxygen	A glowing splint relights.
Hydrogen	A lit splint gives a squeaky pop.
Carbon dioxide	Limewater turns milky or cloudy.
Chlorine	Damp blue litmus paper is bleached white, often after turning red first.

The mole and Avogadro’s constant

On Higher Tier, reactions are linked to the amount of substance, measured in moles, symbol mol.

Definition

Mole

One mole contains $6.02 \times 10^{23}$ particles. This number is the Avogadro constant, $mol−1N_A = 6.02 \times 10^{23}\ \text{mol}^{-1}$ .

The particles might be atoms, molecules, ions or formula units. The relative formula mass, $M_r$ , is found by adding the relative atomic masses in the formula.

The main relationships are:

\frac{m}{M_r}

where $n$ is amount in mol and $m$ is mass in grams.

Flow diagram linking mass, moles, particles, Avogadro constant and concentration

Example

Finding the mass of one molecule

Find the mass of one carbon dioxide molecule. Use C = 12, O = 16.

Calculate the relative formula mass: $Mr(CO2)=12+2(16)=44M_r(\text{CO}_2) = 12 + 2(16) = 44$ .
One mole of CO₂ molecules has a mass of 44 g and contains $6.02 \times 10^{23}$ molecules.
Divide by Avogadro’s constant: $g\frac{44\ \text{g}}{6.02 \times 10^{23}} = 7.31 \times 10^{-23}\ \text{g}$ per molecule.

Concentration of solutions

A solute is the substance dissolved. A solution is the mixture formed when a solute dissolves in a solvent, usually water.

On Higher Tier, concentration can be measured in mol/dm³:

\frac{n}{V}

where $c$ is concentration, $n$ is amount in mol and $V$ is volume in dm³. Remember: 1000 cm³ = 1 dm³.

Example

Calculating mass from concentration

Calculate the mass of sodium chloride needed to make 250 cm³ of 0.200 mol/dm³ NaCl solution. Use Na = 23, Cl = 35.5.

Convert the volume: 250 cm³ = 0.250 dm³.
Calculate moles: $0.200\ \text{mol/dm}^3 \times 0.250\ \text{dm}^3 = 0.0500\ \text{mol}$ .
Calculate mass using $Mr(NaCl)=58.5M_r(\text{NaCl}) = 58.5$ : $nM_r = 0.0500 \times 58.5 = 2.93\ \text{g}$ .

Stoichiometry and limiting reactants

Stoichiometry means using the ratios in a balanced equation. These ratios are mole ratios, not mass ratios.

A limiting reactant is the reactant that runs out first. It limits the maximum amount of product made.

Example

Identifying a limiting reactant

Magnesium reacts with oxygen: 2Mg(s) + O₂(g) → 2MgO(s).
4.8 g of magnesium reacts with 4.8 g of oxygen. Use Mg = 24 and O = 16.

Convert masses to moles: $moln(\text{Mg}) = \frac{4.8}{24} = 0.20\ \text{mol}$ and $moln(\text{O}_2) = \frac{4.8}{32} = 0.15\ \text{mol}$ .
Use the equation ratio: 2 mol Mg needs 1 mol O₂, so 0.20 mol Mg needs 0.10 mol O₂.
Compare with what is available: 0.15 mol O₂ is available, so oxygen is in excess and magnesium is limiting.
Use the product ratio: 2 mol Mg makes 2 mol MgO, so 0.20 mol Mg makes 0.20 mol MgO.
Convert to mass: $gm(\text{MgO}) = 0.20 \times 40 = 8.0\ \text{g}$ .

To deduce a formula from masses, convert each mass to moles, then reduce to the simplest whole-number ratio. This is the same ratio thinking used in balancing equations.

Exam technique

In the exam

For formulae, check charges cancel to zero; for equations, count atoms on both sides.
Always include state symbols when writing chemical equations, and use the exact gas test observations.
For Higher calculations, write the balanced equation first, convert masses to moles, use the mole ratio, then convert back if needed.

Self review

Check yourself

Why can the measured mass decrease when a carbonate reacts with acid in an open flask?
What is the formula of aluminium sulfate, and why are brackets needed?
How would you test for oxygen, hydrogen and carbon dioxide?

What you'll learn

How to write formulae for elements, covalent compounds and ionic compounds.
How balanced chemical equations show conservation of atoms and mass.
How to use state symbols and test for common gases.
If you study Higher Tier, how moles, concentration, Avogadro’s constant and limiting reactants connect equations to real masses.

Reactions are rearrangements of atoms

A chemical reaction happens when substances change into new substances. The starting substances are called reactants. The new substances formed are called products.

Definition

Chemical reaction

A chemical reaction is a change in which atoms are rearranged to make new substances. Atoms are not created, destroyed or changed into different elements.

Balanced reaction showing atoms rearranged in 2H2 plus O2 forming 2H2O

Key Idea

Conservation of mass

In a closed system, no substances enter or leave, so the total mass of reactants equals the total mass of products. This is the law of conservation of mass.

In a non-enclosed system — an open container, for example — the measured mass may change because a gas escapes or a gas from the air joins the reaction.

Common Mistake

Saying mass is lost

If a balance reading goes down during a reaction, do not say “mass has disappeared”. Say that a gas product has escaped from the apparatus, so it is no longer being measured.

Writing formulae

A chemical symbol represents an element, such as H for hydrogen, O for oxygen and Na for sodium. Be careful with capital letters: Co means cobalt, but CO means carbon monoxide.

Simple covalent compounds

A covalent compound is made when non-metal atoms share electrons. Its formula shows the number of atoms in one molecule.

Examples include:

water: H₂O
carbon dioxide: CO₂
methane: CH₄

Ionic compounds

An ion is a charged particle. An ionic compound is made from positive ions and negative ions. The overall charge of the compound must be zero.

Common ions include Na⁺, K⁺, Mg²⁺, Ca²⁺, Al³⁺, Cl⁻, O²⁻, OH⁻, NO₃⁻, SO₄²⁻ and CO₃²⁻.

Example

Deducing an ionic formula

Work out the formula of aluminium sulfate from Al³⁺ and SO₄²⁻.

Aluminium ions have a charge of 3+, while sulfate ions have a charge of 2−.
The smallest total charge that both 3 and 2 fit into is 6, so use two Al³⁺ ions for 6+ and three SO₄²⁻ ions for 6−.
Write the formula as Al₂(SO₄)₃. The brackets show that there are three whole sulfate ions.

Tip

Do not change ion formulae

When writing ionic formulae, change the number of ions, not the formula of the ion itself. For example, sulfate stays as SO₄²⁻; it does not become SO₃.

Balanced chemical equations

A chemical equation uses formulae to show a reaction. A number in front of a formula is called a coefficient. It multiplies the whole formula.

A state symbol shows the physical state:

(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous, meaning dissolved in water

Example

Balancing aluminium oxide

Balance this equation: Al(s) + O₂(g) → Al₂O₃(s)

Oxygen atoms come in pairs on the left but threes on the right, so make 6 oxygen atoms: 3O₂(g) on the left and 2Al₂O₃(s) on the right.
Two Al₂O₃ units contain 4 aluminium atoms, so put 4Al(s) on the left.
The balanced equation is 4Al(s) + 3O₂(g) → 2Al₂O₃(s).

Common Mistake

Changing subscripts when balancing

Never change a formula to balance an equation. Change coefficients only. For example, write 2H₂O, not H₂O₂, unless the substance really is hydrogen peroxide.

Ionic equations and half equations

On Higher Tier, you also need to focus on ions that actually change during reactions.

A spectator ion is an ion present in the reaction mixture that does not change. An ionic equation leaves out spectator ions.

Example

Constructing an ionic equation

Silver nitrate solution reacts with sodium chloride solution to form silver chloride precipitate:
AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Split the aqueous ionic compounds into ions: Ag⁺(aq), NO₃⁻(aq), Na⁺(aq) and Cl⁻(aq).
Identify the ions that stay aqueous and unchanged: Na⁺(aq) and NO₃⁻(aq) are spectator ions.
Remove the spectator ions to get: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

A half equation shows electrons gained or lost by one reactant. For example:

Mg(s) → Mg²⁺(aq) + 2e⁻
Cl₂(g) + 2e⁻ → 2Cl⁻(aq)

Tests for selected gases

You need to know these gas tests accurately.

Gas	Positive test
Oxygen	A glowing splint relights.
Hydrogen	A lit splint gives a squeaky pop.
Carbon dioxide	Limewater turns milky or cloudy.
Chlorine	Damp blue litmus paper is bleached white, often after turning red first.

The mole and Avogadro’s constant

On Higher Tier, reactions are linked to the amount of substance, measured in moles, symbol mol.

Definition

Mole

One mole contains $6.02 \times 10^{23}$ particles. This number is the Avogadro constant, $mol−1N_A = 6.02 \times 10^{23}\ \text{mol}^{-1}$ .

The particles might be atoms, molecules, ions or formula units. The relative formula mass, $M_r$ , is found by adding the relative atomic masses in the formula.

The main relationships are:

\frac{m}{M_r}

where $n$ is amount in mol and $m$ is mass in grams.

Flow diagram linking mass, moles, particles, Avogadro constant and concentration

Example

Finding the mass of one molecule

Find the mass of one carbon dioxide molecule. Use C = 12, O = 16.

Calculate the relative formula mass: $Mr(CO2)=12+2(16)=44M_r(\text{CO}_2) = 12 + 2(16) = 44$ .
One mole of CO₂ molecules has a mass of 44 g and contains $6.02 \times 10^{23}$ molecules.
Divide by Avogadro’s constant: $g\frac{44\ \text{g}}{6.02 \times 10^{23}} = 7.31 \times 10^{-23}\ \text{g}$ per molecule.

Concentration of solutions

A solute is the substance dissolved. A solution is the mixture formed when a solute dissolves in a solvent, usually water.

On Higher Tier, concentration can be measured in mol/dm³:

\frac{n}{V}

where $c$ is concentration, $n$ is amount in mol and $V$ is volume in dm³. Remember: 1000 cm³ = 1 dm³.

Example

Calculating mass from concentration

Calculate the mass of sodium chloride needed to make 250 cm³ of 0.200 mol/dm³ NaCl solution. Use Na = 23, Cl = 35.5.

Convert the volume: 250 cm³ = 0.250 dm³.
Calculate moles: $0.200\ \text{mol/dm}^3 \times 0.250\ \text{dm}^3 = 0.0500\ \text{mol}$ .
Calculate mass using $Mr(NaCl)=58.5M_r(\text{NaCl}) = 58.5$ : $nM_r = 0.0500 \times 58.5 = 2.93\ \text{g}$ .

Stoichiometry and limiting reactants

Stoichiometry means using the ratios in a balanced equation. These ratios are mole ratios, not mass ratios.

A limiting reactant is the reactant that runs out first. It limits the maximum amount of product made.

Example

Identifying a limiting reactant

Magnesium reacts with oxygen: 2Mg(s) + O₂(g) → 2MgO(s).
4.8 g of magnesium reacts with 4.8 g of oxygen. Use Mg = 24 and O = 16.

Convert masses to moles: $moln(\text{Mg}) = \frac{4.8}{24} = 0.20\ \text{mol}$ and $moln(\text{O}_2) = \frac{4.8}{32} = 0.15\ \text{mol}$ .
Use the equation ratio: 2 mol Mg needs 1 mol O₂, so 0.20 mol Mg needs 0.10 mol O₂.
Compare with what is available: 0.15 mol O₂ is available, so oxygen is in excess and magnesium is limiting.
Use the product ratio: 2 mol Mg makes 2 mol MgO, so 0.20 mol Mg makes 0.20 mol MgO.
Convert to mass: $gm(\text{MgO}) = 0.20 \times 40 = 8.0\ \text{g}$ .

To deduce a formula from masses, convert each mass to moles, then reduce to the simplest whole-number ratio. This is the same ratio thinking used in balancing equations.

Exam technique

In the exam

For formulae, check charges cancel to zero; for equations, count atoms on both sides.
Always include state symbols when writing chemical equations, and use the exact gas test observations.
For Higher calculations, write the balanced equation first, convert masses to moles, use the mole ratio, then convert back if needed.

Self review

Check yourself

Why can the measured mass decrease when a carbonate reacts with acid in an open flask?
What is the formula of aluminium sulfate, and why are brackets needed?
How would you test for oxygen, hydrogen and carbon dioxide?

What you'll learn

Reactions are rearrangements of atoms

Writing formulae

Simple covalent compounds

Ionic compounds

Balanced chemical equations

Ionic equations and half equations

Tests for selected gases

The mole and Avogadro’s constant

Concentration of solutions

Stoichiometry and limiting reactants

Review Introducing chemical reactions by teaching Genie

Recap your knowledge with an interactive lesson

Remember key concepts with flashcards

Introducing chemical reactions

What you'll learn

Reactions are rearrangements of atoms

Writing formulae

Simple covalent compounds

Ionic compounds

Balanced chemical equations

Ionic equations and half equations

Tests for selected gases

The mole and Avogadro’s constant

Concentration of solutions

Stoichiometry and limiting reactants

Review Introducing chemical reactions by teaching Genie

Recap your knowledge with an interactive lesson

Remember key concepts with flashcards

Introducing chemical reactions

What you'll learn

Reactions are rearrangements of atoms

Writing formulae

Simple covalent compounds

Ionic compounds

Balanced chemical equations

Ionic equations and half equations

Tests for selected gases

The mole and Avogadro’s constant

Concentration of solutions

Stoichiometry and limiting reactants

Review Introducing chemical reactions by teaching Genie

Recap your knowledge with an interactive lesson

Remember key concepts with flashcards

Introducing chemical reactions Revision Guide

Introducing chemical reactions

What you'll learn

Reactions are rearrangements of atoms

Writing formulae

Simple covalent compounds

Ionic compounds

Balanced chemical equations

Ionic equations and half equations

Tests for selected gases

The mole and Avogadro’s constant

Concentration of solutions

Stoichiometry and limiting reactants

Review Introducing chemical reactions by teaching Genie

Recap your knowledge with an interactive lesson

Remember key concepts with flashcards

Introducing chemical reactions Revision Guide