Home

Chemistry

Atoms, equations and reactions

Balancing equations

Videos

Summary

Exercises

Select Lesson

Exam Board

Select an option

Practical skills

The scientific method

How to plan experiments

Improving repeatability and reproducibility

Presenting results using tables and graphs

Analysing results: anomalies and correlation

Evaluating experiments

Modern analytical techniques II

NMR spectroscopy

Proton NMR spectroscopy

Chromatography

Combining analytical techniques

Organic chemistry III

Benzene: structure and combustion

Electrophilic substitution

Friedel-Crafts reactions

Phenols

Aliphatic and aromatic amines

Reactions of amines

Forming amides

Condensation polymers

Amino acids and zwitterions

Grignard reagents

Synthetic routes

Practical techniques

Percentage composition and combustion analysis

Organic chemistry II

Optical isomers and chirality

Aldehydes and ketones

Carboxylic acids: properties and reactions

Esters: formation, properties and uses

Acyl chlorides: formation and reactions

Kinetics II

Measuring rates of reaction

Determining reaction orders

The initial rates method

Rate equations and the rate constant

The rate determining step

Halogenoalkane reaction mechanisms

Activation energy and the Arrhenius equation

Transition metals

Transition metals: electron configuration and oxidation number

Complex ions: formation, shape and isomerism

Colours of inorganic ions and complexes

Chromium and its reactions

Ligand substitution reactions

Transition metals as catalysts

Redox II

Electrochemical cells and redox

Electrode potentials and calculations

The electrochemical series

Energy storage cells and hydrogen fuel cells

Redox titrations

Energetics II

Lattice enthalpies and Born-Haber cycles

Polarisation and the Pauling scale

Enthalpy change in solution

Entropy and calculating entropy change

Gibbs free energy change

Acid-base equilibria

Acids and bases

pH calculations

The ionic product of water

pH curves and indicators

Buffer solutions and calculations

Equilibrium II

The equilibrium constant

Partial pressure and gas equilibria

Le Chatelier's principle and equilibrium constants

Equilibrium I

Reversible reactions

Le Chatelier's principle

Kinetics I

Collision theory and rates of reaction

Calculating the rate of reaction

Catalysts and the Maxwell-Boltzmann distribution

Energetics I

Enthalpy changes and reaction profiles

Standard enthalpy changes

Hess's Law and enthalpy cycles

Calculating bond enthalpies

Modern analytical techniques I

Mass spectrometry

IR spectroscopy

Organic chemistry I

Basic concepts of organic chemistry

Organic reactions and mechanisms

Structural isomerism

Alkanes and free radical substitution

Fractional distillation of crude oil

Fuels: combustion and emissions

Alkenes: bonding and reactivity

Stereoisomerism

Reactions of alkenes

Addition polymerisation

Halogenoalkanes and their reactions

Alcohol: classification and reactions

Oxidation of alcohols

Organic techniques

Formulae, equations and amounts of substances

The mole and Avogadro's constant

Calculations involving gases

Empirical and molecular formulae

Balancing equations

Acid-base titrations

Calculating uncertainty and errors

Atom economy and percentage yield

Inorganic chemistry and the periodic table

Group 2: ionisation energy and properties

Thermal stability of nitrates and carbonates

Chemical properties of Group 7

Reactions with halogens

Halide ion reactions

Testing for ions

Redox I

Calculating oxidation numbers

Oxidation, reduction and redox reactions

Bonding and structure

Ionic bonding

Covalent bonding

Shapes of molecules

Giant covalent and metallic structures

Electronegativity and polarisation

Intermolecular forces: types and examples

Hydrogen bonding

Solubility

Predicting structures and properties

Atomic structure and the periodic table

The atom: history, structure and isotopes

Relative masses and mass spectrometry

Electronic structure: shells and orbitals

Atomic emission spectra

Ionisation energies

Ionisation energy trends

Periodicity

Explainer Video

Tutor: Aimee LeBrun

Summary

Balancing equations

In a nutshell

Chemical equations are balanced to ensure there is the same number of each type of atom on either side of the reaction.

Balancing an equation

Numbers are placed in front of reactants and products to balance equations and ensure that there are equal numbers of each type of atom on either side of the equation.

procedure

1.	Count the number of atoms on each side of the equation. If the numbers are unequal the equation needs to be balanced.
2.	To balance the equation, add numbers in front of the substances to make the number of each type of atom equal on either side of the reaction.
3.	Before the equation is finished and balanced, be sure to count the numbers of each type of atoms on either side again to be absolutely sure.

Example

Balance the following equation.

$Al \, + \, H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, H_2$

First, count the number of each type of atom on either side of the equation:

type of atom	left-hand side	right-hand side
$Al$	$1$	$2$
$H$	$2$	$2$
$S$	$1$	$3$
$O$	$4$	$12$

Add a two in front of $Al$ on the left-hand side to balance the number of aluminium atoms:

$2Al \, + \, H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, H_2$ 

Recount the number of atoms on either side of the reaction:

$2Al \, + \, H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, H_2$ 

type of atom	left-hand side	right-hand side
$Al$	$2$	$2$
$H$	$2$	$2$
$S$	$1$	$3$
$O$	$4$	$12$

Add a three in front of the $H_2SO_4$ on the left-hand side to balance the number of oxygen and sulfur atoms :

$2Al \, + \, 3H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, H_2$

Recount the number of atoms on either side of the reaction:

$2Al \, + \, 3H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, H_2$ 

type of atom	left-hand side	right-hand side
$Al$	$2$	$2$
$H$	$6$	$2$
$S$	$3$	$3$
$O$	$12$	$12$

The number of hydrogen atoms on the left-hand side has increased, add a three in front of the $H_2$  on the right-hand side to balance the equation:

$2Al \, + \, 3H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, 3H_2$

Finally, be sure to do a final count of each type of atoms either side of the equation to ensure that your equation is balanced correctly.

$2Al \, + \, 3H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, 3H_2$ 

type of atom	left-hand side	right-hand side
$Al$	$2$	$2$
$H$	$6$	$6$
$S$	$3$	$3$
$O$	$12$	$12$

The numbers of each type of atom on either side of the equation are equal and the equation is now balanced; $\underline{2Al \, + \, 3H_2SO_4 \, \rightarrow \, Al_2(SO_4)_3 \, + \, 3H_2}$ .

Finding masses from balanced equations

The stoichiometry discovered when balancing an equation can be used alongside the reacting masses to find the masses of products.

procedure

Use the mass of a reactant to find the number of moles of the substance:

$n \, = \, \frac{m}{M_r}$

The number of moles of the product can be found using the calculated number of moles of the reactant in step 1 along with the mole ratio from the balanced equation.

Use the number of moles of the product to find the amount of mass that should be produced using:

$m \, = \, n \, \times \, M_r$

Ionic equations

Given the full balanced equation of a reaction, ionic equations can be constructed.

PROCEDURE

1.	Distinguish between molecules and ionic compounds.
2.	Distinguish between ionic compounds with ions that can move and substances that are not ions in solution.
3.	Re-write the symbol equation, separating the ions but keeping any solids or substances that are not ions in solution the same.
4.	Cross out ions that are the same on both sides. These are called spectator ions, they are not involved in the reaction.
5.	The ionic equation is left.

Example

Write the ionic equation for the reaction below.

$HCl(aq) \, + \, NaOH(aq) \, \rightarrow \, NaCl(aq) \, + \, H_2O(l)$

First, identify any substance that will not break into ions in solution:

$H_2O(l)$

Next, write out the ions present in this reaction:

$H^+(aq) \, + \, Cl^-(aq) \, + \, Na^+(aq) \, + \, OH^-(aq) \, \rightarrow \, Na^+(aq) \, + \, Cl^- \,(aq) + \, H_2O(l)$

Cross out the ions that are present on both sides of the equation:

$H^+(aq) \, + \, \cancel{Cl^-(aq)} \, + \, \cancel{Na^+(aq)} \, + \, OH^-(aq) \, \rightarrow \, \cancel{Na^+(aq)} \, + \, \cancel{Cl^-(aq)} \, + \, H_2O(l)$

Write the ions that are left:

$H^+(aq) \, + \, OH^-(aq) \, \rightarrow \, H_2O(l)$

The ionic equation for this reaction is $\underline{H^+(aq) \, + \, OH^-(aq) \, \rightarrow \, H_2O(l)}$ .

State symbols

State symbols are used to tell you what state of matter a substance is. State symbols can be used to explain what is happening during a reaction.

STATE	SYMBOL
$Solid$	$(s)$
$Liquid$	$(l)$
$Gas$	$(g)$
$Aqueous ~solution$	$(aq)$

State symbols signify what type of products are produced during reactions.

Precipitate reactions

A precipitate reaction occurs when two aqueous substances react and one of the products is a solid.

Neutralisation reactions

Acids and bases react together to form salt products along with carbon dioxide and water.

Displacement reactions

More reactive metals can displace less reactive metals in a displacement reaction.

Learn with Basics

Length:

Unit 1

Word and symbol equations

Unit 2

Chemical equations

Jump Ahead

Optional

Unit 3