Introducing chemical reactions

What you'll learn

How chemical reactions are represented using symbols, formulae and equations.
Why mass is conserved, even when the measured mass seems to change.
How to write formulae for simple covalent and ionic compounds.
How Higher Tier mole calculations link masses, particles and balanced equations.

Chemical 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.

A chemical formula shows which atoms are present and how many of each there are. For example, H₂O contains hydrogen and oxygen atoms. The small 2 is a subscript; it belongs to the formula and tells you there are two hydrogen atoms in each water molecule.

A chemical equation uses formulae to show a reaction. The arrow → means “reacts to form”.

Definition

Chemical reaction

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

Atoms rearranging in a chemical reaction, showing conservation of hydrogen and oxygen atoms

You should be comfortable using a supplied Periodic Table to find element symbols. Remember that element symbols are case-sensitive: Co is cobalt, but CO is carbon monoxide.

Some elements exist naturally as molecules, such as H₂(g), O₂(g), N₂(g), Cl₂(g), Br₂(l) and I₂(s). Group 0 noble gases, such as He(g), Ne(g) and Ar(g), are single atoms.

Conservation of mass and balanced equations

The law of conservation of mass says that mass is conserved in a chemical reaction. In a closed system, the total mass of reactants equals the total mass of products.

This works because the same atoms are present before and after the reaction — they are just joined together in different ways.

A balanced equation has the same number of each type of atom on both sides. You balance equations by changing the large numbers in front of formulae. These are called coefficients.

Example

Balancing aluminium oxide

Balance: Al(s) + O₂(g) → Al₂O₃(s)

Oxygen comes as O₂ on the left but there are three oxygen atoms in Al₂O₃, so use the lowest common multiple of 2 and 3, which is 6 oxygen atoms.
Put 3 in front of O₂(g) and 2 in front of Al₂O₃(s): Al(s) + 3O₂(g) → 2Al₂O₃(s).
The right side now has four aluminium atoms, so put 4 in front of Al(s): 4Al(s) + 3O₂(g) → 2Al₂O₃(s).

Common Mistake

Changing the formula

Never balance an equation by changing subscripts. Changing Al₂O₃ to Al₂O₂ would make a different substance. Only change coefficients.

Why mass can seem to change

A closed system does not let substances enter or leave. An open system does.

In an open container, the measured mass may change because gas particles can escape or enter from the air. This does not break conservation of mass — it just means the balance is not measuring every particle involved.

Example

Explaining an apparent mass loss

Calcium carbonate is heated in an open test tube:

CaCO₃(s) → CaO(s) + CO₂(g)

The equation shows that carbon dioxide gas is produced.
In an open test tube, CO₂(g) particles can leave the test tube, so the balance reading decreases.
If the escaping CO₂(g) were collected and included, the total mass would be unchanged.

Key Idea

Mass is not really lost

If a reaction appears to lose mass, look for a gas escaping. If it appears to gain mass, look for a gas from the air reacting with the substance.

State symbols

A state symbol shows the physical state of a substance in an equation:

(s) means solid
(l) means liquid
(g) means gas
(aq) means aqueous, which means dissolved in water

For example:

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

Tip

Adding state symbols

Balance the equation first, then add state symbols. State symbols do not affect atom counts.

Writing formulae for ionic compounds

An ion is a charged particle formed when an atom or group of atoms loses or gains electrons. A positive ion is a cation. A negative ion is an anion.

An ionic compound is made from positive and negative ions. Its overall charge must be zero.

Charge balance for magnesium chloride and aluminium oxide ionic formulae

Useful patterns:

Group 1 metal ions, such as Na⁺, have a 1+ charge.
Group 2 metal ions, such as Mg²⁺ and Ca²⁺, have a 2+ charge.
Aluminium forms Al³⁺.
Group 7 halide ions, such as Cl⁻, Br⁻ and I⁻, have a 1− charge.
Common compound ions include OH⁻, NO₃⁻, SO₄²⁻ and CO₃²⁻.

Example

Deducing a formula from ions

Find the formula of magnesium nitrate, made from Mg²⁺ and NO₃⁻ ions.

One Mg²⁺ ion has a charge of 2+, so it needs two nitrate ions, each with charge 1−, to cancel the charge.
Nitrate is a compound ion, so keep NO₃ together and use brackets: Mg(NO₃)₂.
Check the total charge: $+ 2 + 2 (- 1) = 0$ , so the formula is neutral.

Simple covalent compounds are not worked out by charge balance. Their formulae show the atoms in each molecule, such as H₂O, CO₂, CH₄ and NH₃.

Equations involving ions and electrons

This part is Higher Tier only.

A half equation shows either electron loss or electron gain. Electrons are written as e⁻. Balance atoms first, then balance charge using electrons.

Example

Balancing a half equation

Write a half equation for magnesium forming magnesium ions.

Start with one magnesium atom becoming one magnesium ion: Mg(s) → Mg²⁺(aq).
The right side is 2+ overall, so add two electrons to the right to make the total charge zero.
The balanced half equation is: Mg(s) → Mg²⁺(aq) + 2e⁻.

An ionic equation shows only the ions and substances that actually change in a reaction. A spectator ion is an ion that appears unchanged on both sides, so it is left out of the final ionic equation.

Example

Removing spectator ions

Silver nitrate reacts with sodium chloride:

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Split the aqueous ionic compounds into ions: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq).
Identify the spectator ions: Na⁺(aq) and NO₃⁻(aq) appear unchanged on both sides.
Remove them to get the ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

Moles, particles and mass

This part is Higher Tier only.

The amount of substance is measured in moles, symbol mol. One mole contains a very large number of particles.

Definition

Mole and Avogadro constant

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

The relative formula mass, symbol M_r, is found by adding the relative atomic masses, A_r, of all atoms in the formula. For mass calculations, use grams and moles:

n=mMrm=nMr\begin{aligned} n &= \frac{m}{M_r} \\ m &= nM_r \end{aligned}

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

Example

Calculating moles and particles

How many molecules are in 9.00 g of water, H₂O?

Calculate M_r: water has two H atoms and one O atom, so M_r = 18.0.
Convert mass to moles: $\frac{9.00}{18.0} = 0.500\ \text{mol}$ .
Convert moles to particles: $nN_A = 0.500 \times 6.02 \times 10^{23} = 3.01 \times 10^{23}$ molecules.

Tip

Significant figures

Keep extra digits during working, then round the final answer sensibly, usually to the same number of significant figures as the data in the question.

Using balanced equations for masses

This part is also Higher Tier only.

Stoichiometry means the ratio of substances in a balanced chemical equation. For example:

2Mg(s) + O₂(g) → 2MgO(s)

This means 2 mol of magnesium reacts with 1 mol of oxygen to form 2 mol of magnesium oxide.

A limiting reactant is the reactant that runs out first. It limits how much product can be made.

Example

Using mass data and a limiting reactant

Magnesium reacts with oxygen to form magnesium oxide.

From mass data: 4.80 g Mg, 3.20 g O₂ and 8.00 g MgO. Convert to moles: Mg = 0.200 mol, O₂ = 0.100 mol, MgO = 0.200 mol.
Divide by the smallest amount: Mg : O₂ : MgO = 2 : 1 : 2, so the equation is 2Mg(s) + O₂(g) → 2MgO(s).
In a new reaction, 1.20 g Mg gives $\frac{1.20}{24.0} = 0.0500\ \text{mol}$ and 1.00 g O₂ gives $\frac{1.00}{32.0} = 0.0313\ \text{mol}$ .
The equation needs 1 mol O₂ for every 2 mol Mg, so 0.0500 mol Mg needs 0.0250 mol O₂. There is more oxygen than needed, so magnesium is limiting.
The ratio Mg : MgO is 2 : 2, so 0.0500 mol Mg forms 0.0500 mol MgO. Mass of MgO = $0.0500 \times 40.0 = 2.00\ \text{g}$ .

Exam technique

In the exam

Write correct formulae first, then balance by changing coefficients only; add state symbols once the equation is balanced.
For ionic work, check total charge is zero in formulae, and remove only spectator ions when forming ionic equations.
For mass questions, go mass → moles → equation ratio → moles → mass, and always check for the limiting reactant.

Self review

Check yourself

Why must you not change subscripts when balancing a chemical equation?
What formula would you write for aluminium sulfate from Al³⁺ and SO₄²⁻ ions?
Why might the mass decrease when a carbonate is heated in an open test tube?

Introducing chemical reactions

What you'll learn

How chemical reactions are represented using symbols, formulae and equations.
Why mass is conserved, even when the measured mass seems to change.
How to write formulae for simple covalent and ionic compounds.
How Higher Tier mole calculations link masses, particles and balanced equations.

Chemical 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.

A chemical equation uses formulae to show a reaction. The arrow → means “reacts to form”.

Definition

Chemical reaction

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

Atoms rearranging in a chemical reaction, showing conservation of hydrogen and oxygen atoms

You should be comfortable using a supplied Periodic Table to find element symbols. Remember that element symbols are case-sensitive: Co is cobalt, but CO is carbon monoxide.

Some elements exist naturally as molecules, such as H₂(g), O₂(g), N₂(g), Cl₂(g), Br₂(l) and I₂(s). Group 0 noble gases, such as He(g), Ne(g) and Ar(g), are single atoms.

Conservation of mass and balanced equations

The law of conservation of mass says that mass is conserved in a chemical reaction. In a closed system, the total mass of reactants equals the total mass of products.

This works because the same atoms are present before and after the reaction — they are just joined together in different ways.

A balanced equation has the same number of each type of atom on both sides. You balance equations by changing the large numbers in front of formulae. These are called coefficients.

Example

Balancing aluminium oxide

Balance: Al(s) + O₂(g) → Al₂O₃(s)

Oxygen comes as O₂ on the left but there are three oxygen atoms in Al₂O₃, so use the lowest common multiple of 2 and 3, which is 6 oxygen atoms.
Put 3 in front of O₂(g) and 2 in front of Al₂O₃(s): Al(s) + 3O₂(g) → 2Al₂O₃(s).
The right side now has four aluminium atoms, so put 4 in front of Al(s): 4Al(s) + 3O₂(g) → 2Al₂O₃(s).

Common Mistake

Changing the formula

Never balance an equation by changing subscripts. Changing Al₂O₃ to Al₂O₂ would make a different substance. Only change coefficients.

Why mass can seem to change

A closed system does not let substances enter or leave. An open system does.

Example

Explaining an apparent mass loss

Calcium carbonate is heated in an open test tube:

CaCO₃(s) → CaO(s) + CO₂(g)

The equation shows that carbon dioxide gas is produced.
In an open test tube, CO₂(g) particles can leave the test tube, so the balance reading decreases.
If the escaping CO₂(g) were collected and included, the total mass would be unchanged.

Key Idea

Mass is not really lost

If a reaction appears to lose mass, look for a gas escaping. If it appears to gain mass, look for a gas from the air reacting with the substance.

State symbols

A state symbol shows the physical state of a substance in an equation:

(s) means solid
(l) means liquid
(g) means gas
(aq) means aqueous, which means dissolved in water

For example:

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

Tip

Adding state symbols

Balance the equation first, then add state symbols. State symbols do not affect atom counts.

Writing formulae for ionic compounds

An ion is a charged particle formed when an atom or group of atoms loses or gains electrons. A positive ion is a cation. A negative ion is an anion.

An ionic compound is made from positive and negative ions. Its overall charge must be zero.

Charge balance for magnesium chloride and aluminium oxide ionic formulae

Useful patterns:

Group 1 metal ions, such as Na⁺, have a 1+ charge.
Group 2 metal ions, such as Mg²⁺ and Ca²⁺, have a 2+ charge.
Aluminium forms Al³⁺.
Group 7 halide ions, such as Cl⁻, Br⁻ and I⁻, have a 1− charge.
Common compound ions include OH⁻, NO₃⁻, SO₄²⁻ and CO₃²⁻.

Example

Deducing a formula from ions

Find the formula of magnesium nitrate, made from Mg²⁺ and NO₃⁻ ions.

One Mg²⁺ ion has a charge of 2+, so it needs two nitrate ions, each with charge 1−, to cancel the charge.
Nitrate is a compound ion, so keep NO₃ together and use brackets: Mg(NO₃)₂.
Check the total charge: $+ 2 + 2 (- 1) = 0$ , so the formula is neutral.

Simple covalent compounds are not worked out by charge balance. Their formulae show the atoms in each molecule, such as H₂O, CO₂, CH₄ and NH₃.

Equations involving ions and electrons

This part is Higher Tier only.

A half equation shows either electron loss or electron gain. Electrons are written as e⁻. Balance atoms first, then balance charge using electrons.

Example

Balancing a half equation

Write a half equation for magnesium forming magnesium ions.

Start with one magnesium atom becoming one magnesium ion: Mg(s) → Mg²⁺(aq).
The right side is 2+ overall, so add two electrons to the right to make the total charge zero.
The balanced half equation is: Mg(s) → Mg²⁺(aq) + 2e⁻.

Example

Removing spectator ions

Silver nitrate reacts with sodium chloride:

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Split the aqueous ionic compounds into ions: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq).
Identify the spectator ions: Na⁺(aq) and NO₃⁻(aq) appear unchanged on both sides.
Remove them to get the ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

Moles, particles and mass

This part is Higher Tier only.

The amount of substance is measured in moles, symbol mol. One mole contains a very large number of particles.

Definition

Mole and Avogadro constant

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

The relative formula mass, symbol M_r, is found by adding the relative atomic masses, A_r, of all atoms in the formula. For mass calculations, use grams and moles:

n=mMrm=nMr\begin{aligned} n &= \frac{m}{M_r} \\ m &= nM_r \end{aligned}

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

Example

Calculating moles and particles

How many molecules are in 9.00 g of water, H₂O?

Calculate M_r: water has two H atoms and one O atom, so M_r = 18.0.
Convert mass to moles: $\frac{9.00}{18.0} = 0.500\ \text{mol}$ .
Convert moles to particles: $nN_A = 0.500 \times 6.02 \times 10^{23} = 3.01 \times 10^{23}$ molecules.

Tip

Significant figures

Keep extra digits during working, then round the final answer sensibly, usually to the same number of significant figures as the data in the question.

Using balanced equations for masses

This part is also Higher Tier only.

Stoichiometry means the ratio of substances in a balanced chemical equation. For example:

2Mg(s) + O₂(g) → 2MgO(s)

This means 2 mol of magnesium reacts with 1 mol of oxygen to form 2 mol of magnesium oxide.

A limiting reactant is the reactant that runs out first. It limits how much product can be made.

Example

Using mass data and a limiting reactant

Magnesium reacts with oxygen to form magnesium oxide.

From mass data: 4.80 g Mg, 3.20 g O₂ and 8.00 g MgO. Convert to moles: Mg = 0.200 mol, O₂ = 0.100 mol, MgO = 0.200 mol.
Divide by the smallest amount: Mg : O₂ : MgO = 2 : 1 : 2, so the equation is 2Mg(s) + O₂(g) → 2MgO(s).
In a new reaction, 1.20 g Mg gives $\frac{1.20}{24.0} = 0.0500\ \text{mol}$ and 1.00 g O₂ gives $\frac{1.00}{32.0} = 0.0313\ \text{mol}$ .
The equation needs 1 mol O₂ for every 2 mol Mg, so 0.0500 mol Mg needs 0.0250 mol O₂. There is more oxygen than needed, so magnesium is limiting.
The ratio Mg : MgO is 2 : 2, so 0.0500 mol Mg forms 0.0500 mol MgO. Mass of MgO = $0.0500 \times 40.0 = 2.00\ \text{g}$ .

Exam technique

In the exam

Write correct formulae first, then balance by changing coefficients only; add state symbols once the equation is balanced.
For ionic work, check total charge is zero in formulae, and remove only spectator ions when forming ionic equations.
For mass questions, go mass → moles → equation ratio → moles → mass, and always check for the limiting reactant.

Self review

Check yourself

Why must you not change subscripts when balancing a chemical equation?
What formula would you write for aluminium sulfate from Al³⁺ and SO₄²⁻ ions?
Why might the mass decrease when a carbonate is heated in an open test tube?

Introducing chemical reactions

What you'll learn

Chemical reactions are rearrangements of atoms

Conservation of mass and balanced equations

Why mass can seem to change

State symbols

Writing formulae for ionic compounds

Equations involving ions and electrons

Moles, particles and mass

Using balanced equations for masses

Recap questions

You've reached the end

Introducing chemical reactions

What you'll learn

Chemical reactions are rearrangements of atoms

Conservation of mass and balanced equations

Why mass can seem to change

State symbols

Writing formulae for ionic compounds

Equations involving ions and electrons

Moles, particles and mass

Using balanced equations for masses

Recap questions

You've reached the end

Introducing chemical reactions

What you'll learn

Chemical reactions are rearrangements of atoms

Conservation of mass and balanced equations

Why mass can seem to change

State symbols

Writing formulae for ionic compounds

Equations involving ions and electrons

Moles, particles and mass

Using balanced equations for masses

Recap questions

You've reached the end

Introducing chemical reactions Revision Guide

Introducing chemical reactions

What you'll learn

Chemical reactions are rearrangements of atoms

Conservation of mass and balanced equations

Why mass can seem to change

State symbols

Writing formulae for ionic compounds

Equations involving ions and electrons

Moles, particles and mass

Using balanced equations for masses

Recap questions

You've reached the end

Introducing chemical reactions Revision Guide