Part 1 · Molecules and chemical bonds

• Atoms join together by chemical bonds to form molecules.
• A molecule is a group of two or more atoms chemically bonded together.
• In a chemical reaction, bonds in the reactants are broken.
• New bonds then form to create the products.
• No atoms are created or destroyed — they are rearranged.

Part 2 · Chemical symbols

• Every element has a unique letter (or two-letter) symbol on the periodic table.
• The first letter is always a capital; the second (if present) is lower case.
• Most symbols match the English name: e.g. oxygen = O, magnesium = Mg.
• Some come from Latin names: e.g. iron = Fe (Ferrum), gold = Au (Aurum).
• Scientists always have a periodic table available, so symbols need not be memorised.

Part 3 · Diatomic molecules and chemical coefficients

Some elements exist as diatomic molecules — two atoms of the same element bonded together.

• The 7 common diatomic elements: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂.
• Mnemonic: 'Have No Fear Of Ice Cold Beer' (H, N, F, O, I, Cl, Br)
• Subscript number (written below): number of atoms of that element in one molecule.
• Chemical coefficient (written before formula): number of molecules/formula units.
• Example: 2Mg + O₂ → 2MgO (coefficient 2 means 2 magnesium atoms; O₂ means 2 oxygen atoms bonded).

Part 1 · State symbols and interpreting formulae

State symbols added in brackets describe the physical state of each substance.

• (s) = solid  |  (l) = liquid  |  (g) = gas  |  (aq) = aqueous (dissolved in water)
• Subscript numbers apply to the atom immediately before them: H₂O = 2H + 1O.
• Brackets show groups of atoms: Ca(OH)₂ = 1 Ca, 2 O, 2 H.
• Numbers after brackets multiply everything inside: Al₂(SO₄)₃ = 2 Al, 3 S, 12 O.

Part 2 · Balancing equations — principles

• A balanced equation has equal numbers of each type of atom on both sides.
• Only coefficients (numbers in front of formulas) may be changed.
• Subscripts are fixed by the chemical formula and must NEVER be altered.
• Step 1: write the unbalanced equation. Step 2: count atoms on each side. Step 3: adjust coefficients. Step 4: check.
• Example: Zn + O₂ → ZnO → Balanced: 2Zn + O₂ → 2ZnO

Part 3 · Balancing equations — practice

Use the systematic method: list atoms → adjust coefficients → recheck.

Part 1 · Collision theory

For a reaction to happen, reactant particles must collide with sufficient energy.

• The collision must have energy equal to or greater than the activation energy.
• Activation energy: the minimum energy needed for a successful collision.
• No collision → No reaction.
• Collision with insufficient energy → No reaction.
• Collision with sufficient energy (≥ activation energy) → Reaction.
• The more successful collisions per second, the faster the rate of reaction.

Part 2 · Measuring rate of reaction

Rate of reaction = how quickly reactants are converted to products.

• Rate (cm³/s) = amount of product produced (cm³) ÷ time (s)
• Reaction studied: Mg + 2HCl → MgCl₂ + H₂
• Hydrogen gas produced is collected in a gas syringe and volume recorded.
• Rate is fastest at the start (most reactant present); slows as reactant is used up.

Part 3 · Practical: rate of reaction data

• Record the volume of gas produced every 10 seconds.
• Rate is highest at the start — the reactant concentration is greatest.
• Rate slows as reactants are consumed.
• Reaction stops when one reactant is completely used up.

Part 1 · Factors that affect rate

Rate increases if more successful collisions happen per second.

• Achieved by: (a) making collisions more frequent, OR (b) increasing particle energy.
• Four main factors: temperature, concentration, surface area, catalyst.
• Increasing concentration → more particles in same volume → more collisions.
• Increasing pressure (gases) → particles closer → more collisions.
• Increasing surface area → more particles exposed → more collisions.
• Adding a catalyst → lowers activation energy → more collisions are successful.

Part 2 · Effect of concentration

Concentration = amount of substance dissolved in a given volume of solution.

• A more concentrated solution has more particles in the same volume.
• More particles → more frequent collisions between reactant particles.
• More frequent collisions → more successful collisions per second.
• Therefore: increasing concentration increases the rate of reaction.

Part 1 · Effect of temperature on particles

• Heating particles increases their kinetic energy.
• With more kinetic energy, particles move faster in all directions.
• Faster particles collide more frequently.
• More particles also have enough energy to equal or exceed the activation energy.
• Both effects increase the number of successful collisions → faster rate.

Part 2 · Practical: sodium thiosulfate and HCl

Reaction: Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + H₂O(l) + S(s)

• Sulfur precipitate makes the solution go cloudy.
• A cross drawn under the flask disappears when enough sulfur has formed.
• Time for cross to disappear is measured at different temperatures.
• Higher temperature → less time → faster rate.

Part 1 · Surface area and collisions

For a solid reactant, only particles at the surface can collide with solution particles.

• Greater surface area → more particles exposed → more collisions possible.
• More collisions → more successful collisions per second → faster rate.
• Powders have much greater surface area than large lumps of the same substance.
• Powders therefore react faster than large pieces.

Part 2 · Practical: CaCO₃ and HCl

Reaction: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l)

• CO₂ gas is the product that can be collected and measured.
• Test forms of CaCO₃: powder, small chips, and large chips.
• Measure volume of CO₂ every few seconds for each form.
• The form producing gas fastest has the greatest surface area and highest rate.

Part 1 · What is a catalyst?

• A catalyst speeds up a chemical reaction without being used up.
• Catalysts can therefore be used over and over again.
• Catalysts work by providing an alternative reaction pathway.
• This alternative pathway has a lower activation energy.
• Lower activation energy → more collisions have sufficient energy to react.
• More successful collisions per second → faster rate of reaction.

Part 2 · Practical: hydrogen peroxide decomposition

Reaction: 2H₂O₂(aq) → 2H₂O(l) + O₂(g)

• Catalysts tested: iron oxide, manganese dioxide, potato, liver.
• Method: add catalyst to H₂O₂ + washing up liquid; measure foam height after 20 s.
• Foam height indicates volume of O₂ produced. Rate = volume of foam ÷ time (20 s).

Part 3 · Catalysts in industry

• Industrial processes aim to make as much product as cheaply as possible.
• Catalysts speed up reactions → product made faster → greater profit.
• Catalysts lower activation energy → less heating needed → lower energy costs.
• Catalysts are not used up → reused many times → lower material costs.

Part 1 · Energy in chemical reactions: bond breaking and making

• In every chemical reaction, bonds in the reactants must be broken.
• Breaking bonds REQUIRES energy (takes energy IN from surroundings).
• New bonds then form in the products.
• Making bonds RELEASES energy (gives energy OUT to surroundings).
• The overall energy change depends on which process requires more energy.

Part 2 · Endothermic reactions

In an endothermic reaction, more energy is needed to break bonds than is released making bonds.

• The reaction absorbs energy from the surroundings.
• The surroundings become COOLER — temperature decreases.
• Energy is transferred FROM the surroundings INTO the reaction.
• Examples: ice packs, dissolving ammonium nitrate, citric acid + sodium hydrogen carbonate.

Part 1 · Exothermic reactions

In an exothermic reaction, more energy is released making bonds than is needed to break bonds.

• The surplus energy is transferred to the surroundings.
• The surroundings become WARMER — temperature increases.
• Energy is transferred FROM the reaction INTO the surroundings.
• Examples: combustion, neutralisation, hand warmers, metals reacting with acid.

Part 2 · Identifying reaction type from temperature change

• Temperature INCREASE → exothermic reaction.
• Temperature DECREASE → endothermic reaction.
• Exothermic: energy released making bonds > energy needed to break reactant bonds.
• Endothermic: energy needed to break bonds > energy released making bonds.

Part 1 · What is a reaction profile?

A reaction profile (energy profile diagram) shows energy changes during a reaction.

• Y-axis = potential energy; X-axis = progress of reaction (reaction pathway).
• All reactions show an initial energy peak — this represents the activation energy.
• Endothermic: products are at a HIGHER energy level than reactants.
• Exothermic: products are at a LOWER energy level than reactants.
• The height of the peak above the reactant line = activation energy.

Part 2 · Drawing and labelling reaction profiles

• Key labels: Reactants (left level), Activation energy (peak), Products (right level), Energy absorbed/released.
• Endothermic profile: start low, rise to peak, finish HIGHER than start.
• Exothermic profile: start high, rise to peak, finish LOWER than start.
• A catalyst lowers the peak (activation energy) but does NOT change reactant or product energy levels.

Part 1 · What is thermal decomposition?

Thermal decomposition: heat breaks a compound into smaller/simpler molecules.

• Example: calcium carbonate → calcium oxide + carbon dioxide
• CaCO₃ → CaO + CO₂
• Generally endothermic — more energy needed to break bonds than is released.
• Occasionally exothermic — released energy can trigger thermal runaway (explosions).

Part 2 · Identifying thermal decomposition reactions

• Thermal decomposition has ONLY ONE reactant, which is heated.
• Products are simpler/smaller than the reactant.
• Signs: colour change, gas production (test CO₂ with limewater — goes milky).
• Copper carbonate (green) + heat → copper oxide (black) + CO₂

Part 3 · Energy changes in thermal decomposition

• Thermal decomposition of CaCO₃ is endothermic.
• Energy is absorbed from the heat source (Bunsen flame) into the reaction.
• More energy is needed to break bonds in CaCO₃ than is released forming CaO and CO₂.
• Reaction profile: products (CaO + CO₂) at a higher energy level than reactants (CaCO₃).

Part 1 · Specific heat capacity

Specific heat capacity (c): energy needed to raise 1 g of a substance by 1°C.

• Units: J / (°C g)  |  Symbol: c
• Low c → heats up quickly (e.g. copper: c = 0.385 J/°C g).
• High c → heats up slowly (e.g. water: c = 4.18 J/°C g).
• This is why food cooks faster in oil than in water.

Part 2 · Calculating energy change: Q = mcΔT

Heat energy change (Q) is calculated using: Q = mcΔT

• Q = heat energy change (J)
• m = mass of liquid (g)
• c = specific heat capacity (J / °C g)
• ΔT = temperature change (°C)
• Example: Q = 50 × 4.18 × 10 = 2090 J (50 g water, ΔT = 10°C)

Part 3 · Practical: measuring energy change (HCl + Mg)

• Reaction: 2HCl + Mg → MgCl₂ + H₂
• Use a polystyrene cup + lid to minimise heat loss.
• Polystyrene is a poor conductor — reduces energy lost to surroundings.
• Record start temperature, add Mg, record maximum temperature; calculate ΔT and Q.