GCSE Chemistry

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Revision Topics

• Topic 1: Atomic Structure & Periodic Table 5
  • 1.1 Atoms, Elements & Compounds 5
  • 1.2 Separation Techniques 6
  • 1.3 Chromatography & Rf Values 7
  • 1.4 Changing Model of the Atom 7
  • 1.5 Subatomic Particles 8
  • 1.6 Atomic Size & Mass Distribution 8
  • 1.7 Isotopes & Relative Atomic Mass 9
  • 1.8 Electron Configuration 10
  • 1.9 The Periodic Table 10
  • 1.10 Groups in the Periodic Table 10
• Topic 2: Bonding, Structure & Properties 12
  • 2.1 Chemical Bonds 12
  • 2.2 Ionic Bonding 13
  • 2.3 Giant Ionic Lattice 13
  • 2.4 Covalent Bonding 14
  • 2.5 Small Molecule Covalent Substances 14
  • 2.6 Polymers 15
  • 2.7 Giant Covalent Structures 15
  • 2.8 Graphene & Fullerenes Chemistry Only 16
  • 2.9 Metallic Bonding 16
  • 2.10 Alloys 16
  • 2.11 Three States of Matter 17
  • 2.12 Nanoparticles Chemistry Only 17
• Topic 3: Quantitative Chemistry 18
  • 3.1 Conservation of Mass 18
  • 3.2 Relative Formula Mass (Mr) 18
  • 3.3 Moles Chemistry Only 19
  • 3.4 Amounts in Equations Chemistry Only 19
  • 3.5 Limiting Reactants Chemistry Only 20
  • 3.6 Concentration of Solutions 20
  • 3.7 Percentage Yield Chemistry Only 20
  • 3.8 Atom Economy Chemistry Only 21
  • 3.9 Volume of Gases Chemistry Only 21
• Topic 4: Chemical Changes 22
  • 4.1 Metal Oxides 22
  • 4.2 The Reactivity Series 22
  • 4.3 Extraction of Metals 22
  • 4.4 Reactions of Acids 23
  • 4.5 Soluble Salts Practical 23
  • 4.6 The pH Scale 24
  • 4.7 Strong & Weak Acids Chemistry Only 24
  • 4.8 Electrolysis 24
  • 4.9 Aqueous Electrolysis Chemistry Only 25
  • 4.10 Half Equations Chemistry Only 25
• Topic 5: Energy Changes 26
  • 5.1 Exothermic & Endothermic Reactions 26
  • 5.2 Reaction Profiles 26
  • 5.3 Bond Energy Calculations Chemistry Only 27
  • 5.4 Chemical Cells & Batteries Chemistry Only 27
  • 5.5 Fuel Cells Chemistry Only 28
• Topic 6: Rate & Extent of Chemical Change 29
  • 6.1 Calculating Rates of Reaction 29
  • 6.2 Collision Theory 29
  • 6.3 Factors Affecting Rate 30
  • 6.4 Catalysts 30
  • 6.5 Reversible Reactions 31
  • 6.6 Equilibrium Chemistry Only 32

• Topic 7: Organic Chemistry 33
  • 7.1 Crude Oil & Hydrocarbons 33
  • 7.2 Fractional Distillation 34
  • 7.3 Combustion of Hydrocarbons 35
  • 7.4 Cracking & Alkenes 35
  • 7.5 Reactions of Alkenes Chemistry Only 36
  • 7.6 Addition Polymerisation Chemistry Only 36
  • 7.7 Alcohols Chemistry Only 36
  • 7.8 Carboxylic Acids Chemistry Only 37
  • 7.9 Condensation Polymerisation Chemistry Only 37
  • 7.10 Natural Polymers Chemistry Only 37
• Topic 8: Chemical Analysis 38
  • 8.1 Pure Substances & Mixtures 38
  • 8.2 Formulations 38
  • 8.3 Chromatography 38
  • 8.4 Gas Tests 39
  • 8.5 Cation Tests (Flame) Chemistry Only 39
  • 8.6 Cation Tests (Precipitate) Chemistry Only 40
  • 8.7 Anion Tests Chemistry Only 40
  • 8.8 Instrumental Methods Chemistry Only 40
• Topic 9: Chemistry of the Atmosphere 41
  • 9.1 The Atmosphere Today 41
  • 9.2 The Earth's Early Atmosphere 42
  • 9.3 How Oxygen Increased 42
  • 9.4 How Carbon Dioxide Decreased 42
  • 9.5 Greenhouse Gases 43
  • 9.6 Pollutants from Fuels 43
• Topic 10: Using Resources 44
  • 10.1 Sustainable Development 44
  • 10.2 Potable Water 44
  • 10.3 Waste Water Treatment 45
  • 10.4 Alternative Metal Extraction Chemistry Only 45
  • 10.5 Life Cycle Assessment (LCA) 46
  • 10.6 Reduce, Reuse, Recycle 46
  • 10.7 Corrosion Chemistry Only 46
  • 10.8 Alloys & Materials Chemistry Only 47
  • 10.9 The Haber Process Chemistry Only 47
• Quick-Check Answers
  • Topic 1 Answers 48
  • Topic 2 Answers 49
  • Topic 3 Answers 50
  • Topic 4 Answers 51
  • Topic 5 Answers 52
  • Topic 6 Answers 53
  • Topic 7 Answers 54
  • Topic 8 Answers 55
  • Topic 9 Answers 56
  • Topic 10 Answers 57
• Glossary of Key Terms 58
  Required Practicals Hub 59
  Periodic Table Insert 60
• Practice Exams
  • Paper 1 - Higher Tier (Topics 1-5) 61
  • Paper 2 - Higher Tier (Topics 6-10) 78
  • Mark Scheme - Paper 1 95
  • Mark Scheme - Paper 2 111

1.1 Atoms, Elements & Compounds

All substances are made of atoms. An atom is the smallest part of an element that can exist.

Compounds can only be separated by chemical reactions, involving new substances and energy changes.

Chemical Equations

Word equations name reactants (left) and products (right). Balanced symbol equations show atom counts satisfying conservation of mass.

1.2 Separation Techniques

1.3 Chromatography & Rf Values

Chromatography separates mixtures based on relative solubility in a mobile phase. The R_f value shows this relationship:

1.4 Changing Model of the Atom

The model of the atom has changed over time as new experimental evidence was discovered:

• Dalton (1803): Atoms proposed as solid, indivisible spheres.
• Thomson (1897): Discovered electron; proposed Plum Pudding model (positive sphere with embedded negative charges).
• Rutherford (1909): Nuclear model: positive nucleus surrounded by empty space.
• Bohr (1913): Electrons orbit nucleus in fixed shells.
• Chadwick (1932): Discovered the neutron.

1.5 Subatomic Particles

An atom consists of a central positive nucleus (protons and neutrons) surrounded by negative electrons in shells.

1.6 Atomic Size & Mass Distribution

Atoms are tiny. The atomic radius is approximately 1 × 10^-10 metres (0.1 nm). The radius of the nucleus is about 1 × 10^-14 metres (10,000 times smaller than the atom).

Nearly all of the atom's mass is concentrated in the nucleus, as protons and neutrons have a relative mass of 1, whereas electron mass is negligible.

1.7 Isotopes & Relative Atomic Mass

1.8 Electron Configuration

Electrons occupy energy levels (shells). Capacities: 1st shell = 2, 2nd shell = 8, 3rd shell = 8. Electrons fill lowest available shells first.

Configurations show the electron structure: Sodium (11 electrons) is 2.8.1; Chlorine (17 electrons) is 2.8.7.

1.9 The Periodic Table

Elements are arranged in order of atomic number. Early periodic tables were arranged by atomic weight. Mendeleev left gaps for undiscovered elements, predicting their properties. Metals (left/bottom) lose electrons to form positive ions. Non-metals (right/top) gain/share electrons.

1.10 Groups in the Periodic Table

Group 0: The Noble Gases

Helium, Neon, Argon. Full outer shells (stable electronic configuration). Monatomic and chemically inert. Boiling points increase down the group as relative atomic mass increases (stronger intermolecular forces require more energy to break).

Group 1: The Alkali Metals

Li, Na, K. 1 outer electron. Reactivity increases down the group as the outer electron is further from the nucleus, more shielded by inner shells, and lost more easily. React with water to form metal hydroxide and hydrogen:

Group 7: The Halogens

F, Cl, Br, I. 7 outer electrons. Diatomic molecules (X₂). Reactivity decreases down the group as it is harder to attract an incoming electron due to increased distance and shielding. Boiling points increase down the group.

A more reactive halogen displaces a less reactive halide from its salt solution:

Observation: the solution turns orange-brown as bromine is displaced.

Transition Metals Chemistry Only

Central block of the periodic table. Compared to Group 1 metals: harder, denser, stronger, higher melting points, and much less reactive.

Coloured compounds:

• Cu²⁺ ions: blue (e.g. copper sulfate solution)
• Fe³⁺ ions: yellow-brown (e.g. iron(III) chloride solution)
• Fe²⁺ ions: green (e.g. iron(II) sulfate solution)

Variable oxidation states: most transition metals form ions with different charges, e.g. iron forms Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)).

Catalytic activity:

• Iron is the catalyst in the Haber process (making ammonia).
• Vanadium(V) oxide (V₂O₅) is the catalyst in the Contact process (making sulfuric acid).

2.1 Chemical Bonds

Compounds form when atoms join by chemical bonds. Atoms gain, lose or share electrons to get a stable full outer shell:

• Ionic: Metal + non-metal. Electrons transferred.
• Covalent: Non-metals. Electrons shared.
• Metallic: Metals/alloys. Delocalised electrons.

2.2 Ionic Bonding

Metal atoms lose outer electrons to become positive ions (cations). Non-metal atoms gain these electrons to become negative ions (anions). They are held by strong electrostatic attractions.

2.3 Giant Ionic Lattice

Ionic compounds form a 3D giant ionic lattice of alternating positive and negative ions held by strong, multi-directional electrostatic forces.

Properties:

• High melting/boiling points: Large amount of energy needed to break strong electrostatic attractions between oppositely charged ions.
• Electrical conductivity: Do not conduct when solid (ions fixed in place). Conduct when molten or dissolved in water because ions are free to move and carry charge.

2.4 Covalent Bonding

Covalent bonding occurs when two non-metal atoms share one or more pairs of electrons to achieve stable full outer shells. These bonds are very strong.

Covalent substances can consist of small molecules (like water) or giant covalent structures (like diamond).

2.5 Small Molecule Covalent Substances

Substances with small covalent molecules (e.g. H₂O, CO₂, CH₄, NH₃) have very low melting and boiling points.

• Bonding structure: Atoms within the molecules are joined by very strong covalent bonds. However, the forces between separate molecules (intermolecular forces) are very weak.
• Melting/boiling: When these substances melt or boil, it is only the weak intermolecular forces that are broken, not the strong covalent bonds. Very little energy is needed.
• Conductivity: Do not conduct electricity because the molecules are neutral (no overall charge) and have no free delocalised electrons or ions to carry charge.

2.6 Polymers

Polymers consist of very large molecules made of long chains of repeating units (monomers) joined by strong covalent bonds.

• Intermolecular forces: Because polymer molecules are very large, the sum of the intermolecular forces between the chains is relatively strong.
• State: This makes polymers solid at room temperature. However, they melt at lower temperatures than giant ionic or giant covalent structures because these forces are still weaker than chemical bonds.

2.7 Giant Covalent Structures

Giant covalent structures (macromolecules) contain huge networks of atoms joined by strong covalent bonds. They have very high melting/boiling points because breaking these structures requires breaking many strong covalent bonds.

• Diamond: Each carbon atom forms 4 strong covalent bonds in a rigid tetrahedral structure. This makes diamond extremely hard and unable to conduct electricity (no free electrons).
• Graphite: Each carbon forms 3 bonds, creating hexagonal layers. The fourth electron is delocalised and free to move along the layers, so graphite conducts electricity and heat. Weak forces between layers let them slide, making graphite soft and slippery.
• Silicon dioxide (Silica, SiO₂): Similar structure to diamond. Each silicon atom is bonded to 4 oxygen atoms, and each oxygen to 2 silicon atoms. Extremely high melting point and hard.

2.8 Graphene & Fullerenes Chemistry Only

• Graphene: A single 2D layer of graphite, one atom thick. Very strong, light, and conduct electricity. Used in electronics and composite materials.
• Fullerenes: Carbon molecules with hollow shapes (e.g. Buckminsterfullerene, C₆₀, a sphere). Used for drug delivery, lubricants, and catalysts.
• Carbon nanotubes: Cylindrical fullerenes. High tensile strength and electrical conductivity. Used in nanotechnology and structural materials.

2.9 Metallic Bonding

Metals consist of a giant lattice of positive metal ions surrounded by a "sea" of delocalised electrons. This is held together by strong electrostatic attractions.

Properties: High melting/boiling points. Delocalised electrons carry thermal energy and electrical charge, so metals are excellent conductors. Pure metals are malleable because layers of ions can slide over each other.

2.10 Alloys

Alloys are mixtures of a metal with other elements. Pure metals are soft because their atoms are arranged in regular layers that slide easily. In alloys, different-sized atoms distort these layers, making it harder for them to slide, which makes alloys harder than pure metals.

2.11 Three States of Matter

The three states of matter are solid, liquid, and gas. State changes depend on the strength of forces between particles, which is determined by the bonding and structure of the substance.

2.12 Nanoparticles Chemistry Only

Nanoparticles have a diameter of 1 to 100 nm. They have an extremely high surface area to volume ratio, giving them properties that are different from the bulk material.

• SA:V Ratio: As particles decrease in size, their surface area to volume ratio increases dramatically, making them highly reactive.
• Uses: Sun creams (ZnO nanoparticles block UV, transparent on skin), catalysts (large surface area), medicine (targeted drug delivery), and electronics.
• Risks: May enter the body and cells (toxic effects), or accumulate in environments. Long-term risks are not yet fully understood.

3.1 Conservation of Mass

No atoms are lost or made during a chemical reaction, so the total mass of the products is always equal to the total mass of the reactants. This is why symbol equations must be balanced.

Apparent mass changes: If the mass seems to increase, a gaseous reactant has joined (e.g. a metal reacting with oxygen to form an oxide). If it seems to decrease, a gas has escaped (e.g. thermal decomposition of a metal carbonate releasing CO₂).

3.2 Relative Formula Mass (M_r)

The relative formula mass (M_r) of a compound is the sum of the relative atomic masses (A_r) of the atoms in the numbers shown in the formula.

3.3 Moles Chemistry Only

Chemical amounts are measured in moles (mol). One mole of any substance contains 6.02 × 10²³ particles (Avogadro's constant). The mass of one mole of a substance in grams is equal to its relative formula mass.

3.4 Amounts in Equations Chemistry Only

Balanced chemical equations show the ratio of moles that react together. We can use these ratios to calculate the mass of a product formed or reactant needed.

3.5 Limiting Reactants Chemistry Only

In a reaction, the reactant that is completely used up is the limiting reactant. Any other reactants are in excess. The amount of product formed is directly proportional to the amount of limiting reactant used.

3.6 Concentration of Solutions

Concentration tells us how much solute is dissolved in a volume of solvent. It can be expressed in g/dm³ or mol/dm³ (where 1 dm³ = 1000 cm³).

3.7 Percentage Yield Chemistry Only

The yield is the amount of product obtained. Percentage yield compares actual yield to maximum theoretical yield:

Why yield is less than 100%: Reaction may be reversible; some product lost in separation/transfer; side reactions may occur; reactants may not be pure.

3.8 Atom Economy Chemistry Only

Atom economy measures the proportion of starting materials that become useful products. High atom economy reduces waste and is more sustainable.

3.9 Volume of Gases Chemistry Only

One mole of any gas occupies 24 dm³ at room temperature and pressure (RTP, 20 °C and 1 atm pressure).

4.1 Metal Oxides

Metals react with oxygen to form metal oxides. This is an oxidation reaction because the metal gains oxygen.

4.2 The Reactivity Series

Metals are arranged in order of reactivity based on their reactions with water and dilute acids. A more reactive metal displaces a less reactive metal from its compound.

4.3 Extraction of Metals

Unreactive metals (like gold) are found as pure elements. Metals less reactive than carbon are extracted from their oxides by heating with carbon (carbon reduces the metal oxide by removing oxygen):

Metals more reactive than carbon must be extracted using electrolysis, which requires large amounts of energy.

4.4 Reactions of Acids

Acids react with metals and bases (metal oxides, metal hydroxides, and metal carbonates) to produce salts.

Hydrochloric acid produces chlorides, sulfuric acid produces sulfates, and nitric acid produces nitrates.

4.5 Soluble Salts Practical

The ionic equation for any neutralisation reaction between an acid and an alkali is:

4.6 The pH Scale

The pH scale (0 to 14) measures acidity or alkalinity. Acids produce H⁺ ions in solution; alkalis produce OH⁻ ions in solution. A pH of 7 is neutral.

4.7 Strong & Weak Acids Chemistry Only

• Strong acids: Completely ionise/dissociate in aqueous solution, releasing all H⁺ ions (e.g. HCl, HNO₃, H₂SO₄).
• Weak acids: Only partially ionise in aqueous solution (e.g. ethanoic, citric, carbonic acids). The reaction is reversible:
CH₃COOH(aq) ⇌ CH₃COO^-(aq) + H⁺(aq)
• pH scale relationship: As pH decreases by 1 unit, the hydrogen ion concentration increases by a factor of 10.
• Strength vs Concentration: Acid strength refers to the degree of ionisation. Acid concentration refers to the mass of acid dissolved in a given volume of solution. A strong acid can be dilute, and a weak acid can be concentrated.

4.8 Electrolysis

Electrolysis is the decomposition of an electrolyte using electricity. Ions move to electrodes: positive ions (cations) move to the cathode (negative electrode) where they gain electrons (reduction). Negative ions (anions) move to the anode (positive electrode) where they lose electrons (oxidation).

4.9 Aqueous Electrolysis Chemistry Only

In aqueous solutions, water molecules ionise into H⁺ and OH⁻ ions alongside the compound's ions. Rules for products:

• Cathode (-): Hydrogen is produced if the metal is more reactive than hydrogen. The metal is produced only if it is less reactive than hydrogen (e.g. copper).
• Anode (+): Oxygen is produced (from OH⁻) unless halide ions (Cl⁻, Br⁻, I⁻) are present, in which case the halogen is formed.

4.10 Half Equations Chemistry Only

Write half equations to show reactions at electrodes (e.g. electrolysis of aqueous NaCl):

5.1 Exothermic & Endothermic Reactions

Energy is conserved in chemical reactions. The total energy before a reaction is equal to the total energy after.

• Exothermic: Transfers energy to surroundings (surrounding temperature increases). Examples: combustion, neutralisation, hand warmers.
• Endothermic: Takes in energy from surroundings (surrounding temperature decreases). Examples: thermal decomposition, citric acid + sodium hydrogencarbonate, sports injury cold packs.

5.2 Reaction Profiles

Reaction profiles show the relative energies of reactants and products, the activation energy, and the overall energy change.

5.3 Bond Energy Calculations Chemistry Only

Chemical reactions involve breaking bonds in reactants (takes in energy, endothermic) and forming new bonds in products (releases energy, exothermic).

If the result is negative, the reaction is exothermic. If positive, it is endothermic.

5.4 Chemical Cells & Batteries Chemistry Only

A simple chemical cell consists of two different metals in an electrolyte. The difference in reactivity creates a potential difference (voltage). A battery contains two or more cells connected in series.

• Non-rechargeable cells: The chemical reactions stop when one of the reactants is used up (e.g. alkaline batteries).
• Rechargeable cells: The chemical reactions can be reversed by connecting the cell to an external electrical current.

5.5 Fuel Cells Chemistry Only

Fuel cells are supplied by an external source of fuel (e.g. hydrogen) and oxygen. The fuel is oxidised electrochemically to produce a voltage continuously.

The overall reaction in a hydrogen fuel cell is the combustion of hydrogen to form water:

Half equations (alkaline electrolyte):

Advantages and Disadvantages

Fuel Cells vs Rechargeable Batteries

6.1 Calculating Rates of Reaction

The rate of a chemical reaction tells us how fast reactants turn into products. It can be calculated as:

Units are usually g/s, cm³/s, or mol/s. On a graph of amount against time, the rate at any moment equals the gradient of the tangent at that point. The reaction is fastest at the start and the line levels off when a reactant is used up.

Measuring rate (required practical)

• Gas given off: collect it with a gas syringe and measure the volume over time.
• Mass change: stand the flask on a balance and record the mass lost as gas escapes.
• Cloudiness (turbidity): time how long a cross marked under the flask takes to disappear (e.g. sodium thiosulfate + hydrochloric acid).

6.2 Collision Theory

Chemical reactions can only occur when reacting particles collide with each other with sufficient energy (the activation energy).

6.3 Factors Affecting Rate

• Temperature: Particles move faster, colliding more frequently and with more energy. A higher proportion of collisions have energy ≥ activation energy.
• Concentration / Pressure: More particles in a given volume, increasing collision frequency.
• Surface Area: Breaking a solid reactant into smaller pieces increases exposed surface area, increasing collision frequency.

6.4 Catalysts

A catalyst changes the rate of a chemical reaction but is not used up. It provides an alternative pathway with a lower activation energy, increasing the proportion of successful collisions.

6.5 Reversible Reactions

In a reversible reaction, products can react to reform the original reactants. Shown by the double arrow (⇌).

If a reversible reaction is exothermic in one direction, it is endothermic in the opposite direction. The same amount of energy is transferred in both directions.

Key Examples

1. Ammonium chloride:

On heating, white ammonium chloride decomposes. On cooling, the gases recombine and white solid reforms higher up the tube.

2. Hydrated copper sulfate:

Forward (heating): blue crystals turn to white powder (endothermic). Reverse (add water): white powder turns blue (exothermic).

6.6 Equilibrium Chemistry Only

In a closed system, dynamic equilibrium is reached when the forward and reverse reactions occur at exactly the same rate, and the concentrations of reactants and products remain constant.

Le Chatelier's Principle

If a system at equilibrium is subjected to a change in conditions, the system will adjust to counteract the change:

• Temperature: If temperature is increased, the system moves in the endothermic direction to absorb heat. If decreased, it moves in the exothermic direction.
• Pressure (gases): If pressure is increased, the system shifts to the side with fewer gas molecules to reduce pressure. If decreased, it shifts to the side with more molecules.
• Concentration: If concentration of a reactant is increased, more products form to use it up.

7.1 Crude Oil & Hydrocarbons

Crude oil is a finite resource found in rocks, formed over millions of years from the remains of ancient biomass (mainly plankton). It is a mixture of hydrocarbons.

Hydrocarbons are compounds containing hydrogen and carbon atoms only. Most are alkanes, which are saturated hydrocarbons (each carbon has single covalent bonds only).

The alkanes are a homologous series: a family of compounds with the same general formula, where each member differs from the next by CH₂. Members have similar chemical properties and show a gradual trend in physical properties as the chain lengthens.

7.2 Fractional Distillation

Crude oil is separated into fractions (groups of hydrocarbons with similar boiling points) by fractional distillation. The column is hot at the bottom and cold at the top.

Properties of hydrocarbons:

• Short-chain: Low boiling point, low viscosity (runny), highly flammable. Condense at the cool top.
• Long-chain: High boiling point, high viscosity (thick), low flammability. Condense at the hot bottom.

7.3 Combustion of Hydrocarbons

Combustion releases energy. The carbon and hydrogen in the fuel are oxidised.

• Complete: Plenty of oxygen. Products: carbon dioxide and water.
propane + oxygen → carbon dioxide + water
C₃H₈ + 5O₂ → 3CO₂ + 4H₂O
• Incomplete: Limited oxygen. Products: carbon monoxide (CO), soot (carbon), and water.

7.4 Cracking & Alkenes

Long-chain hydrocarbons are cracked into shorter, more useful alkanes and alkenes. Methods: catalytic cracking (high temperature, zeolite catalyst) and steam cracking (higher temperature, steam).

Alkenes are unsaturated hydrocarbons containing a carbon-carbon double bond (C=C).

Test for Alkenes: Orange bromine water is decolourised (goes colourless) when mixed with an alkene (unsaturated), but stays orange with an alkane (saturated).

7.5 Reactions of Alkenes Chemistry Only

Alkenes are highly reactive due to the C=C bond, reacting via addition reactions:

• With halogens: E.g. ethene + bromine → dibromoethane.
• With hydrogen (hydrogenation): E.g. ethene + H₂ → ethane (requires 150 °C, nickel catalyst).
• With water (hydration): E.g. ethene + steam ⇌ ethanol (requires high temp, phosphoric acid catalyst).

7.6 Addition Polymerisation Chemistry Only

Alkenes act as monomers. In addition polymerisation, many small alkene molecules join together to form polymers. The double bond opens up to form single bonds between monomers. Only one product forms.

7.7 Alcohols Chemistry Only

Alcohols contain the functional group -OH. The first four are methanol, ethanol, propanol, and butanol.

• Reactions: Burn in air (combustion); react with sodium (releases hydrogen); oxidise to form carboxylic acids.
• Fermentation: Glucose → ethanol + carbon dioxide (requires yeast, warm, anaerobic conditions).

7.8 Carboxylic Acids Chemistry Only

Carboxylic acids contain the functional group -COOH (e.g. ethanoic acid). They react with carbonates to produce a salt, water, and CO₂. They are weak acids (only partially ionise in solution).

React with alcohols (in presence of an acid catalyst) to form esters (functional group -COO-):

7.9 Condensation Polymerisation Chemistry Only

Involves monomers with two functional groups. When they react, they join together and lose small molecules such as water (unlike addition polymerisation which forms no by-products).

Example: Polyester from a dicarboxylic acid and a diol:

7.10 Natural Polymers Chemistry Only

• Amino Acids: Have two different functional groups (basic amine group -NH₂ and acidic carboxylic acid group -COOH). Join by condensation polymerisation to form polypeptides and proteins.
• DNA (Deoxyribonucleic acid): Large molecule essential for life. Consists of two polymer chains made from four different monomers (nucleotides) in a double helix.
• Starch & Cellulose: Polymers made from glucose monomers.

8.1 Pure Substances & Mixtures

In chemistry, a pure substance is a single element or compound, not mixed with any other substance.

Melting/Boiling Points: Pure substances melt and boil at specific, sharp temperatures. Impure substances (mixtures) melt and boil over a range of temperatures, and at lower melting points / higher boiling points than the pure substance.

8.2 Formulations

A formulation is a mixture that has been designed as a useful product. Every component is carefully measured to ensure it has the required properties.

Examples: fuels, cleaning agents, paints, medicines, alloys, fertilisers, and foods.

8.3 Chromatography

Paper chromatography separates mixtures based on their relative solubility in a mobile phase (solvent) and attraction to a stationary phase (paper).

A pure substance produces a single spot on a chromatogram. An impure substance produces multiple spots.

8.4 Gas Tests

8.5 Cation Tests (Flame) Chemistry Only

Identify metal ions (cations) using flame tests (nichrome wire dipped in HCl):

• Lithium (Li⁺): Crimson flame
• Sodium (Na⁺): Yellow flame
• Potassium (K⁺): Lilac flame
• Calcium (Ca²⁺): Orange-red flame
• Copper (Cu²⁺): Green flame

8.6 Cation Tests (Precipitate) Chemistry Only

Identify metal ions by adding sodium hydroxide (NaOH) solution:

• Al³⁺, Ca²⁺, Mg²⁺: Form white precipitates. Only Al(OH)₃ dissolves in excess NaOH to form a clear solution.
• Copper (Cu²⁺): Blue precipitate: Cu²⁺ + 2OH^- → Cu(OH)₂(s)
• Iron(II) (Fe²⁺): Green precipitate: Fe²⁺ + 2OH^- → Fe(OH)₂(s)
• Iron(III) (Fe³⁺): Brown precipitate: Fe³⁺ + 3OH^- → Fe(OH)₃(s)

8.7 Anion Tests Chemistry Only

• Carbonates (CO₃^2-): Add dilute acid. Fizzing (effervescence) occurs because CO₂ is released. Test with limewater.
• Halides (Cl^-, Br^-, I^-): Add dilute nitric acid and silver nitrate. Chloride forms a white precipitate (AgCl); Bromide forms cream (AgBr); Iodide forms yellow (AgI).
• Sulfates (SO₄^2-): Add dilute hydrochloric acid and barium chloride. A white precipitate of barium sulfate forms.

8.8 Instrumental Methods Chemistry Only

Instrumental methods are rapid, sensitive, and accurate. Flame emission spectroscopy is an instrumental method used to analyze metal ions in solution. The sample is heated in a flame, and the emitted light is analyzed through a spectroscope to produce a line spectrum.

9.1 The Atmosphere Today

For about 200 million years, the proportions of gases in the Earth's atmosphere have been roughly stable:

• Nitrogen (N₂): about 78%
• Oxygen (O₂): about 21%
• Other gases: argon (~0.9%), carbon dioxide (~0.04%), and small, variable amounts of water vapour.

The proportion of carbon dioxide is small but is now increasing because of human activity - mainly burning fossil fuels and deforestation (covered in 9.5 Greenhouse Gases).

9.2 The Earth's Early Atmosphere

The Earth was formed 4.6 billion years ago. Theories suggest:

1. Volcanic Activity: Intense volcanic activity released gases that formed the early atmosphere - mainly carbon dioxide, with little or no oxygen, water vapour, nitrogen, and small amounts of methane and ammonia.
2. Ocean Formation: The Earth cooled and water vapour condensed to form the oceans. Large amounts of CO₂ dissolved in the oceans, reacting to form insoluble carbonate compounds.

9.3 How Oxygen Increased

Algae and plants produced oxygen by photosynthesis over billions of years:

As oxygen levels rose, animals evolved.

9.4 How Carbon Dioxide Decreased

CO₂ levels decreased because:

• Photosynthesis absorbed CO₂.
• CO₂ dissolved in oceans and formed carbonate precipitates.
• Carbon was locked up in sedimentary rocks (e.g. limestone) and fossil fuels (coal, oil, gas) formed from dead plant and animal remains.

9.5 Greenhouse Gases

Greenhouse gases (carbon dioxide, water vapour, methane) maintain temperatures on Earth high enough to support life.

The Greenhouse Effect: Short-wavelength radiation from the Sun passes through the atmosphere and heats the Earth. The Earth emits long-wavelength (infrared) radiation, which is absorbed by greenhouse gases in the atmosphere, trapping heat.

Human activities (burning fossil fuels, deforestation, agriculture) have increased CO₂ and methane levels, leading to global climate change.

9.6 Pollutants from Fuels

Combustion of fuels releases pollutants that harm health and the environment:

• Carbon Dioxide (CO₂): Greenhouse gas, causes climate change.
• Carbon Monoxide (CO): Toxic, colourless, odourless gas. Binds to haemoglobin, reducing oxygen carriage in the blood.
• Sulfur Dioxide (SO₂) & Nitrogen Oxides (NO_x): Cause respiratory problems and acid rain, which damages plants and buildings.
• Particulates (soot): Solid carbon particles. Cause global dimming and lung damage.

10.1 Sustainable Development

Humans use Earth's resources for warmth, shelter, food, and transport. Resources are finite (will run out, e.g. metal ores, fossil fuels) or renewable (reformed at a rate similar to use, e.g. timber).

10.2 Potable Water

Potable water is water that is safe to drink. It is not pure water in a chemical sense (which contains H₂O molecules only) as it contains low levels of dissolved salts and microbes.

Treatment of fresh water (groundwater/rivers):

1. Filtration: Passed through filter beds to remove insoluble solids.
2. Sterilisation: Treated with chlorine, ozone, or UV light to kill microbes.

If fresh water is scarce, desalination of salty water (seawater) is used by distillation or reverse osmosis. Both require large amounts of energy.

10.3 Waste Water Treatment

Sewage, agricultural, and industrial waste water must be treated before release to prevent pollution:

1. Screening: Removes large solid objects.
2. Sedimentation: Settles into solid sludge (sinks) and liquid effluent (floats).
3. Aerobic digestion: Effluent is treated with aerobic bacteria to break down organic matter.
4. Anaerobic digestion: Sludge is digested by anaerobic bacteria (produces biogas).

10.4 Alternative Metal Extraction Chemistry Only

Copper ores are finite. Phytomining and bioleaching extract copper from low-grade ores, avoiding traditional mining waste.

• Phytomining: Plants are grown on soil containing copper compounds. The plants absorb copper and accumulate it in their tissues. They are harvested and burned. The ash containing copper compounds is dissolved in acid to make a leachate, and copper is displaced by scrap iron or extracted by electrolysis.
• Bioleaching: Bacteria are used to produce acidic solutions (leachate) containing copper ions from the ore. Copper is then extracted using displacement or electrolysis.

10.5 Life Cycle Assessment (LCA)

An LCA assesses the environmental impact of a product over its entire life cycle:

1. Extracting and processing raw materials.
2. Manufacturing and packaging.
3. Use and operation during its lifetime.
4. Disposal at the end of its life (incineration, landfill, recycling).

LCAs are not purely objective; while energy, water, and waste can be quantified, impacts like visual pollution require subjective value judgements.

10.6 Reduce, Reuse, Recycle

Reduces the use of limited resources, energy consumption, waste, and environmental impacts. Glass, metals, and plastics are commonly recycled by melting and reshaping.

10.7 Corrosion Chemistry Only

Corrosion is the destruction of materials by chemical reactions with substances in the environment (e.g. rusting of iron). Rusting requires both oxygen and water.

Rusting equation: iron + oxygen + water → hydrated iron(III) oxide

Prevention: Barriers (paint, grease, electroplating); sacrificial protection (connecting to a more reactive metal like zinc, which corrodes instead of the iron); galvanising (coating iron in zinc, acting as both barrier and sacrificial protection).

10.8 Alloys & Materials Chemistry Only

• Bronze: Copper + tin. Hard and corrosion-resistant.
• Brass: Copper + zinc. Malleable, used for musical instruments.
• Gold alloys: Gold + silver/copper/zinc. Harder than pure gold (measured in carats, where 24 carat is 100% pure).
• Steels: Iron + carbon/other metals. High-carbon steel is strong but brittle. Low-carbon steel is softer and easily shaped. Stainless steel (with chromium/nickel) is corrosion-resistant.
• Polymers: Thermosoftening polymers melt when heated (contain individual chains with weak intermolecular forces). Thermosetting polymers do not melt when heated (contain cross-links between chains).

10.9 The Haber Process Chemistry Only

Manufactures ammonia (NH₃) from nitrogen (from air) and hydrogen (from natural gas). Ammonia is used mainly to make fertilisers (e.g. ammonium nitrate, NH₄NO₃).

Conditions (a compromise between rate and yield)

• Temperature: ~450 °C. The forward reaction is exothermic, so a lower temperature would give a higher yield, but the rate would be too slow.
• Pressure: ~200 atm. High pressure shifts equilibrium to the right (fewer gas moles), increasing yield. Very high pressures are expensive and hazardous.
• Catalyst: iron. Speeds up both forward and reverse reactions equally. Does not change the yield, but equilibrium is reached faster.

The yield per pass is only about 15%. Unreacted nitrogen and hydrogen are recycled back through the reactor to improve overall conversion. The ammonia is cooled, liquefied, and removed.

Topic 1 Answers

Topic 1 Answers

1. Compound vs Mixture: Elements in a compound are chemically bonded in fixed proportions and can only be separated chemically. A mixture consists of substances that are not chemically bonded and can be separated by physical methods.
2. Mendeleev Gaps: He left gaps for undiscovered elements so that elements with similar chemical properties remained in the same groups/columns.
3. Group 1 Reactivity: Reactivity increases down the group because the atomic radius increases and there is more shielding from inner shells. Therefore, the single outer electron is less attracted to the nucleus and is lost more easily.
4. Sodium + Water reaction: Word: sodium + water → sodium hydroxide + hydrogen. Symbol: 2Na + 2H₂O → 2NaOH + H₂.
5. Group 0 Boiling Points: Boiling points increase down the group because the atoms become larger and have more electrons, which increases the strength of the weak intermolecular forces. More energy is required to overcome these forces.
6. Relative Atomic Mass (A_r) calculation: A_r = ((63 × 69) + (65 × 31)) / 100 = (4347 + 2015) / 100 = 6362 / 100 = 63.6.

Topic 1 Revision Checklist

Check off each area once you are confident you can do it:

• Write word and balanced symbol equations (including state symbols).
• Describe physical separation methods and calculate R_f values.
• Outline the changing model of the atom from Dalton to Chadwick.
• State relative charge/mass of protons, neutrons, and electrons.
• Calculate neutrons from mass/atomic number and relative atomic mass (A_r).
• Work out electronic structures (e.g. 2.8.1) and link to Periodic groups/periods.
• Explain reactivity and physical property trends in Groups 1, 7, and 0.
• Compare Transition Metals with Group 1 alkali metals.

Topic 2 Answers

Topic 2 Answers

1. Ionic conductivity: In a solid, ions are fixed in a giant lattice and cannot move. When molten or dissolved, the lattice breaks and ions are free to move and carry electrical charge.
2. Small molecules boiling point: While the atoms inside the molecules are held by strong covalent bonds, the forces between the molecules (intermolecular forces) are very weak and require little thermal energy to break.
3. Diamond vs Graphite: Diamond has a rigid 3D lattice where each carbon forms 4 bonds. Graphite has layers of carbon bonded to 3 others. The 4th electron is delocalised and free to move along layers, allowing it to conduct electricity.
4. Malleability & Alloys: Pure metals have layers of identical ions that slide easily. Alloys contain different-sized atoms that distort layers, making it harder for them to slide.
5. Polymers solid state: Polymers are long chains of repeating monomer units. Their large molecular size means their intermolecular forces are relatively strong, keeping them solid at room temperature.

Topic 2 Revision Checklist

Check off each area once you are confident you can do it:

• Describe ionic, covalent, and metallic bonding in terms of electrons.
• Draw dot-and-cross diagrams for simple ionic and covalent compounds.
• Explain properties of ionic lattices (high melting points, solubility, conductivity).
• Differentiate between small molecules, giant covalent structures, and polymers.
• Explain why pure metals are malleable while alloys are hard.
• Describe the three states of matter and the particle model limitations.
• Explain the unique properties of nanoparticles based on surface area to volume ratio.

Topic 3 Answers

Topic 3 Answers

1. Conservation of Mass: Mass is conserved in reactions because no atoms are lost or made. The mass of the products equals the mass of the reactants.
2. sulfuric acid M_r: M_r = (1 × 2) + 32 + (16 × 4) = 2 + 32 + 64 = 98.
3. Avogadro's Constant: The number of atoms, molecules, or ions in one mole of any substance, equal to 6.02 × 10²³.
4. Yield vs Atom Economy: Percentage yield measures the efficiency of the reaction process in obtaining product (actual vs theoretical mass). Atom economy measures the proportion of reactant atoms that end up in the desired product rather than waste.
5. Gas volume calculation: Volume = moles × 24 = 0.5 × 24 = 12 dm³.

Topic 3 Revision Checklist

Check off each area once you are confident you can do it:

• Explain apparent mass changes in reactions involving gases.
• Calculate relative formula masses (M_r) of compounds.
• Convert mass to moles and vice versa using formula mass.
• Use balanced symbol equations to calculate masses of reactants or products.
• Identify the limiting reactant in a reaction and calculate excess amounts.
• Calculate percentage yield and outline why it is rarely 100%.
• Calculate atom economy and explain its environmental and cost benefits.

Topic 4 Answers

Topic 4 Answers

1. Oxidation vs Reduction: Oxidation is the gain of oxygen by a substance; reduction is the loss of oxygen from a substance.
2. Metal extraction: Copper oxide can be reduced by heating with carbon because copper is less reactive than carbon. Magnesium is more reactive than carbon, so it must be extracted via electrolysis.
3. Neutralisation equation: H⁺(aq) + OH^-(aq) → H₂O(l).
4. Strong vs Weak acids: Strong acids completely ionise/dissociate in water to release all H⁺ ions. Weak acids only partially ionise in water.
5. Na₂SO₄ electrolysis: Cathode (-): Hydrogen gas (H₂) is formed because sodium is more reactive than hydrogen. Anode (+): Oxygen gas (O₂) is formed because no halide ions are present.

Topic 4 Revision Checklist

Check off each area once you are confident you can do it:

• Write word and balanced symbol equations for reactions of acids.
• Outline the Required Practical to prepare a pure, dry sample of a soluble salt.
• Explain how the pH of a solution relates to hydrogen ion concentration.
• Explain acid strength vs acid concentration.
• Describe the processes of oxidation and reduction in terms of electron transfer (OIL RIG).
• Predict the electrolysis products of molten and aqueous compounds.
• Write half equations for anode and cathode reactions.

Topic 5 Answers

Topic 5 Answers

1. Exothermic vs Endothermic: Exothermic reactions release energy to the surroundings, causing the surrounding temperature to increase. Endothermic reactions take in energy from the surroundings, causing the temperature to decrease.
2. Activation energy on profile: It is shown as an upward arrow from the energy level of the reactants to the peak of the curve.
3. Exothermic bond energies: Energy is taken in to break bonds, and energy is released when new bonds form. In an exothermic reaction, the energy released when making bonds is greater than the energy required to break bonds.
4. Cell voltage factors: The identity/reactivity difference of the two metal electrodes, and the concentration/type of electrolyte.
5. Hydrogen fuel cell product: Water (H₂O).

Topic 5 Revision Checklist

Check off each area once you are confident you can do it:

• Classify chemical reactions as exothermic or endothermic from temperature data.
• Draw and label reaction profiles for exothermic and endothermic processes.
• Outline the Required Practical to investigate temperature changes.
• Calculate energy changes of reactions using bond energy values.
• Describe how chemical cells generate electricity.
• Evaluate the use of hydrogen fuel cells compared to rechargeable batteries.

Topic 6 Answers

Topic 6 Answers

1. Collision Criteria: Reacting particles must collide with each other, and they must have energy greater than or equal to the activation energy.
2. Surface area effect: Breaking a solid increases exposed surface area, which increases the frequency of collisions between reactant particles.
3. How catalysts work: A catalyst provides an alternative reaction pathway with a lower activation energy, so a higher proportion of collisions are successful.
4. Reversible reaction energy: It must be exothermic, and the quantity of energy released is equal to the energy absorbed in the forward reaction.
5. Pressure change on equilibrium: Increasing pressure shifts the equilibrium to the right (products side) because there are fewer gas molecules on the right (2 moles) than on the left (4 moles). This counteracts the increase in pressure.

Topic 6 Revision Checklist

Check off each area once you are confident you can do it:

• Calculate mean rate of reaction from experimental measurements.
• Determine rate at a specific time from a tangent line on a graph.
• Outline Required Practicals to measure reaction rates (gas volume or colour change).
• Explain factors affecting reaction rates using collision theory.
• Draw and interpret reaction profiles with and without catalysts.
• Define dynamic equilibrium and state conditions required to reach it.
• Predict shifts in equilibrium using Le Chatelier's Principle.

Topic 7 Answers

Topic 7 Answers

1. Alkane formula: C_nH_2n+2.
2. Hydrocarbon trends: As chain length increases, boiling point increases, viscosity increases (becomes thicker), and flammability decreases.
3. Methane combustion: CH₄ + 2O₂ → CO₂ + 2H₂O.
4. Test for alkenes: Shake with bromine water. Alkenes decolourise orange bromine water to colourless; alkanes do not.
5. Organic functional groups: (a) Alcohol: -OH. (b) Carboxylic acid: -COOH. (c) Ester: -COO-.

Topic 7 Revision Checklist

Check off each area once you are confident you can do it:

• Explain how fractional distillation separates crude oil.
• Write balanced chemical equations for complete and incomplete combustion.
• Describe the process of cracking and explain why it is carried out.
• Draw displayed formulae for the first four alkanes and alkenes.
• Describe the addition reactions of alkenes with halogens, water, and hydrogen.
• Compare addition and condensation polymerisation.
• Describe structure and importance of DNA, proteins, starch, and cellulose.

Topic 8 Answers

Topic 8 Answers

1. Pure substance definition: A single element or compound that is not mixed with any other substance.
2. Melting point check: Pure substances melt at a sharp, specific temperature. Impure substances melt over a range of temperatures.
3. R_f formula: R_f = distance moved by substance / distance moved by solvent.
4. Chlorine gas test: Place damp litmus paper in the gas. The litmus paper is bleached and turns white.
5. Bromide precipitate: A cream-coloured precipitate of silver bromide (AgBr) is formed.

Topic 8 Revision Checklist

Check off each area once you are confident you can do it:

• Explain how paper chromatography separates mixtures.
• Calculate R_f values from chromatograms.
• Describe chemical tests for hydrogen, oxygen, carbon dioxide, and chlorine gases.
• Carry out flame tests to identify metal cations.
• Use sodium hydroxide solution to distinguish between metal cations.
• Identify carbonate, halide, and sulfate anions using chemical tests.
• State the advantages of instrumental analysis methods compared to chemical tests.

Topic 9 Answers

Topic 9 Answers

1. Atmosphere percentages: Nitrogen: ~78%, Oxygen: ~21%.
2. Early volcanic gases: Mostly carbon dioxide, water vapour, nitrogen, and small amounts of methane and ammonia.
3. Photosynthesis equation: carbon dioxide + water → glucose + oxygen.
4. Greenhouse effect: Short-wavelength radiation from the Sun passes through the atmosphere. The Earth absorbs this and emits long-wavelength infrared radiation, which is absorbed by greenhouse gases, trapping heat.
5. Carbon monoxide: It is colourless and odourless. It is dangerous because it binds to haemoglobin, preventing oxygen transport in the blood.

Topic 9 Revision Checklist

Check off each area once you are confident you can do it:

• State the composition of the modern atmosphere.
• Describe theories for the evolution of the Earth's atmosphere.
• Explain how photosynthesis increased oxygen and decreased carbon dioxide.
• Outline how carbon became locked up in rocks and fossil fuels.
• Describe the greenhouse effect in terms of radiation wavelengths.
• List human activities that increase greenhouse gas concentrations.
• State the sources and effects of CO, SO₂, NO_x, and particulates.

Topic 10 Answers

Topic 10 Answers

1. Pure vs Potable water: Pure water is chemically 100% H₂O. Potable water is water that is safe to drink, containing low levels of dissolved salts and no harmful microbes.
2. Fresh water treatment: Filtration through sand/gravel beds to remove insoluble solids, followed by sterilisation using chlorine, ozone, or UV light to kill bacteria and microbes.
3. LCA stages: (1) Raw material extraction/processing, (2) manufacturing and packaging, (3) use and operation, and (4) disposal.
4. Rusting conditions: Both oxygen (from air) and water are required.
5. Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). Conditions: 450 °C, 200 atm pressure, iron catalyst.

Topic 10 Revision Checklist

Check off each area once you are confident you can do it:

• Distinguish between finite and renewable resources.
• Explain how potable water is produced from fresh water and sea water.
• Describe how sewage water is treated in stages.
• Describe phytomining and bioleaching of copper.
• Compare the environmental impact of products using LCAs.
• Explain corrosion prevention methods and write rusting equations.
• Explain the compromise conditions of temperature and pressure in the Haber process.

GCSE Chemistry Glossary

AQA GCSE Chemistry Required Practicals

This hub provides quick-links to all key Required Practicals (RPs) featured in the AQA specification. Click any practical to jump directly to its revision notes.