Learn Chemistry with understanding

1. States of matter

Learning outcomes: distinguish solids, liquids and gases; describe particle separation, arrangement and motion; explain changes of state; relate gas volume to temperature and pressure; and explain diffusion using kinetic particle theory.

1.1 The particle model

All matter is made of tiny particles (atoms, molecules or ions). The particles are in constant motion. Their spacing, arrangement, speed and attractive forces help explain the properties of each state.

1.2 Changes of state

Melting is the change from solid to liquid; freezing is liquid to solid; boiling and evaporation are liquid to gas; condensation is gas to liquid; sublimation is solid directly to gas. During a change of state, energy is transferred but the temperature of a pure substance remains constant while the change is taking place at a fixed pressure.

1.3 Kinetic particle theory and heating curves

When a substance is heated, its particles gain energy. In a single state, the average kinetic energy and temperature increase. At the melting or boiling point, energy is used to overcome attractive forces and separate particles rather than increase their average kinetic energy; therefore the temperature stays constant until the change is complete. During cooling, the reverse occurs and energy is released to the surroundings.

1.4 Temperature and pressure effects on gas volume

Temperature: at constant pressure, heating a fixed amount of gas increases its volume. Gas particles move faster and collide more frequently/forcefully; the gas expands until its pressure balances the external pressure. Cooling reduces the volume.

Pressure: at constant temperature, increasing external pressure compresses a gas and decreases its volume. Decreasing pressure allows the gas to expand. Gases are compressible because there is substantial empty space between their particles.

1.5 Diffusion

Diffusion is the net movement of particles from a region of higher concentration to a region of lower concentration due to their random motion, until the particles become more evenly distributed. It occurs in gases and liquids and is generally faster in gases because gas particles move rapidly and are far apart.

Effect of relative molecular mass (Supplement): at the same temperature, lighter gas molecules generally diffuse faster than heavier molecules because their average speeds are higher. For example, ammonia (Mᵣ = 17) diffuses faster than hydrogen chloride (Mᵣ ≈ 36.5). In a tube, the white ammonium chloride ring forms closer to the HCl end because NH₃ travels farther in the same time.

Worked example

Question: A sealed flexible balloon is warmed while the external pressure remains approximately constant. Explain why its volume increases.

Solution: Heating transfers energy to the gas particles. Their average kinetic energy and speed increase. They collide more frequently and forcefully with the balloon walls. The balloon expands, increasing its volume, until the internal pressure again balances the external pressure.

1.6 Evaporation: factors that affect its rate

Evaporation happens when particles at the surface of a liquid have enough energy to escape into the gas state. It can occur at temperatures below the boiling point. The rate increases when:

• Temperature increases: a greater proportion of particles have enough kinetic energy to escape.
• Surface area increases: more particles are exposed at the surface at the same time.
• Air movement increases: moving air carries vapour away from the surface, reducing the amount of vapour close to the liquid.
• Humidity decreases: drier air can accept more water vapour before becoming saturated.

1.7 Gas pressure explained by collisions

Gas particles move randomly and collide with the walls of their container. Each collision changes the particles' momentum and exerts a force on the wall. The combined force per unit area is the gas pressure. At a fixed temperature, compressing a gas into a smaller volume makes wall collisions more frequent, so pressure increases.

1.8 Diffusion investigation: ammonia and hydrogen chloride

When cotton wool soaked in concentrated aqueous ammonia and cotton wool soaked in concentrated hydrochloric acid are placed at opposite ends of a long glass tube, the gases diffuse towards each other. A white ring of solid ammonium chloride forms where they meet.

NH₃(g) + HCl(g) → NH₄Cl(s)

1.9 Exam-style application questions

Common misconceptions to avoid

• Particles themselves do not expand significantly when heated; the average separation changes.
• Particles in a solid are not stationary; they vibrate about fixed positions.
• During melting or boiling of a pure substance at fixed pressure, temperature does not rise until the change is complete.
• Diffusion is not the same as bulk movement caused by stirring or wind; it results from random particle motion.

Quick knowledge check

Self-assessment: explain each answer aloud, draw the particle diagrams from memory, and practise applying the model to unfamiliar situations.

2. Atoms, elements and compounds

Learning outcomes: distinguish elements, compounds and mixtures; describe the structure of atoms; use proton number and nucleon number; interpret isotopes; write electronic configurations for the first 20 elements; explain ion formation and ionic bonding; describe covalent and metallic bonding; and relate bonding and structure to physical properties.

2.1 Elements, compounds and mixtures

An element is a substance consisting of only one type of atom. A compound contains two or more different elements chemically bonded together in fixed proportions. A mixture contains two or more substances that are not chemically bonded and can be present in any proportion.

2.2 Inside the atom

An atom has a tiny, dense, positively charged nucleus containing protons and neutrons. Negatively charged electrons occupy shells (energy levels) around the nucleus. Most of the atom is empty space. In a neutral atom, the number of protons equals the number of electrons.

2.3 Proton number and nucleon number

Proton number (atomic number, Z) is the number of protons in the nucleus. It identifies the element. Nucleon number (mass number, A) is the total number of protons and neutrons in the nucleus.

Number of neutrons = A − Z

For a neutral atom: number of electrons = number of protons. For an ion, the proton number remains unchanged but the number of electrons differs.

2.4 Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have the same proton number and different nucleon numbers. Neutral isotopes have the same number of electrons and therefore very similar chemical properties, because chemical reactions depend mainly on electron arrangement. Their physical properties, such as mass, can differ.

2.5 Electronic configuration and the Periodic Table

Electrons occupy shells around the nucleus. For the first 20 elements at this level, the shell capacities used in simple configurations are 2 in the first shell, 8 in the second, 8 in the third, then 2 in the fourth. Electrons fill the lower-energy shells first. The number of occupied shells indicates the period; for main-group elements, the number of outer-shell electrons helps identify the group.

2.6 Ions and ionic bonding

An ion is a charged particle formed when an atom or group of atoms loses or gains electrons. Metals usually lose electrons to form positive ions (cations); non-metals usually gain electrons to form negative ions (anions). The nucleus does not change during ordinary ion formation.

Ionic bonding is the strong electrostatic attraction between oppositely charged ions. It commonly forms when electrons transfer from a metal atom to a non-metal atom. The resulting ionic compound is electrically neutral overall.

Common ion charges: Group I metals form 1+ ions; Group II metals form 2+ ions; Group III metals commonly form 3+ ions; Group VI non-metals commonly form 2− ions; Group VII non-metals form 1− ions. Learn the ions specified by your syllabus and practise balancing total charge.

Worked formula example: magnesium chloride. Magnesium forms Mg²⁺ and chlorine forms Cl⁻. Two chloride ions are needed to balance one magnesium ion: (+2) + 2(−1) = 0. The formula is MgCl₂.

2.7 Properties of ionic compounds

• High melting and boiling points: strong electrostatic attractions between oppositely charged ions in a giant ionic lattice require much energy to overcome.
• Conduct electricity when molten or in aqueous solution: ions are free to move and carry charge.
• Do not conduct when solid: ions are held in fixed positions and cannot move through the lattice.
• Often brittle: shifting layers can bring like charges next to each other; repulsion causes the crystal to split.

2.8 Covalent bonding

A covalent bond is a shared pair of electrons between two atoms. It forms mainly between non-metal atoms. Each atom contributes one electron to each shared pair in a simple covalent bond. Atoms share electrons to achieve a more stable outer-shell arrangement.

Simple molecular substances: molecules contain strong covalent bonds within each molecule, but the forces between molecules are usually much weaker. Therefore many simple molecular substances have low melting and boiling points and do not conduct electricity because they have no mobile ions or delocalised electrons.

Worked example: methane, CH₄. Carbon has four outer-shell electrons. It shares one pair with each of four hydrogen atoms, forming four single covalent bonds. Each hydrogen achieves two electrons in its first shell; carbon achieves eight in its outer shell.

2.9 Giant covalent structures

In giant covalent structures, many atoms are joined by a continuous network of covalent bonds. A large amount of energy is needed to break these bonds, so they generally have very high melting points.

2.10 Metallic bonding

Metals consist of a giant lattice of positive metal ions surrounded by a sea of delocalised electrons. Metallic bonding is the strong electrostatic attraction between the positive ions and the delocalised electrons.

Metals usually conduct electricity because delocalised electrons can move through the structure and carry charge. They conduct heat because mobile electrons transfer energy. They are malleable and ductile because layers of ions can slide while attraction to the electron sea maintains the metallic bond.

2.11 Compare bonding and structure

2.12 Exam technique: how to earn explanation marks

• For ionic bonding, use the words electron transfer, oppositely charged ions and electrostatic attraction.
• For covalent bonding, state that atoms share pairs of electrons.
• For metallic bonding, mention positive ions, delocalised electrons and electrostatic attraction.
• When explaining melting points, identify the structure and the forces/bonds that must be overcome.
• When explaining conductivity, identify the mobile charged particles (ions or electrons).

Worked examples

Common misconceptions

• Isotopes are not different elements: they have the same proton number.
• When an atom forms an ion, electrons change; the number of protons does not.
• Ionic compounds are not made of separate molecules; they form giant ionic lattices.
• Strong covalent bonds within molecules are not the same as the weaker forces between simple molecules.
• Graphite conducts because of delocalised electrons, not because its layers slide.
• Diamond does not conduct electricity because it has no mobile charged particles.

Topic 2 knowledge and exam check

Revision routine: draw an atom, an isotope pair, an ionic transfer diagram, a covalent dot-and-cross diagram, a graphite layer and a metallic lattice from memory. Then answer the questions without viewing the marking points and award only the marks supported by precise chemistry statements.

3. Stoichiometry

Learning outcomes: write and interpret chemical formulae; construct and balance symbol equations; calculate relative formula masses; use mole ratios to calculate reacting masses and gas volumes; determine empirical formulae; and calculate solution concentrations.

3.1 Chemical formulae and balancing equations

A chemical formula shows the elements present and the simplest whole-number ratio of their atoms (or ions). In a balanced chemical equation, the number of atoms of each element is the same on both sides. This follows the law of conservation of mass: atoms are rearranged, not created or destroyed.

Balancing method: write correct formulae first; count each type of atom; add whole-number coefficients in front of formulae; recount; reduce coefficients to the smallest whole-number ratio. Never change subscripts to balance an equation because that changes the substance.

3.2 Relative atomic mass and relative formula mass

Relative atomic mass (Aᵣ) is the weighted mean mass of an atom of an element compared with one-twelfth of the mass of a carbon-12 atom. Relative molecular mass (Mᵣ) or relative formula mass is the sum of the relative atomic masses of all atoms shown in a formula. It has no unit.

3.3 The mole and Avogadro constant

The mole (mol) is the amount of substance containing the Avogadro number of specified particles. The Avogadro constant is approximately 6.02 × 10²³ mol⁻¹. The particles may be atoms, molecules, ions or formula units; identify which particles are being counted.

For a pure substance, amount in moles (n) = mass in grams (m) ÷ relative mass in g mol⁻¹. Rearranging gives m = n × Mᵣ. For elements, use Aᵣ in place of Mᵣ.

3.4 Using a balanced equation: mole ratios

The coefficients in a balanced equation give the reacting mole ratio. Convert the known quantity into moles, use the coefficient ratio to find the required moles, then convert to the requested unit.

3.5 Empirical and molecular formulae

The empirical formula is the simplest whole-number ratio of atoms in a compound. To find it from percentage composition or masses: (1) divide each mass or percentage by its Aᵣ; (2) divide all results by the smallest; (3) multiply to obtain whole numbers if necessary.

If the molecular mass is known, molecular formula = empirical formula × factor, where factor = molecular mass ÷ empirical formula mass. The factor must be a whole number.

3.6 Gas volumes and the mole

At room temperature and pressure (r.t.p.), one mole of any gas occupies approximately 24 dm³ (24 000 cm³). Use the molar gas volume stated in the question or required syllabus convention. Always convert units consistently.

For gases reacting together, use the balanced equation's mole ratio. At the same temperature and pressure, gas volume ratios equal mole ratios.

3.7 Solution concentration

Concentration describes the amount of solute dissolved in a given volume of solution. Molar concentration is measured in mol dm⁻³.

Rearrangements: n = c × V and V = n ÷ c. Convert cm³ to dm³ by dividing by 1000. Convert dm³ to cm³ by multiplying by 1000.

3.8 A reliable calculation pathway

3.9 Common errors to avoid

• Changing formula subscripts instead of adding coefficients when balancing equations.
• Forgetting to multiply every atom inside brackets when calculating Mᵣ.
• Using mass directly in a mole ratio; convert to moles first.
• Using an unbalanced equation to obtain mole ratios.
• Forgetting that 1000 cm³ = 1 dm³.
• Writing units incorrectly: mass in g, volume in dm³ or cm³, concentration in mol dm⁻³.
• Rounding too early; keep extra digits during working and round the final answer appropriately.

3.10 Topic 3 exam practice

Self-assessment: attempt each question without opening the marking points. Award only marks earned by correct formulae, substitutions, units and explanations. Rework any calculation where the mole ratio or volume conversion was missed.

4. Electrochemistry

Learning outcomes: explain electrolysis in terms of mobile ions; identify anode, cathode and electrolyte; predict products for selected molten and aqueous electrolytes; write ionic half-equations; describe electroplating; and explain the operation of a hydrogen–oxygen fuel cell.

4.1 What is electrolysis?

Electrolysis is the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of direct electric current. The liquid or solution that conducts because it contains mobile ions is the electrolyte. Electrolysis requires ions that are free to move: ions in a solid ionic lattice are fixed in position, so solid sodium chloride does not conduct by ion movement.

• Cathode (negative electrode): positive ions (cations) move towards it and gain electrons. Reduction occurs at the cathode.
• Anode (positive electrode): negative ions (anions) move towards it and lose electrons. Oxidation occurs at the anode.
• Electrons move through the external wires and power supply; ions carry charge through the electrolyte.

4.2 Electrolysis of molten lead(II) bromide

Molten lead(II) bromide, PbBr₂, contains mobile Pb²⁺ and Br⁻ ions. At the cathode, lead(II) ions gain electrons to form lead. At the anode, bromide ions lose electrons to form bromine gas.

Overall equation: PbBr₂(l) → Pb(l) + Br₂(g). The electrolyte must be molten so its ions can move to the electrodes.

4.3 Electrolysis of aqueous solutions

An aqueous electrolyte contains ions from the dissolved compound and from water. At the cathode, a metal less reactive than hydrogen may be deposited; otherwise hydrogen is often produced. At the anode, halide ions may form halogens when present in sufficiently concentrated solution; otherwise oxygen is often produced. Always apply the syllabus rules and conditions stated in a question.

4.4 Copper(II) sulfate solution with inert electrodes

Using inert graphite or platinum electrodes, aqueous copper(II) sulfate contains Cu²⁺, SO₄²⁻, H⁺ and OH⁻ ions. Copper is less reactive than hydrogen, so Cu²⁺ ions are discharged at the cathode. Oxygen is produced at the anode from hydroxide ions / water.

The sulfate ions remain in solution. With inert electrodes, the solution becomes more acidic as water is decomposed and copper(II) ions are removed from solution.

4.5 Concentrated aqueous sodium chloride (brine)

Brine contains Na⁺, Cl⁻, H⁺ and OH⁻ ions. At the cathode, hydrogen is formed rather than sodium because sodium is much more reactive than hydrogen. At the anode, concentrated chloride ions form chlorine gas.

Sodium hydroxide remains in solution. The overall products are hydrogen, chlorine and sodium hydroxide. Chlorine is toxic: tests and demonstrations must be carried out only with suitable teacher supervision, ventilation and safety controls.

4.6 Writing and checking half-equations

A half-equation shows the electron transfer at one electrode. Check that both atoms and total charge are balanced. Electrons appear on the right when they are lost (oxidation) and on the left when they are gained (reduction).

4.7 Electroplating

Electroplating coats an object with a thin layer of another metal using electrolysis. The object to be plated is the cathode (negative electrode); the plating metal is commonly the anode (positive electrode); and the electrolyte contains ions of the plating metal.

1. Clean the object thoroughly so the coating adheres evenly.
2. Connect the object to the negative terminal and the plating metal to the positive terminal.
3. Immerse both in a suitable solution containing ions of the plating metal.
4. Pass direct current. Metal ions gain electrons at the object and form a metal coating.

Uses include corrosion protection, improving appearance and giving a cheaper object a surface with the properties of a more expensive metal. The object must be clean; poor cleaning can produce an uneven or poorly attached coating.

4.8 Hydrogen–oxygen fuel cells

A hydrogen–oxygen fuel cell converts chemical energy directly into electrical energy. Hydrogen is supplied to one electrode and oxygen to the other. The gases react to form water, while electrons flow through the external circuit and provide an electric current. A fuel cell continues operating while reactants are supplied; it is not simply a rechargeable battery.

Overall reaction: 2H₂ + O₂ → 2H₂O. The only chemical product of the ideal hydrogen–oxygen cell is water. In practice, the environmental impact depends on how hydrogen is produced, compressed, transported and stored.

4.9 Comparison: electrolysis and fuel cells

4.10 Topic 4 exam practice

Practical safety: wear eye protection; use low-voltage d.c. apparatus under teacher supervision; keep gases away from flames unless a controlled gas test is being demonstrated; chlorine must never be inhaled; and follow the laboratory risk assessment for toxic, corrosive or flammable substances.

5. Chemical energetics

Learning outcomes: distinguish exothermic and endothermic reactions; describe energy transfer between a reacting system and its surroundings; interpret reaction profile diagrams; identify activation energy and the enthalpy change; explain the energy changes involved in bond breaking and bond formation; and analyse simple temperature-change experiments.

5.1 Energy changes during chemical reactions

Chemical reactions involve changes in the arrangement of atoms and in the chemical bonds between them. Energy is transferred during these changes. In a typical school experiment, the reaction mixture and its surroundings may become warmer or cooler.

• Exothermic reaction: transfers energy to the surroundings. The surroundings usually become warmer and the measured temperature rises. The products have less chemical energy than the reactants.
• Endothermic reaction: takes in energy from the surroundings. The surroundings usually become cooler and the measured temperature falls. The products have more chemical energy than the reactants.

5.2 Reaction profile diagrams

A reaction profile shows how the energy of a reacting system changes as reactants are converted into products. The vertical axis is energy; the horizontal axis is the reaction pathway (not necessarily time).

Activation energy (Eₐ) is the minimum energy that reacting particles must have for a successful reaction. On a reaction profile, it is the energy difference between the reactants and the top of the energy barrier. The enthalpy change (ΔH) is the energy difference between products and reactants: ΔH = energy of products − energy of reactants.

5.3 Bond breaking and bond formation

Breaking a chemical bond requires energy: it is an endothermic process. Forming a chemical bond releases energy: it is an exothermic process. The overall energy change depends on the balance between energy absorbed to break bonds in the reactants and energy released when new bonds form in the products.

If more energy is released by bond formation than is absorbed by bond breaking, the reaction is exothermic. If more energy is absorbed to break bonds than is released by forming new bonds, the reaction is endothermic.

Worked example 1 – bond-energy calculation: In a simplified reaction, 500 kJ mol⁻¹ is absorbed breaking bonds and 650 kJ mol⁻¹ is released forming bonds. ΔH = 500 − 650 = −150 kJ mol⁻¹. The negative value means the reaction is exothermic.

Worked example 2 – interpret the sign: A reaction has ΔH = +85 kJ mol⁻¹. The positive sign means the products have 85 kJ mol⁻¹ more energy than the reactants; the reaction is endothermic and takes in energy.

5.4 Catalysts and activation energy

A catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy. It is not used up overall. A catalyst does not change the energy of the reactants or products and therefore does not change the overall enthalpy change, ΔH.

5.5 Investigating temperature changes

A simple method for comparing energy changes is to measure the temperature of a reaction mixture before and after a reaction. A temperature rise suggests an exothermic process; a temperature fall suggests an endothermic process.

1. Measure a known volume of solution using a measuring cylinder or pipette.
2. Record the initial temperature with a thermometer or temperature probe.
3. Mix the reactants, stir consistently and record the highest or lowest temperature reached.
4. Calculate the temperature change: ΔT = final temperature − initial temperature.
5. Repeat the experiment and compare results, keeping volumes, concentrations and starting temperatures consistent.

Worked example 3: A mixture warms from 21.5 °C to 29.0 °C. ΔT = 29.0 − 21.5 = +7.5 °C. The temperature rises, so the reaction is exothermic.

Worked example 4: A solution cools from 24.0 °C to 18.5 °C. ΔT = 18.5 − 24.0 = −5.5 °C. The temperature falls, so the process is endothermic.

5.6 Improving reliability and evaluating results

• Use an insulated polystyrene cup and a lid to reduce energy transfer to or from the surroundings.
• Use a thermometer or data logger with suitable resolution; stir gently so the temperature is uniform.
• Keep the mass or volume, concentration and initial temperature of reactants constant.
• Record temperature at short intervals to avoid missing the maximum or minimum.
• Repeat measurements and calculate a mean; identify anomalous results rather than automatically including them.
• State that heat exchange with surroundings and heat absorbed by the container make the measured temperature change less than the ideal value.

5.7 Quick recall

5.8 Topic 5 exam practice

Exam technique: label both axes on every reaction profile, show reactant and product levels clearly, draw activation energy from the reactant level to the peak, and show ΔH between reactant and product levels. Check whether a question asks for a temperature change (°C) or an enthalpy change (kJ mol⁻¹); they are not interchangeable.

6. Chemical reactions

Learning outcomes: distinguish physical and chemical changes; describe common reaction patterns; explain reaction rates using collision theory; investigate factors affecting rate; interpret rate graphs; explain reversible reactions and dynamic equilibrium; and identify oxidation and reduction.

6.1 Physical changes and chemical changes

A physical change changes the state, shape or arrangement of a substance without producing a new substance. Examples include melting ice, boiling water, dissolving sugar and cutting a metal. Many physical changes can be reversed by a physical process.

A chemical change produces one or more new substances with different chemical properties. Examples include burning magnesium, rusting iron, reacting an acid with a carbonate and decomposing hydrogen peroxide. Chemical changes often involve energy transfer and may show colour change, gas formation or precipitate formation, but no single observation proves a chemical reaction on its own.

6.2 Common reaction patterns

Recognising reaction patterns helps you predict products and write balanced equations. Learn the general pattern, then check the conditions and reactivity of the substances involved.

Worked example 1: Magnesium reacts with oxygen. The word equation is magnesium + oxygen → magnesium oxide. The balanced symbol equation is 2Mg + O₂ → 2MgO. Two magnesium atoms and two oxygen atoms are present on each side.

6.3 Rate of reaction

The rate of reaction measures how quickly reactants are used up or products are formed. A reaction is faster when a measurable change occurs in a shorter time. Rate can be followed by measuring gas volume, mass loss, concentration, colour intensity or the time taken for a visible change.

For a simple average-rate calculation:

When a graph shows product volume against time, the gradient represents rate. A steeper gradient means a faster reaction. The curve becomes horizontal when no more product is being formed, so the reaction has stopped or reached completion under the conditions used.

6.4 Collision theory

For a reaction to occur, reacting particles must collide. A collision only leads to reaction if particles collide with sufficient energy (at least the activation energy) and with a suitable orientation. Increasing the number of successful collisions per second increases the reaction rate.

6.5 Factors affecting reaction rate

Worked example 2: Calcium carbonate powder reacts faster with dilute hydrochloric acid than the same mass of large chips. The powder has a larger total exposed surface area, so acid particles can collide with the solid at more locations each second.

Worked example 3: Increasing acid concentration increases the reaction rate because there are more acid particles per unit volume, increasing the frequency of successful collisions with the solid.

6.6 Investigating reaction rates

Method A – gas collection: React a measured mass of marble chips (calcium carbonate) with a measured volume and concentration of dilute hydrochloric acid in a flask connected to a gas syringe. Record the carbon dioxide volume at regular time intervals. Keep the mass and size of chips, acid volume, acid concentration and temperature controlled when testing another variable.

Method B – loss of mass: Place a flask containing acid on a balance, add marble chips and record the mass at regular intervals. Carbon dioxide escapes, so the measured mass decreases. Use a cotton-wool plug to reduce splashing while allowing gas to escape; do not seal a gas-producing flask.

Method C – disappearing cross: Mix sodium thiosulfate solution with dilute hydrochloric acid over a marked cross. Sulfur forms and makes the mixture cloudy. Record the time until the cross can no longer be seen. Use the same viewing conditions and volumes. This method gives a comparative rate, often estimated as 1 ÷ time.

6.7 Reversible reactions

A reversible reaction can proceed in both directions. The forward reaction converts reactants into products; the reverse reaction converts products back into reactants. The symbol ⇌ is used to show that a reaction is reversible.

Example: hydrated copper(II) sulfate can be dehydrated by heating to form white anhydrous copper(II) sulfate. Adding water to the anhydrous solid reforms the blue hydrated compound:

Another familiar example is the hydration of cobalt(II) chloride, which changes colour as water is removed or added. Colour observations depend on the particular compound and conditions.

6.8 Dynamic equilibrium

In a closed system, a reversible reaction may reach dynamic equilibrium. At equilibrium, the forward and reverse reactions continue at equal rates. The concentrations of reactants and products remain constant, although the reactions have not stopped. The amounts of reactants and products are not necessarily equal.

6.9 Factors that affect equilibrium position

For reversible reactions where equilibrium shifts are considered, use the following qualitative rules (Le Chatelier's principle): a system responds to a change by shifting in the direction that opposes that change.

• Concentration: adding a reactant tends to shift equilibrium towards products; removing a product also tends to shift towards products.
• Pressure (gaseous systems): increasing pressure tends to favour the side with fewer moles of gas; decreasing pressure tends to favour the side with more moles of gas. If both sides have the same number of gas molecules, pressure has no effect on the position.
• Temperature: increasing temperature favours the endothermic direction; decreasing temperature favours the exothermic direction.
• Catalyst: speeds up both forward and reverse reactions equally and helps equilibrium be reached sooner; it does not change the equilibrium position.

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) is exothermic in the forward direction. Higher pressure favours ammonia because there are fewer gas molecules on the product side (2 rather than 4). Lower temperature favours the exothermic forward reaction, although it also makes the reaction slower.

6.10 Oxidation and reduction

Oxidation and reduction occur together in redox reactions. In terms of electrons, oxidation is loss of electrons and reduction is gain of electrons. The memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain.

In reactions involving oxygen, oxidation can also be described as gain of oxygen and reduction as loss of oxygen. In terms of hydrogen, oxidation can involve loss of hydrogen and reduction can involve gain of hydrogen. Use the definition most appropriate to the question.

Worked example 4: In CuO + H₂ → Cu + H₂O, copper(II) oxide is reduced because it loses oxygen; hydrogen is oxidised because it gains oxygen. Both processes happen in the same reaction.

6.11 Topic 6 exam practice

Exam technique: on rate graphs, state what each axis represents and use the gradient to compare rates. For equilibrium, say “equal rates” rather than “equal concentrations”. In redox questions, state clearly whether you are using electron transfer or oxygen transfer.

Scope note: Equilibrium-position rules are included as an explanatory extension for learners meeting reversible reactions; check the current examination syllabus and the exact depth required for your course and examination session.

7. Acids, bases and salts

Learning outcomes: describe acids and bases; distinguish strong and weak acids from concentrated and dilute solutions; explain pH and indicators; write equations for acid reactions and neutralisation; describe preparation of soluble and insoluble salts; and understand the purpose and basic method of acid–alkali titration.

7.1 Acids, bases and alkalis

An acid is a substance that produces hydrogen ions, H⁺(aq), when dissolved in water. An alkali is a soluble base that produces hydroxide ions, OH⁻(aq), in water. A base neutralises an acid; bases include metal oxides, metal hydroxides and some carbonates. Not all bases are soluble, so not all bases are alkalis.

Strong and weak acids: strength describes the extent to which an acid ionises in water. A strong acid ionises almost completely; a weak acid ionises only partially. Concentrated and dilute describe the amount of acid dissolved per unit volume. Strength and concentration are different ideas: a dilute strong acid is possible, as is a concentrated weak acid.

7.2 Characteristic reactions of acids

Some metals do not react with dilute non-oxidising acids under ordinary laboratory conditions. The reactivity series helps predict whether hydrogen will be displaced. Copper, for example, does not produce hydrogen with dilute hydrochloric acid.

Gas test: carbon dioxide turns limewater milky. Hydrogen gives a squeaky pop when a small sample is tested with a lit splint. Use only small quantities and teacher-approved safety procedures.

7.3 pH scale and indicators

The pH scale is used to describe how acidic or alkaline an aqueous solution is. At about 25 °C, pH 7 is neutral; values below 7 are acidic and values above 7 are alkaline. A lower pH indicates a more acidic solution; a higher pH indicates a more alkaline solution. The pH scale is logarithmic: a change of one pH unit corresponds to a tenfold change in hydrogen-ion concentration in dilute aqueous solutions.

Universal indicator is useful for estimating pH across a wide range. Litmus distinguishes acidic from alkaline solutions but does not give an accurate pH. In titration, a suitable indicator is selected to give a clear colour change near the end-point.

7.4 Neutralisation and ionic equations

Neutralisation occurs when hydrogen ions from an acid react with hydroxide ions from an alkali to form water. The net ionic equation is:

For hydrochloric acid and sodium hydroxide:

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

Chloride and sodium ions are spectator ions; they remain unchanged in solution. A complete ionic equation is H⁺ + Cl⁻ + Na⁺ + OH⁻ → Na⁺ + Cl⁻ + H₂O. Removing spectator ions gives the net ionic equation above.

Worked example 1: Sulfuric acid reacts with potassium hydroxide. The balanced equation is H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O. Two moles of KOH are required for each mole of H₂SO₄ because sulfuric acid can supply two hydrogen ions per formula unit.

7.5 Naming salts

The name of a salt depends on the metal (or ammonium) ion and the acid used:

• Hydrochloric acid forms chlorides.
• Nitric acid forms nitrates.
• Sulfuric acid forms sulfates.

Examples: magnesium + hydrochloric acid → magnesium chloride; copper(II) oxide + sulfuric acid → copper(II) sulfate; sodium hydroxide + nitric acid → sodium nitrate.

7.6 Preparing a soluble salt from an insoluble base

To prepare a soluble salt such as copper(II) sulfate from an insoluble base such as copper(II) oxide:

1. Warm a measured volume of dilute sulfuric acid gently in a beaker. Do not boil it.
2. Add copper(II) oxide in small portions while stirring until some solid remains unreacted. This ensures the acid has been used up.
3. Filter the mixture to remove excess unreacted copper(II) oxide.
4. Heat the filtrate gently to concentrate it, stopping before it becomes dry.
5. Allow the solution to cool so crystals form. If needed, leave it in a warm place to evaporate slowly.
6. Filter the crystals and dry them between filter papers or in a suitable warm place.

Why add excess base? It ensures all the acid reacts. The excess solid can then be removed by filtration, leaving a salt solution without excess acid.

7.7 Preparing a soluble salt by titration

If both reactants are soluble, such as hydrochloric acid and sodium hydroxide, excess reactant cannot be removed by filtration. Titration is used to find the exact volumes that neutralise one another. The accurate neutralisation is first carried out with an indicator. The same measured volumes are then mixed again without indicator to obtain a purer salt solution for crystallisation.

1. Rinse the burette with the solution it will contain, then fill it and remove any air bubble from the jet.
2. Use a pipette to transfer a fixed volume of alkali into a conical flask; add a few drops of a suitable indicator.
3. Record the initial burette reading at eye level.
4. Add acid while swirling the flask. Near the end-point, add it drop by drop until the indicator just changes colour permanently.
5. Record the final burette reading. Titre = final reading − initial reading.
6. Repeat until concordant titres are obtained (commonly within 0.10–0.20 cm³, according to the practical instructions), then calculate a suitable mean, excluding a rough titre and any justified anomalous result.

Worked example 2: Initial burette reading = 1.20 cm³; final reading = 24.65 cm³. Titre = 24.65 − 1.20 = 23.45 cm³.

End-point and equivalence point: the end-point is the observed indicator colour change. The equivalence point is when reactants have combined in the exact stoichiometric ratio. A suitable indicator is chosen so its colour-change range is close to the equivalence point.

7.8 Preparing an insoluble salt by precipitation

An insoluble salt can be prepared by mixing two soluble solutions that provide the required ions. The insoluble product forms as a precipitate and is separated by filtration.

Example – barium sulfate: mix aqueous barium chloride with aqueous sodium sulfate. A white precipitate of barium sulfate forms.

BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

Net ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

1. Prepare separate solutions of two suitable soluble salts.
2. Mix the solutions and stir; the insoluble salt forms as a precipitate.
3. Filter the mixture to collect the solid on filter paper.
4. Wash the precipitate with distilled water to remove soluble impurities.
5. Dry the solid between filter papers or in a suitable drying oven.

7.9 Salt solubility and selecting a preparation method

Use a solubility table when provided. Do not assume every salt is soluble or every carbonate is insoluble without checking the relevant rules for the course.

7.10 Practical safety and quality

• Wear safety goggles and a laboratory coat; handle dilute acids and alkalis carefully because they can irritate or damage skin and eyes.
• Use a pipette filler, never pipette by mouth. Read burette scales at eye level to avoid parallax error.
• Heat solutions gently using appropriate apparatus; point test tubes away from people and avoid heating a closed vessel.
• Do not taste chemicals or directly inhale gases. Wash spills promptly using the teacher's approved procedure.
• For purer crystals, remove insoluble excess reactant by filtration, concentrate carefully, allow crystals to form, then wash and dry as instructed.

7.11 Topic 7 exam practice

Exam technique: identify the acid first to determine the salt ending (chloride, nitrate or sulfate); balance the equation before calculating; include state symbols when requested; and show every subtraction in burette titre calculations.

8. The Periodic Table

Learning outcomes: explain how elements are arranged by proton number; interpret groups and periods using electronic configurations; describe the properties and trends of Group I, Group VII and noble gases; explain halogen displacement reactions; and recognise characteristic properties of transition elements.

8.1 Arrangement of the Periodic Table

The modern Periodic Table arranges elements in order of increasing proton (atomic) number. The horizontal rows are called periods; the vertical columns are called groups. Elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell (for the main-group elements).

• Period number: tells you the number of occupied electron shells in a simple atom.
• Group number: for Groups I–VII, tells you the number of outer-shell electrons. Group 0 (also called Group 18) contains noble gases with full outer shells.
• Across a period: proton number increases by one from one element to the next. The outer shell being filled remains the same shell.
• Down a group: atoms have an additional occupied electron shell at each step; atomic size generally increases.

8.2 Group I – the alkali metals

Lithium, sodium and potassium are soft, reactive metals. They have one electron in the outer shell and form +1 ions by losing that electron. They are stored under oil to prevent contact with oxygen and water. Their reactions with water become more vigorous down the group.

General equation: 2M + 2H₂O → 2MOH + H₂, where M represents an alkali metal. For example: 2Na + 2H₂O → 2NaOH + H₂. The solution becomes alkaline because a metal hydroxide forms, and hydrogen gas is released.

Why does reactivity increase down Group I? The outer electron is in a shell farther from the nucleus and is more shielded by inner electrons. The attraction between the nucleus and this outer electron becomes weaker, so it is lost more easily.

8.3 Group VII – the halogens

Fluorine, chlorine, bromine and iodine are non-metals. They exist as diatomic molecules (F₂, Cl₂, Br₂ and I₂), have seven outer-shell electrons, and gain one electron to form halide ions with a −1 charge. Their colours become darker down the group and their physical states change from gases to liquid to solid at room temperature.

Trend in reactivity: halogen reactivity decreases down Group VII. A halogen atom gains an electron during reaction. Down the group, the outer shell is farther from the nucleus and shielding increases, so the attraction for an incoming electron is weaker.

8.4 Halogen displacement reactions

A more reactive halogen displaces a less reactive halogen from a solution of its halide. Chlorine displaces bromide and iodide ions; bromine displaces iodide ions; iodine does not displace chloride or bromide ions.

8.5 Group 0 / Group 18 – noble gases

Helium, neon, argon, krypton and xenon are colourless, monatomic gases with very low chemical reactivity. Their outer electron shells are full (helium has a full first shell of two electrons; the others have eight outer-shell electrons). They do not normally form ions or compounds in the conditions studied at this level.

Trend down the group: noble gases have increasing boiling points down the group. Their atoms become larger and have more electrons, so temporary induced dipole attractions (London dispersion forces) become stronger. More energy is needed to separate the atoms.

8.6 Transition elements

Transition elements are metals found in the central block of the Periodic Table. Compared with Group I metals, they are generally harder, stronger and denser, have higher melting points, and are less reactive. Many form coloured compounds, show variable oxidation states, and act as catalysts.

• Variable oxidation states: an element can form ions with different charges, for example iron(II), Fe²⁺, and iron(III), Fe³⁺.
• Coloured compounds: many transition-metal compounds are coloured; for example, copper(II) compounds are often blue or blue-green.
• Catalysts: iron is used in the Haber process; nickel is used in hydrogenation; transition-metal compounds can catalyse other reactions.
• Physical properties: transition metals are typically strong, dense and have high melting points, although there are exceptions.

8.7 Comparing the important groups

8.8 Common misconceptions and exam technique

• Do not confuse group (vertical column) with period (horizontal row).
• Group VII halogens are diatomic molecules, not single atoms, when present as elemental substances.
• Do not say halogen reactivity increases down the group: it decreases.
• Do not say Group I reactivity decreases down the group: it increases.
• In displacement questions, compare the reactivity of the halogens, not the potassium or sodium ions in the salt.
• Balance displacement ionic equations and include state symbols when asked.
• Use “outer-shell electrons” rather than “number of shells” when explaining the group number of a main-group element.

8.9 Worked exam examples

8.10 Topic 8 exam practice

Attempt the questions before opening the marking points. Give clear explanations and balanced equations where required.

9. Metals

Learning outcomes: describe the physical properties of metals; use the reactivity series to predict reactions and displacement; explain the extraction of iron and aluminium; describe corrosion and rust prevention; and explain why alloys are often more useful than pure metals.

9.1 Physical properties and uses of metals

Most metals are good conductors of heat and electricity, are strong, malleable (can be hammered into shape) and ductile (can be drawn into wires). Many have high melting points and high densities, although there are exceptions such as mercury, which is liquid at room temperature, and some low-density metals.

9.2 The reactivity series

The reactivity series ranks metals by how readily they take part in reactions. A useful school-level order is:

A metal higher in the series can displace a less reactive metal from a solution of its salt. A metal above hydrogen can usually react with dilute hydrochloric acid or dilute sulfuric acid to produce hydrogen gas; metals below hydrogen generally do not. The exact reaction rate also depends on conditions and the metal's surface.

9.3 Reactions with water and dilute acids

General acid reaction: metal + dilute acid → salt + hydrogen. For example: Mg + 2HCl → MgCl₂ + H₂. With sulfuric acid: Mg + H₂SO₄ → MgSO₄ + H₂. The salt name depends on the acid: hydrochloric acid forms chlorides; sulfuric acid forms sulfates; nitric acid forms nitrates (nitric acid can behave differently with metals and should not be treated as a simple hydrogen-producing acid in all cases).

9.4 Displacement reactions

A more reactive metal displaces a less reactive metal from a solution containing its ions. For example, zinc displaces copper from copper(II) sulfate solution:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Ionic equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). Blue copper(II) sulfate solution becomes paler as Cu²⁺ ions are removed, and a reddish-brown copper deposit forms on the zinc. Copper cannot displace zinc from zinc sulfate because copper is less reactive.

9.5 Extraction of metals from ores

Most metals occur naturally in compounds called ores. The method used to extract a metal depends on its position in the reactivity series and the stability of its compounds.

• Very reactive metals such as aluminium are extracted by electrolysis of molten compounds. Carbon is not a sufficiently strong reducing agent for extracting aluminium from aluminium oxide.
• Less reactive metals such as iron can be extracted by reduction of their oxides using carbon or carbon monoxide in a furnace.
• Very unreactive metals such as gold may occur native (uncombined) in the Earth's crust.

9.6 Extraction of iron in the blast furnace

Iron is extracted from haematite, mainly iron(III) oxide, Fe₂O₃. A blast furnace is charged with iron ore, coke (carbon) and limestone (calcium carbonate). A blast of hot air enters near the bottom.

1. Coke burns in oxygen and releases heat: C + O₂ → CO₂.
2. Carbon dioxide reacts with hot coke to form carbon monoxide: CO₂ + C → 2CO.
3. Carbon monoxide reduces iron(III) oxide to iron: Fe₂O₃ + 3CO → 2Fe + 3CO₂.
4. Limestone thermally decomposes: CaCO₃ → CaO + CO₂.
5. Calcium oxide reacts with acidic silica impurities to form slag: CaO + SiO₂ → CaSiO₃.

9.7 Extraction of aluminium by electrolysis

Aluminium is extracted from purified aluminium oxide (alumina), Al₂O₃, dissolved in molten cryolite. Cryolite lowers the operating temperature and helps the mixture conduct electricity. The cell uses carbon electrodes; the carbon anodes are gradually consumed because oxygen produced at the anode reacts with hot carbon.

• Cathode (negative): Al³⁺ + 3e⁻ → Al. Molten aluminium forms and collects at the bottom.
• Anode (positive): 2O²⁻ → O₂ + 4e⁻. Oxygen reacts with carbon: C + O₂ → CO₂ (some CO may also form).

Overall simplified equation: 2Al₂O₃(l) → 4Al(l) + 3O₂(g). The carbon anodes must be replaced periodically. Large amounts of electrical energy are required, which makes aluminium extraction expensive.

9.8 Recycling metals

Recycling metals conserves finite mineral resources, reduces the need for mining, often uses less energy than extraction from ores, and reduces waste and environmental damage. However, collection, sorting, transport and purification also require energy and money. Aluminium recycling is especially valuable because primary aluminium extraction uses large amounts of electricity.

9.9 Corrosion and rusting

Corrosion is the gradual destruction of a metal by chemical reactions with substances in its environment. Rusting is the corrosion of iron. Iron rusts when both oxygen and water are present. Salt solutions speed up rusting because they increase electrical conductivity and help the electrochemical processes occur.

Rust is hydrated iron(III) oxide, often written as Fe₂O₃·xH₂O. Rust is porous and flakes away, exposing fresh iron to further corrosion.

9.10 Preventing rust and corrosion

9.11 Alloys

An alloy is a mixture of a metal with one or more other elements. Alloying changes the arrangement and sizes of atoms in the metal, making it harder for layers of atoms to slide over one another. Many alloys are therefore harder or stronger than the pure metal.

9.12 Common exam errors

• Do not state that all metals react with cold water or dilute acids. Compare the metal with hydrogen and use the specified conditions.
• In a displacement reaction, the more reactive metal is oxidised and its ions enter solution; the less reactive metal ions gain electrons and form the solid metal.
• Do not confuse extraction by reduction with extraction by electrolysis. Aluminium needs electrolysis; iron oxide can be reduced by carbon monoxide.
• Rusting requires both oxygen and water. Salt speeds up rusting but is not one of the two essential conditions.
• Galvanising uses zinc. Zinc can give sacrificial protection even when the coating is scratched.
• An alloy is a mixture, not a pure compound with a fixed chemical formula.

9.13 Topic 9 exam practice

Answer each question using equations and explanations where requested. Open the marking points only after attempting the question.

10. Chemistry of the environment

Learning outcomes: describe the composition of clean, dry air; identify major air pollutants and their sources and effects; explain the formation and consequences of acid rain and enhanced greenhouse effect; describe water pollution and treatment; and evaluate practical ways to reduce environmental damage.

10.1 The composition of clean, dry air

Air is a mixture of gases. The approximate composition of clean, dry air near sea level is about 78% nitrogen, 21% oxygen, 0.9% argon and 0.04% carbon dioxide, with other gases present in very small amounts. Water vapour varies with weather and location, so it is not included in the percentages for dry air.

10.2 Pollutants in the air

10.3 Complete and incomplete combustion

When a hydrocarbon burns in a plentiful supply of oxygen, complete combustion produces carbon dioxide and water. When oxygen is limited, incomplete combustion can produce carbon monoxide and/or carbon (soot), as well as water.

Complete combustion: CH₄ + 2O₂ → CO₂ + 2H₂O.

Incomplete combustion to carbon monoxide: 2CH₄ + 3O₂ → 2CO + 4H₂O.

Incomplete combustion to carbon: CH₄ + O₂ → C + 2H₂O.

10.4 Acid rain

Sulfur dioxide and nitrogen oxides dissolve in atmospheric water and undergo reactions that produce acids. These acidic substances can return to the Earth's surface in rain, snow, fog or dry deposition. Acid deposition can lower the pH of lakes and soils, harm aquatic organisms and plants, and accelerate the corrosion of limestone buildings and some metals.

Representative equations include:

• Sulfur dioxide dissolves and can be oxidised to sulfuric acid: 2SO₂ + O₂ + 2H₂O → 2H₂SO₄.
• Nitrogen dioxide can form nitric acid in the presence of oxygen and water: 4NO₂ + O₂ + 2H₂O → 4HNO₃.
• Acids react with calcium carbonate in limestone: CaCO₃ + 2H⁺ → Ca²⁺ + CO₂ + H₂O.

10.5 The greenhouse effect and climate change

The natural greenhouse effect keeps Earth's surface warmer than it would otherwise be. Short-wave radiation from the Sun passes through the atmosphere and warms the surface. The surface emits infrared radiation; greenhouse gases absorb and re-emit some of this outgoing infrared energy. Human activities increase concentrations of greenhouse gases, strengthening this effect and contributing to global warming and climate change.

Possible consequences of climate change include rising average temperatures, changing rainfall patterns, more frequent or intense heat extremes in many regions, melting glaciers and ice sheets, sea-level rise and impacts on ecosystems, agriculture, water availability and human health. The exact effects vary by region.

10.6 Water pollution

10.7 Eutrophication

Eutrophication occurs when excess nutrients, especially nitrates and phosphates, enter a water body. These nutrients promote rapid algal growth. When algae die, decomposer microorganisms break down the organic matter and use dissolved oxygen. Oxygen levels may fall so low that fish and other aquatic organisms die.

10.8 Water treatment and purification

Water treatment makes water suitable for its intended use. A typical drinking-water treatment plant may use several stages. The exact sequence varies according to the quality of the source water and local regulations.

1. Screening: removes large debris such as leaves and rubbish.
2. Sedimentation: allows heavier suspended particles to settle; a coagulant such as aluminium sulfate may be added to help fine particles clump together.
3. Filtration: water passes through sand, gravel or other filter media to remove remaining suspended particles.
4. Disinfection: chlorine, ozone or ultraviolet treatment is used to reduce harmful microorganisms. A controlled chlorine residual may protect water in distribution systems.
5. Storage and testing: treated water is checked against safety standards before supply.

Important: Boiling kills many disease-causing microorganisms but does not remove dissolved salts, heavy metals or many chemical pollutants. Filtration alone does not guarantee that water is microbiologically safe. Safe treatment requires methods appropriate to the contaminants present.

10.9 Reducing environmental pollution

10.10 Practical investigation: measuring water quality

Water quality can be investigated using simple measurements, but no single test proves that water is safe to drink.

• pH: use a calibrated pH meter or suitable indicator to test acidity or alkalinity.
• Turbidity: compare clarity using a turbidity tube or meter. Cloudiness suggests suspended particles but does not identify the contaminant.
• Dissolved oxygen: use a dissolved-oxygen probe or approved test kit. Low dissolved oxygen can indicate organic pollution or eutrophication.
• Nitrate / phosphate: use suitable test kits to detect nutrient enrichment.
• Microbiological testing: trained personnel use approved methods to test for indicator organisms such as E. coli. Do not culture unknown environmental samples at home or taste test water.

10.11 Interpreting environmental data

In exam questions, identify the trend, quote supporting data and then explain a chemical or biological mechanism. For example, if a graph shows rising carbon dioxide concentration and rising average temperature over time, state that both increase, quote values or dates, and explain that carbon dioxide absorbs outgoing infrared radiation. A correlation alone does not prove that one measured variable is the only cause of the other.

10.12 Common misconceptions and exam technique

• Air is a mixture, not a compound. Nitrogen is the largest component of dry air, not oxygen.
• Carbon monoxide is formed by incomplete combustion and is poisonous; carbon dioxide is a greenhouse gas.
• Acid rain is mainly linked to sulfur dioxide and nitrogen oxides, not directly to carbon dioxide.
• The natural greenhouse effect is necessary for a habitable Earth; the enhanced effect results from increased greenhouse-gas concentrations.
• Ozone depletion and climate change are separate issues, although some substances can affect both.
• Eutrophication begins with nutrient enrichment; oxygen depletion occurs when decomposers break down excess organic matter.
• Water that looks clear is not necessarily safe to drink. Microorganisms and dissolved chemicals may be invisible.
• When evaluating a control method, mention both its benefit and a possible limitation, cost or energy requirement.

10.13 Topic 10 exam practice

Attempt all questions. Use balanced equations, data and scientific explanations where required.

11. Organic chemistry

Course unit map: fuels and crude oil; alkanes and alkenes; homologous series; functional groups; reactions and tests; alcohols; carboxylic acids; polymers and addition polymerisation.

Displayed formulae, reaction pathways, polymer diagrams and structured practice are pending development.

12. Experimental techniques and chemical analysis

Course unit map: experimental design and apparatus; acid–base titrations; chromatography; filtration, crystallisation and distillation; identification of ions and gases; flame tests; laboratory safety and data handling.

Fully labelled apparatus diagrams, observation tables, separation-method comparisons and practical examination tasks are pending development.