GCSE Chemistry

Key Reaction Types

• Combustion: substance reacts with oxygen, releasing heat and light. E.g. CH₄ + 2O₂ → CO₂ + 2H₂O
• Thermal decomposition: a substance breaks down when heated. E.g. CaCO₃ → CaO + CO₂
• Neutralisation: acid + base → salt + water. E.g. HCl + NaOH → NaCl + H₂O
• Displacement: a more reactive element replaces a less reactive one. E.g. Zn + CuSO₄ → ZnSO₄ + Cu
• Oxidation: gain of oxygen or loss of electrons (OIL — Oxidation Is Loss)
• Reduction: loss of oxygen or gain of electrons (RIG — Reduction Is Gain)
• Precipitation: two solutions react to form an insoluble solid (precipitate). E.g. AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Chemical equations must be balanced — the same number of each type of atom on both sides (conservation of mass).

Steps to Balance an Equation

1. Write the unbalanced equation with correct formulae
2. Count atoms of each element on both sides
3. Adjust coefficients (big numbers in front) to balance — never change the formulae
4. Check all atoms balance

Example

State Symbols

The pH Scale

• pH 0-6: Acidic (lower = stronger acid)
• pH 7: Neutral
• pH 8-14: Alkaline (higher = stronger alkali)

Indicators: Litmus (red in acid, blue in alkali), Universal indicator (range of colours), Phenolphthalein (colourless in acid, pink in alkali).

Strong vs Weak Acids

• Strong acids (HCl, H₂SO₄, HNO₃) are fully dissociated (ionised) in water
• Weak acids (CH₃COOH, H₂CO₃) are only partially dissociated — equilibrium lies to the left

Reactions of Acids

Examples

• HCl + NaOH → NaCl + H₂O
• 2HCl + Mg → MgCl₂ + H₂
• H₂SO₄ + CuO → CuSO₄ + H₂O
• 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂

Naming Salts

• Hydrochloric acid (HCl) makes chlorides
• Sulfuric acid (H₂SO₄) makes sulfates
• Nitric acid (HNO₃) makes nitrates

Making Salts

• Soluble salts: react acid with excess insoluble base/carbonate, filter off excess, evaporate
• Titration method: for soluble base + acid — use indicator to find exact volumes

Electrolysis

Electrolysis is the decomposition of a compound using electricity. The compound must be molten or dissolved so ions are free to move.

• Anode = positive electrode — negative ions (anions) are attracted here and oxidised
• Cathode = negative electrode — positive ions (cations) are attracted here and reduced

Electrolysis of Molten Compounds

E.g. Molten lead bromide (PbBr₂):

• Cathode: Pb²⁺ + 2e⁻ → Pb (lead metal formed)
• Anode: 2Br⁻ → Br₂ + 2e⁻ (bromine gas formed)

Electrolysis of Solutions

Water provides H⁺ and OH⁻ ions. At GCSE level:

• At the cathode: if the metal is more reactive than hydrogen, hydrogen gas is produced. If less reactive, the metal is deposited.
• At the anode: if a halide ion is present, the halogen is produced. Otherwise, oxygen is produced.

Electroplating

Coating an object with a thin layer of metal using electrolysis. The object to be plated is the cathode. The plating metal is the anode. The electrolyte contains ions of the plating metal.

Electrolysis of Molten Compounds

When an ionic compound is melted, the ions are free to move. Only two ions are present, so the products are straightforward:

Electrolysis of Aqueous Solutions

In aqueous solutions, water (H₂O) also provides ions: H⁺ and OH⁻. This means there is competition at each electrode.

Rules for Products at Electrodes

Half Equations at Electrodes

At the cathode (reduction — gain of electrons):

At the anode (oxidation — loss of electrons):

Examples of Aqueous Electrolysis

Extraction of Aluminium

Aluminium is too reactive to be extracted by reduction with carbon. Instead, it is extracted by electrolysis of molten aluminium oxide (Al₂O₃) dissolved in cryolite (Na₃AlF₆).

• Cryolite lowers the melting point of aluminium oxide from ~2050 °C to ~950 °C, saving energy and cost
• Cathode (carbon): Al³⁺ + 3e⁻ → Al (molten aluminium sinks to the bottom and is tapped off)
• Anode (carbon): 2O²⁻ → O₂ + 4e⁻ (oxygen gas produced)
• The carbon anodes burn away (react with oxygen at high temperature: C + O₂ → CO₂) and must be regularly replaced
• The process is expensive due to the high energy requirement (large electric current needed)

Electroplating

Electroplating is coating an object with a thin layer of metal using electrolysis.

• Cathode: the object to be plated
• Anode: the plating metal (it dissolves to replace metal ions in solution)
• Electrolyte: a solution containing ions of the plating metal

Reasons for electroplating: appearance (silver/gold plating for jewellery), corrosion resistance (chromium plating on steel), hardness (chromium plating on tools).

Industrial Uses of Electrolysis

Collision Theory

For a reaction to occur, particles must collide with sufficient energy. The minimum energy needed is the activation energy. The rate of reaction depends on the frequency of successful collisions.

Factors Affecting Rate

Catalyst: a substance that speeds up a reaction without being used up. It is not consumed and can be recovered unchanged.

Exothermic Reactions

• Transfer energy to the surroundings (temperature increases)
• Examples: combustion, neutralisation, oxidation, hand warmers
• Energy released when bonds are formed is greater than energy needed to break bonds

Endothermic Reactions

• Take in energy from the surroundings (temperature decreases)
• Examples: thermal decomposition, photosynthesis, cold packs (ammonium nitrate dissolving)
• Energy needed to break bonds is greater than energy released when bonds are formed

Energy Profiles

Exothermic: products are at a lower energy level than reactants. Energy change is negative.

Endothermic: products are at a higher energy level than reactants. Energy change is positive.

Both profiles show an activation energy hump that must be overcome for the reaction to proceed.

Bond Energies

Breaking bonds = endothermic (energy IN). Making bonds = exothermic (energy OUT).

If the answer is negative → exothermic. If positive → endothermic.

Reversible Reactions

A reversible reaction is one that can go in both directions — products can react to re-form the original reactants. The symbol ⇌ is used instead of →.

Example: Heating hydrated copper sulfate:

The forward reaction is endothermic (heating drives off water). The reverse reaction is exothermic (adding water releases heat). In any reversible reaction, if the forward reaction is exothermic, the reverse is endothermic by the same amount, and vice versa.

Dynamic Equilibrium

When a reversible reaction occurs in a closed system (nothing can enter or leave), it reaches dynamic equilibrium:

• The rate of the forward reaction equals the rate of the reverse reaction
• The concentrations of reactants and products remain constant (but not necessarily equal)
• Both reactions are still occurring — it is dynamic, not static

Le Chatelier's Principle

Le Chatelier's principle: if a system at equilibrium is subjected to a change, the equilibrium will shift to oppose that change.

The Haber Process

The industrial manufacture of ammonia:

Conditions chosen as a compromise:

Unreacted nitrogen and hydrogen are recycled back through the reactor to improve overall yield.

The Contact Process

The industrial manufacture of sulfuric acid (H₂SO₄). The key equilibrium step is:

Conditions: temperature of 450 °C, 1-2 atm pressure, vanadium(V) oxide (V₂O₅) catalyst. The SO₃ produced is then dissolved in concentrated H₂SO₄ and diluted with water to form more sulfuric acid.

Definitions of Oxidation and Reduction

Oxidation and reduction always happen together — these are called redox reactions.

Oxidation States (Oxidation Numbers)

Oxidation states are a way of tracking electron transfer. Rules:

• Elements in their natural state have an oxidation state of 0 (e.g. O₂, Fe, Na)
• Simple ions have an oxidation state equal to their charge (e.g. Na⁺ = +1, Cl⁻ = -1)
• Oxygen is usually -2 in compounds
• Hydrogen is usually +1 in compounds
• The sum of oxidation states in a compound = 0
• The sum in a polyatomic ion = the charge on the ion

If an element's oxidation state increases, it has been oxidised. If it decreases, it has been reduced.

Identifying Redox Reactions

Example: Mg + CuO → MgO + Cu

• Magnesium goes from 0 to +2 (oxidation state increases) — it is oxidised (loses electrons)
• Copper goes from +2 to 0 (oxidation state decreases) — it is reduced (gains electrons)
• Magnesium gains oxygen — oxidised. Copper loses oxygen — reduced.

Displacement Reactions as Redox

All displacement reactions are redox reactions:

• Zinc is oxidised: Zn → Zn²⁺ + 2e⁻ (loses electrons, oxidation state 0 → +2)
• Copper is reduced: Cu²⁺ + 2e⁻ → Cu (gains electrons, oxidation state +2 → 0)

A more reactive metal will displace a less reactive metal from a solution of its salt. This is because the more reactive metal is a better electron donor (more easily oxidised).

Reactivity Series and Redox

Rusting as a Redox Process

Rusting is the corrosion of iron. It requires both oxygen and water.

• Iron is oxidised: Fe → Fe³⁺ + 3e⁻
• Oxygen is reduced: O₂ + 2H₂O + 4e⁻ → 4OH⁻

Prevention of Rusting

Kinetic Theory of Gases

In a gas, particles are in constant random motion. They collide with each other and with the walls of the container. The pressure of a gas is caused by particles colliding with the container walls.

• Increasing temperature makes particles move faster → more frequent and harder collisions → pressure increases
• Decreasing volume means particles are closer together → more frequent collisions → pressure increases
• Increasing the number of particles increases the frequency of collisions → pressure increases

Boyle's Law (constant temperature)

At constant temperature, the pressure of a gas is inversely proportional to its volume.

If you double the pressure, the volume halves (and vice versa).

Charles's Law (constant pressure)

At constant pressure, the volume of a gas is directly proportional to its absolute temperature (in Kelvin).

If you double the temperature (in Kelvin), the volume doubles.

Pressure Law (constant volume)

At constant volume, the pressure of a gas is directly proportional to its absolute temperature (in Kelvin).

Combined Gas Law

When temperature, pressure, and volume all change:

Summary of Gas Laws

Relative atomic mass (Ar) is the average mass of an atom of an element compared to 1/12 of a carbon-12 atom. Found on the periodic table.

Relative formula mass (Mr) is the sum of all the Ar values in a formula.

A mole is the amount of substance that contains 6.02 × 10²³ particles (Avogadro's number).

The empirical formula is the simplest whole number ratio of atoms of each element in a compound.

Steps

1. Write the mass (or percentage) of each element
2. Divide each by the Ar of that element (gives moles)
3. Divide all values by the smallest value
4. If needed, multiply to get whole numbers

Method

1. Write the balanced equation
2. Calculate moles of the substance you know
3. Use the mole ratio from the equation to find moles of the substance you want
4. Convert moles back to mass

Titration Method

1. Measure a known volume of alkali into a conical flask using a pipette
2. Add a few drops of indicator
3. Add acid from a burette until the indicator changes colour (end point)
4. Record the volume of acid used (titre)
5. Repeat until concordant results (within 0.10 cm³)

Titration Calculation Steps

1. Calculate moles of the substance you know: moles = concentration × volume
2. Use the mole ratio from the balanced equation
3. Calculate the unknown concentration or volume

Percentage Yield

Yield is always less than 100% because of: incomplete reactions, side reactions, loss during transfer/filtration/evaporation.

Atom Economy

High atom economy = less waste. Addition reactions have 100% atom economy. Industry prefers high atom economy for cost and environmental reasons.

At room temperature and pressure (RTP) — 20 °C and 1 atmosphere — one mole of any gas occupies 24 dm³ (24 000 cm³).

Crude oil is a fossil fuel formed from the remains of ancient marine organisms over millions of years. It is a mixture of hydrocarbons (compounds containing only hydrogen and carbon).

Fractional Distillation

Crude oil is separated into fractions by fractional distillation:

1. Crude oil is heated until it evaporates
2. Vapours enter the fractionating column (hot at bottom, cool at top)
3. Different fractions condense at different heights based on their boiling points
4. Shorter chain hydrocarbons rise higher (lower boiling point)

Trends in Hydrocarbons

Uses of Fractions

• Gases (LPG): heating, cooking
• Petrol/gasoline: fuel for cars
• Kerosene: jet fuel
• Diesel: fuel for lorries, buses
• Fuel oil: ships, power stations
• Bitumen: roads, roofing

Alkanes are a family of saturated hydrocarbons (all single bonds, no double bonds).

First Four Alkanes

Combustion of Alkanes

Complete combustion requires plenty of oxygen. Incomplete combustion occurs with limited oxygen and produces toxic carbon monoxide and/or soot (carbon particles).

Alkenes are unsaturated hydrocarbons containing at least one carbon-carbon double bond (C=C).

First Three Alkenes

Test for Alkenes

Add bromine water (orange/brown). If it decolourises (turns colourless), an alkene (C=C double bond) is present. This is an addition reaction.

Addition Reactions

• + Bromine: ethene + Br₂ → dibromoethane (test for unsaturation)
• + Hydrogen: ethene + H₂ → ethane (hydrogenation, using a nickel catalyst)
• + Water (steam): ethene + H₂O → ethanol (hydration, using phosphoric acid catalyst)

Addition polymerisation: many small alkene molecules (monomers) join together to form a long chain (polymer). The double bond opens up to form single bonds linking the monomers.

Common Polymers and Uses

Environmental Issues

• Most polymers are non-biodegradable — they persist in the environment for hundreds of years
• Burning plastics releases toxic gases (e.g. HCl from PVC)
• Plastics in oceans harm marine wildlife
• Solutions: recycling, developing biodegradable polymers, reducing plastic use

Cracking is the process of breaking down long-chain hydrocarbons into shorter, more useful molecules. It produces alkanes (for fuels) and alkenes (for making polymers and other chemicals).

Why Is Cracking Needed?

• Fractional distillation produces more long-chain fractions than there is demand for
• There is high demand for short-chain fuels (petrol) and alkenes (for polymers)
• Cracking converts surplus long-chain molecules into these high-demand products

Thermal Cracking

• Carried out at high temperatures (700-1000 °C) and high pressures
• Produces a high proportion of alkenes

Catalytic Cracking

• Uses a hot catalyst (e.g. zeolite or aluminium oxide) at about 500 °C
• Lower temperature than thermal cracking
• Produces more branched alkanes and aromatic hydrocarbons (used in fuels)

Example

The products always include at least one alkene (with a C=C double bond). You can verify cracking has occurred by testing the products with bromine water — it will decolourise if alkenes are present.

A homologous series is a family of compounds with the same functional group, the same general formula, and similar chemical properties. Each member differs by CH₂ from the next.

Properties of a Homologous Series

• Same general formula
• Same functional group — determines chemical properties
• Similar chemical properties
• Gradual change in physical properties (e.g. boiling points increase as chain length increases)
• Each member differs from the next by CH₂

Key Homologous Series

Alcohols

Alcohols contain the -OH functional group. The first three are methanol (CH₃OH), ethanol (C₂H₅OH), and propanol (C₃H₇OH).

• Dissolve in water to form neutral solutions
• React with sodium to produce hydrogen
• Burn in air to produce CO₂ + H₂O (used as fuels)
• Used as solvents and in alcoholic drinks (ethanol)

Carboxylic Acids

Carboxylic acids contain the -COOH functional group. They are weak acids (only partially ionised in water).

• React with metals: e.g. 2CH₃COOH + Mg → (CH₃COO)₂Mg + H₂
• React with carbonates: e.g. 2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂
• React with alcohols to form esters (see below)
• Examples: methanoic acid (HCOOH), ethanoic acid (CH₃COOH — found in vinegar)

Esters

Esters are formed when a carboxylic acid reacts with an alcohol in the presence of an acid catalyst (e.g. concentrated sulfuric acid). This is a condensation reaction (water is also produced).

• Esters have fruity/sweet smells
• Used in flavourings, perfumes, and solvents

There are two methods for producing ethanol:

1. Fermentation

Fermentation uses yeast enzymes to convert sugars (glucose) into ethanol and carbon dioxide:

Conditions required:

• Temperature of about 30-40 °C (optimum for yeast enzymes)
• Anaerobic conditions (no oxygen — otherwise ethanol is oxidised to ethanoic acid)
• Yeast as a biological catalyst

Advantages: uses renewable sugar crops, low temperature, low energy cost.

Disadvantages: slow process, produces dilute ethanol (needs distillation to concentrate), batch process.

2. Hydration of Ethene

Conditions: high temperature (300 °C), high pressure (60-70 atm), phosphoric acid catalyst.

Advantages: fast, continuous process, produces pure ethanol.

Disadvantages: uses non-renewable crude oil (ethene from cracking), high energy cost.

Addition Polymerisation (Recap)

• Monomers must have a C=C double bond (alkenes)
• The double bond opens up and monomers join together
• Only one product (the polymer) — no by-product
• 100% atom economy

Condensation Polymerisation

• Monomers have two functional groups (one at each end of the molecule)
• Monomers join together with the loss of a small molecule (usually water, H₂O)
• Two products: the polymer and the small molecule
• Examples: polyesters (from dicarboxylic acid + diol), polyamides/nylon (from dicarboxylic acid + diamine), proteins (from amino acids)

Current Composition of the Atmosphere

• Nitrogen: ~78%
• Oxygen: ~21%
• Argon: ~0.9%
• Carbon dioxide: ~0.04%
• Plus small amounts of water vapour and other gases

Evolution of Earth's Atmosphere

1. Early atmosphere (4 billion years ago): mainly CO₂ and water vapour from volcanic activity. Little or no oxygen. Similar to Mars and Venus today.
2. Water vapour condensed to form oceans. CO₂ dissolved in the oceans.
3. Photosynthetic organisms (cyanobacteria, then algae and plants) evolved and produced oxygen:
   6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
4. Oxygen built up, CO₂ decreased. Marine organisms locked CO₂ into shells (formed limestone/chalk). Dead organisms formed fossil fuels over millions of years.

Greenhouse Effect and Climate Change

Greenhouse gases (CO₂, methane, water vapour) absorb infrared radiation re-emitted from Earth's surface and re-radiate it in all directions, warming the atmosphere.

Human activities increasing greenhouse gases: burning fossil fuels (CO₂), deforestation (less CO₂ absorbed), agriculture and landfill (methane).

Consequences: rising global temperatures, melting ice caps, rising sea levels, extreme weather, habitat loss.

Water Treatment

1. Sedimentation: large particles settle out
2. Filtration: smaller particles removed by passing through sand beds
3. Chlorination: chlorine added to kill bacteria and make water safe to drink

Distillation can produce pure water but is too expensive and energy-intensive for large-scale use.

Testing for Gases

Balance each equation. Click to reveal the answer.

1. Saying ionic compounds contain "molecules" — they form giant ionic lattices, not molecules.
2. Saying "covalent bonds are weak" — the bonds are strong! It is the intermolecular forces that are weak in simple molecules.
3. Changing subscript numbers when balancing — only change the big numbers (coefficients) in front of formulae.
4. Forgetting to convert cm³ to dm³ — divide by 1000 before using concentration formulae.
5. Saying temperature "gives particles more energy to react" — say that more particles exceed the activation energy.
6. Mixing up Group 1 and Group 7 reactivity trends — metals get more reactive going down, halogens get less reactive.
7. Confusing oxidation and reduction — remember OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons).
8. Writing "electrons are shared" in ionic bonding — in ionic bonding, electrons are transferred, not shared.
9. Saying a catalyst "gives energy" to a reaction — a catalyst lowers the activation energy by providing an alternative pathway.
10. Confusing greenhouse effect with ozone layer depletion — these are completely separate environmental issues.
11. Forgetting state symbols — many CCEA mark schemes require (s), (l), (g), (aq).
12. Saying "atoms want to have a full outer shell" — atoms do not have desires. Say atoms are "more stable" with a full outer shell.