Chemical Reactions
Every reaction here can be interrogated with the same four questions: is it a chemical change, how fast does it go, does it stop halfway, and where are the electrons moving?
Physical vs chemical changes
Every change falls into one of two categories. In a physical change, matter is only rearranged; in a chemical change, it is transformed — new substances appear.
Physical changes
A physical change alters how a substance looks or where it sits but not what it is. Water freezing into ice is still H₂O; sugar dissolving in tea is still sucrose. The particles are the same as before — only their arrangement, spacing or partners have shifted, and the change is usually easy to reverse. Common examples: state changes, dissolving, and simple mixing.
Chemical changes
A chemical change produces one or more new substances with different properties from the reactants, and reversing it usually needs a separate reaction, not a physical trick — rusting iron doesn’t turn back into iron and oxygen by cooling. Four signs suggest a reaction has happened.
| Sign | Example |
|---|---|
| Colour change | Copper turning silver-salt solution blue |
| Temperature change | Neutralisation warms the flask; NH₄NO₃ dissolving cools it |
| Effervescence (fizzing) | Metal + acid releasing hydrogen gas |
| Precipitate forms | A soluble salt reacting to make an insoluble one |
A single sign is usually enough to be confident. Two or more together settle the matter.
Rate of reaction — the four factors
Four factors change how fast a reaction proceeds: the concentration of reactants (or pressure of gases), the surface area of solids, the temperature, and whether a catalyst is present.
Collision theory
For a reaction to happen, reactant particles must collide — and not just any collision will do. They need to hit with enough energy to break existing bonds (exceeding the activation energy Ea) and with the correct orientation. Anything that increases the number of successful collisions per second increases the rate.
The four factors
| Factor (increased) | Effect | Why |
|---|---|---|
| Concentration / gas pressure | Rate increases | More particles per unit volume → more collisions per second |
| Surface area (smaller pieces) | Rate increases | More reactant surface exposed → more particles can collide at once |
| Temperature | Rate increases | Particles move faster → more collisions, and more exceed Ea |
| Catalyst | Rate increases | Provides an alternative route with a lower Ea → more collisions succeed |
Temperature is the strongest of the four: roughly, a 10 °C rise doubles the rate, because a small temperature increase pushes a large fraction of collisions over the activation-energy threshold.
Catalysts don’t change the amount of product formed — only how quickly it gets there. On a reaction pathway diagram a catalyst lowers the Ea peak but leaves the reactant and product levels, and therefore ΔH, unchanged.
Measuring & interpreting rate
Two standard experiments measure how fast a reaction is happening: watching mass fall, or watching gas volume rise. Either way, the raw data is a curve on a graph — and the rate is its gradient.
Two ways to measure
When a reaction produces a gas, either method works. If the gas escapes into the air, place the flask on an electronic balance and record the falling mass at fixed intervals. If the gas is collected, attach a gas syringe (or an inverted measuring cylinder over water) and record the rising volume. In both cases time goes on the x-axis and the measured quantity on the y-axis. The curve is steep at the start (fast rate), bends over as reactants are used up, and plateaus when the reaction is complete.
Reading a rate–time graph
The rate at any moment is the gradient of the curve at that point. It is highest at t = 0 and falls to zero at the plateau. Comparing two conditions on the same axes: the steeper the initial gradient, the faster the reaction; the sooner it plateaus, the sooner it is complete; the final height gives the total amount of product.
An overall rate can be calculated as rate = amount of product formed ÷ time, in cm³/s, g/s or mol/s depending on what is measured.
ExtendedLight and enzymes
Two special cases sit alongside the four main factors. Some reactions require light as their energy source rather than heat: photosynthesis captures sunlight into glucose; silver halides on old camera film darken when exposed to light. Others are catalysed by enzymes — biological catalysts made of protein, extremely efficient at body temperature but denatured if the temperature climbs too high.
Reversible reactions
Not every reaction goes to completion. Some reach a point where the products themselves start converting back into the reactants — and the arrow between them has to run both ways.
The ⇌ symbol
A single-arrow equation A + B → C + D says the reactants are converted into products until one runs out. A double-arrow equation A + B ⇌ C + D says the reaction can go either way: C and D can react to give A and B back. The forward and reverse steps have equal magnitude but opposite sign of ΔH — if the forward reaction releases 100 kJ, the reverse absorbs 100 kJ.
Hydrated salts — the test for water
Two reversible reactions between coloured salts and water underpin the standard laboratory tests for water. The hydrated salt contains water of crystallisation locked into its structure; the anhydrous salt has none. Copper(II) sulfate: CuSO₄·5H₂O (blue) ⇌ CuSO₄ (white) + 5H₂O. Cobalt(II) chloride: CoCl₂·6H₂O (pink) ⇌ CoCl₂ (blue) + 6H₂O.
Heating either hydrated salt drives off the water — the endothermic forward direction, giving the anhydrous solid. Adding water reverses the change (exothermic) and restores the colour. So the standard test for water is a colour change from white to blue on anhydrous CuSO₄, or from blue to pink on anhydrous CoCl₂.
Equilibrium & industry
In a closed vessel, a reversible reaction eventually reaches dynamic equilibrium: both directions run at the same rate and nothing appears to change. Applying pressure, changing temperature or adjusting concentration all shift where that balance sits.
ExtendedLe Chatelier’s principle
| Change applied | Equilibrium shifts to |
|---|---|
| Increase reactant concentration | The right — using up the added reactant |
| Increase temperature | The endothermic direction — absorbing the added heat |
| Increase pressure (gases) | The side with fewer gas moles — reducing total pressure |
| Add a catalyst | No shift — it only speeds up reaching equilibrium |
ExtendedThe Haber process
The industrial synthesis of ammonia, the source of nitrogen fertilisers: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), forward exothermic, 4 moles of gas → 2 moles. Conditions: 450 °C, 200 atm, iron catalyst. A lower temperature would push equilibrium further right (favouring products in an exothermic reaction) but slow the rate too much; a higher pressure would push it right (fewer gas moles on the product side) but is expensive and dangerous. 450 °C and 200 atm are the compromise.
ExtendedThe Contact process
The industrial route to sulfuric acid via sulfur trioxide: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), forward exothermic, 3 moles of gas → 2 moles. Conditions: 450 °C, 2 atm, vanadium(V) oxide (V₂O₅) catalyst. The relatively low pressure is enough because the equilibrium already lies well to the right.
Redox — oxygen & electron transfer
The oldest definition of oxidation and reduction is the simplest: follow the oxygen. A substance is oxidised when it gains oxygen and reduced when it loses it. Because oxygen has to come from somewhere and go somewhere, the two always occur together — a redox reaction.
Standard examples
Extraction of iron in a blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂. Iron(III) oxide loses oxygen (reduced to iron); carbon monoxide gains oxygen (oxidised to CO₂). Magnesium in copper(II) oxide: Mg + CuO → MgO + Cu. Magnesium is oxidised, copper(II) oxide is reduced — a displacement reaction that is also redox. Any combustion is redox too: the fuel is oxidised, the O₂ supplies the oxygen.
ExtendedThe electron view
The deeper definition is about electrons, not oxygen — some redox reactions have no oxygen at all. Consider iron displacing copper: Fe + Cu²⁺ → Fe²⁺ + Cu (the sulfate is a spectator). The Fe atom loses 2 electrons (oxidised); the Cu²⁺ ion gains 2 electrons (reduced). OIL RIG — Oxidation Is Loss, Reduction Is Gain of electrons.
ExtendedOxidation state & agents
The oxidation state tracks each atom’s charge if all bonds were ionic. An increase means oxidation, a decrease means reduction: going from FeCl₃ (iron +3) to FeCl₂ (iron +2), iron has been reduced. An oxidising agent causes another substance to be oxidised and is itself reduced; a reducing agent gives up its electrons and is itself oxidised.
ExtendedTests using colour change
| Test | Colour change | Detects |
|---|---|---|
| Acidified KMnO₄ | Purple → colourless | A reducing agent (it reduces MnO₄⁻) |
| Aqueous KI | Colourless → brown | An oxidising agent (it oxidises I⁻ to I₂) |
Exam advice
Common mistakes
Model answer
Recall checklist
- Distinguish physical from chemical change and list the four common signs of a reaction.
- Name the four factors that affect rate and explain each using collision theory.
- Describe two experimental methods to measure rate and interpret a rate–time graph.
- Use the ⇌ symbol and give the equations and colour changes for the two hydrated-salt tests for water.
- Define dynamic equilibrium and apply Le Chatelier’s principle.
- State the equations and conditions for the Haber and Contact processes and justify each condition.
- Define oxidation and reduction in terms of oxygen and of electrons.
- Identify the oxidising agent and reducing agent in a redox reaction.
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