Chemical Changes (Reactivity, Electrolysis, Acids)

Quick-Fire Definitions

Oxidation: Gain of oxygen OR loss of electrons (OIL).
Reduction: Loss of oxygen OR gain of electrons (RIG).
Displacement: A more reactive element replaces a less reactive element from a compound.
Electrolyte: A molten or dissolved ionic compound that conducts electricity.
Cathode: The negative electrode - attracts cations; reduction occurs here.
Anode: The positive electrode - attracts anions; oxidation occurs here.
Strong acid: An acid that fully ionises in water (e.g. HCl, H₂SO₄, HNO₃).
Weak acid: An acid that only partially ionises in water (e.g. Ethanoic acid, citric acid).

4.4.1 Metal Oxides

When metals react with oxygen, they form metal oxides. The vigour of this reaction depends on the metal's position in the reactivity series.

2Mg(s) + O₂(g) → 2MgO(s)

Magnesium burns in air with a bright white flame to produce the white powder magnesium oxide.

At this stage, we define oxidation as the gain of oxygen and reduction as the loss of oxygen. In the reaction above, magnesium is oxidised (it gains oxygen).

CuO(s) + C(s) → Cu(s) + CO₂(g)

Here, copper oxide is reduced (it loses oxygen) and carbon is oxidised (it gains oxygen). The carbon acts as a reducing agent.

Metal oxides are basic oxides. They react with acids in neutralisation reactions to produce a salt and water.

The Reactivity Series

The reactivity series ranks metals in order of how vigorously they react. A metal's reactivity is determined by how readily it loses its outer shell electrons to form positive ions (cations).

Order (most to least reactive)

The reactivity series showing common extraction methods. Carbon and hydrogen are included for reference.

Carbon and hydrogen are included in the reactivity series even though they are non-metals. Carbon is used as a benchmark for extraction methods, and hydrogen is used to compare acid reactions: only metals above hydrogen react with dilute acids.

Reactions with water and dilute acid

Potassium: reacts violently with cold water, ignites with a lilac flame. Too dangerous for acid reactions.
Sodium: vigorous fizzing with cold water, melts into a ball on the surface. Too dangerous for acid reactions.
Calcium: steady bubbling with cold water, producing a cloudy solution of calcium hydroxide.
Magnesium: very slow reaction with cold water. Reacts vigorously with steam and with dilute acid (rapid fizzing).
Zinc: no reaction with cold water. Slow, steady bubbling with dilute acid.
Iron: no reaction with cold water. Very slow reaction with dilute acid.
Copper: no reaction with cold water or dilute acid.

Metal	Reaction with Water	Reaction with Dilute Acid
Potassium	Violent - lilac flame	Too dangerous
Sodium	Vigorous fizzing, melts	Too dangerous
Calcium	Steady bubbling, cloudy solution	Vigorous fizzing
Magnesium	Very slow with cold water	Rapid fizzing
Zinc	No reaction	Slow, steady bubbles
Iron	No reaction	Very slow
Copper	No reaction	No reaction

Reactivity series - observations with water and dilute acid.

Remember the mnemonic Please Stop Letting Cows Moo All Continuously Zipping In Heavy Copper Shoes Going for the order of the reactivity series. Other mnemonics work just as well; use whichever sticks for you.

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See where each metal sits on our Interactive Periodic Table and compare their properties.

Displacement Reactions

A more reactive metal can displace a less reactive metal from a compound in solution. This happens because the more reactive metal has a greater tendency to form positive ions.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

When an iron nail is placed in blue copper sulfate solution, the solution gradually fades from blue to pale green (iron(II) sulfate), and a reddish-brown coating of copper metal appears on the nail.

Predicting displacement

Will zinc react with magnesium chloride solution?

Step 1: Check the reactivity series. Magnesium is above zinc.

Step 2: A less reactive metal cannot displace a more reactive metal. Therefore, no reaction occurs.

Predicting displacement (reaction occurs)

What happens when magnesium ribbon is placed in zinc sulfate solution?

Step 1: Magnesium is above zinc in the reactivity series.

Step 2: A more reactive metal displaces a less reactive metal. So magnesium will displace zinc.

Equation: Mg(s) + ZnSO₄(aq) → MgSO₄(aq) + Zn(s)

Observations: The magnesium dissolves. A grey coating of zinc metal appears. The solution warms up (exothermic).

If asked to predict whether a displacement reaction occurs, always compare the two metals in the reactivity series. The more reactive metal must be the one being added, not the one already in solution.

Extraction of Metals

Most metals are found in the Earth's crust as ores, which are rocks containing enough metal to make extraction economically worthwhile.

Method depends on reactivity

Metals below carbon (zinc, iron, copper): extracted by reduction with carbon in a blast furnace.
Metals above carbon (aluminium, sodium, potassium): too reactive for carbon reduction, so must be extracted by electrolysis.
Very unreactive metals (gold, platinum): found native in the Earth's crust and need no chemical extraction.

Reduction with carbon

2Fe₂O₃(s) + 3C(s) → 4Fe(l) + 3CO₂(g)

The carbon removes oxygen from the metal oxide (reducing it), and is itself oxidised to carbon dioxide. This is a redox reaction.

Biological methods

Phytomining uses plants that absorb metal compounds from the soil. The plants are harvested and burned, and copper is extracted from the ash. This is useful for low-grade ores that are uneconomical to mine traditionally.

Bioleaching uses bacteria to produce a leachate solution containing dissolved metal compounds. The copper ions can then be recovered from the leachate, for example by displacement with scrap iron or by electrolysis.

Phytomining and bioleaching are slower than traditional mining but are more environmentally friendly. They allow metals to be extracted from low-grade ores, reducing the need for quarrying.

Choosing the extraction method

How would you extract (a) iron, (b) aluminium, (c) gold from their ores?

(a) Iron is below carbon in the reactivity series → extract by reduction with carbon (blast furnace). 2Fe₂O₃ + 3C → 4Fe + 3CO₂

(b) Aluminium is above carbon → too reactive for carbon reduction, so extract by electrolysis of Al₂O₃ dissolved in molten cryolite.

(c) Gold is very unreactive → found native (uncombined) in the Earth's crust. No chemical extraction needed.

Oxidation & Reduction (OIL RIG)

At a basic level, oxidation means gaining oxygen and reduction means losing oxygen. These complementary processes always happen together in a redox reaction.

Redox in Terms of Electron Transfer Higher Tier

For Higher Tier, oxidation and reduction are defined more precisely in terms of electron transfer:

Oxidation Is Loss of electrons (OIL).
Reduction Is Gain of electrons (RIG).

A substance that is oxidised is the reducing agent (it causes another substance to be reduced by donating electrons). A substance that is reduced is the oxidising agent.

Half-equations

In the displacement of copper by magnesium:

Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s)

This can be split into two half-equations showing the electron transfer:

Mg(s) → Mg²⁺(aq) + 2e⁻ (oxidation)

Cu²⁺(aq) + 2e⁻ → Cu(s) (reduction)

Magnesium loses two electrons (oxidised). Copper ions gain two electrons (reduced). The electrons lost by one species are gained by the other.

In ionic half-equations, the number of electrons lost in the oxidation half-equation must equal the number gained in the reduction half-equation. Always check the charges balance on both sides.

4.4.2 Reactions of Acids

Acids react with metals, metal oxides, metal hydroxides, and metal carbonates. The type of salt produced depends on which acid is used.

Acid + Metal → Salt + Hydrogen

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

Only metals above hydrogen in the reactivity series react with dilute acids. The hydrogen gas can be tested with a lighted splint, which produces a squeaky pop.

Acid + Metal Oxide → Salt + Water

CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)

Acid + Metal Hydroxide → Salt + Water

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

Acid + Metal Carbonate → Salt + Water + CO₂

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

Carbon dioxide can be tested by bubbling it through limewater, which turns milky (cloudy white).

Naming salts

The salt produced depends on the acid used: hydrochloric acid → chloride salts, sulfuric acid → sulfate salts, nitric acid → nitrate salts. The metal in the salt comes from the metal, metal oxide, metal hydroxide or metal carbonate.

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Struggle with salt names? Common naming errors are covered in 10 Most Common Mistakes in GCSE Chemistry Exams.

Predicting the products: Zn + H₂SO₄

Step 1: Identify the acid. Sulfuric acid → produces a sulfate salt.

Step 2: Identify the metal. Zinc. So the salt is zinc sulfate (ZnSO₄).

Step 3: Metal + acid → salt + hydrogen.

Balanced equation: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

Observation: Zinc dissolves, bubbles of gas produced (squeaky pop with lighted splint).

Neutralisation

Neutralisation is the reaction between an acid and a base (or alkali) to produce a salt and water. A base is any substance that neutralises an acid. An alkali is a soluble base that produces hydroxide ions (OH⁻) in solution.

acid + base → salt + water

In terms of H⁺ and OH⁻ ions:

H⁺(aq) + OH⁻(aq) → H₂O(l)

The hydrogen ions from the acid combine with the hydroxide ions from the base to form water. This is why the solution becomes less acidic (or less alkaline) during neutralisation.

This ionic equation for neutralisation is one of the most commonly tested equations. Always include the state symbols. Remember: an alkali is a soluble base.

Making Soluble Salts Required Practical

To make a pure, dry sample of a soluble salt from an insoluble reactant (such as a metal oxide, hydroxide, or carbonate):

Warm dilute acid in a beaker using a Bunsen burner.
Add the insoluble base (metal oxide, hydroxide, or carbonate) in small amounts, stirring after each addition, until no more dissolves and excess solid remains.
Filter the mixture to remove the excess unreacted solid. The filtrate is the salt solution.
Pour the filtrate into an evaporating basin. Heat gently until about half of the water has evaporated.
Leave the solution to cool and crystallise slowly. Pat the crystals dry with filter paper.

The key to this practical is using excess base. This ensures all the acid reacts, so the final salt solution is pure (not contaminated with unreacted acid). The excess solid is then removed by filtration.

This is Required Practical 1 in the AQA specification. You need to know the method, the apparatus used, and why each step is performed.

Naming the salt: What salt is produced when magnesium oxide reacts with hydrochloric acid?

Step 1: The metal comes from material MgO → magnesium.

Step 2: Hydrochloric acid produces chloride salts.

Step 3: The salt is magnesium chloride (MgCl₂).

Full equation: MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)

pH Scale & Indicators

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is.

pH 0 to 6: Acidic (lower pH = stronger acid). Acids produce H⁺ ions in solution.
pH 7: Neutral.
pH 8 to 14: Alkaline (higher pH = stronger alkali). Alkalis produce OH⁻ ions in solution.

Universal indicator is a mixture of dyes that changes colour across the full pH range (red for strong acid, green for neutral, purple for strong alkali). A pH probe connected to a pH meter gives a more precise numerical reading.

A single indicator such as litmus only tells you whether a solution is acidic or alkaline. Universal indicator or a pH probe is needed to determine the strength of the acid or alkali.

Strong & Weak Acids

Full Ionisation vs Partial Ionisation Higher Tier

Strong acids

Strong acids fully ionise (dissociate) in water. Every molecule splits to release H⁺ ions. This is an irreversible process.

Examples: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), nitric acid (HNO₃).

HCl(aq) → H⁺(aq) + Cl⁻(aq)

Weak acids

Weak acids only partially ionise in water. A reversible equilibrium exists, with most molecules remaining undissociated.

Examples: ethanoic acid (CH₃COOH), citric acid, carbonic acid.

CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq)

Strong vs concentrated

These terms mean different things:

Strong/weak refers to the degree of ionisation (how much the acid splits into ions).
Concentrated/dilute refers to the amount of acid dissolved per unit volume of solution.

A dilute strong acid and a concentrated weak acid are both valid concepts.

pH and hydrogen ion concentration

The pH scale is logarithmic. As the pH decreases by 1 unit, the hydrogen ion concentration increases by a factor of 10. A decrease of 2 pH units means the H⁺ concentration is 100 times greater.

At the same concentration, a strong acid will have a lower pH than a weak acid because more H⁺ ions are present in solution.

Comparing H⁺ concentration

Two acids both have a concentration of 0.1 mol/dm³. Acid A has pH 1, Acid B has pH 3. Which is the strong acid?

Step 1: pH 1 is lower than pH 3, so Acid A has a higher H⁺ concentration.

Step 2: The pH difference is 2 units. The pH scale is logarithmic, so Acid A has 10² = 100 times more H⁺ ions than Acid B.

Step 3: At the same concentration, the acid with more H⁺ ions is the strong acid. Acid A is strongly ionised; Acid B is weakly ionised.

A common 6-mark question asks you to compare strong and weak acids. Always mention: degree of ionisation, reversible vs irreversible, relative pH at the same concentration, and rate of reaction with metals.

Titrations HT / Required Practical

Required Practical 2: Neutralisation by Titration Higher Tier

A titration is used to find the exact volume of acid needed to neutralise a known volume of alkali (or vice versa). This is Required Practical 2.

Method

Use a pipette to measure a fixed volume of alkali (e.g. 25.0 cm³) into a conical flask.
Add a few drops of a suitable indicator (e.g. Phenolphthalein or methyl orange).
Fill a burette with the acid and record the starting volume.
Add the acid dropwise to the alkali, swirling the flask continuously.
Stop adding acid when the indicator permanently changes colour (the end point).
Record the final burette reading and calculate the volume of acid used (the titre).
Repeat until you get concordant results (titres within 0.10 cm³ of each other).

Key points for exam questions

Use a white tile under the flask to see colour changes more clearly.
The titre is: final burette reading minus initial burette reading.
Take the mean of concordant results only (ignore anomalous values).
You must read the burette at the bottom of the meniscus.

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Titration calculations use moles. Need a refresher? Read The Mole Explained: The One Concept That Unlocks All of GCSE Chemistry.

4.4.3 Electrolysis

Electrolysis uses electricity to decompose an ionic compound. For electrolysis to work, the ions must be free to move. This means the compound must be either molten or dissolved in water (aqueous).

In the solid state, the strong electrostatic forces in the ionic lattice hold the ions in fixed positions, preventing them from carrying charge. When melted or dissolved, the lattice breaks down and the ions become mobile.

Key terminology

Electrolyte: the ionic compound, either molten or in aqueous solution.
Electrode: a solid conductor through which the current enters and leaves the electrolyte. Usually made of inert materials like graphite or platinum.
Cathode (negative electrode): attracts positive ions (cations). Reduction occurs here (cations gain electrons).
Anode (positive electrode): attracts negative ions (anions). Oxidation occurs here (anions lose electrons).

Use the mnemonic PANIC: Positive Anode, Negative Is Cathode. Or remember AN OX, RED CAT: Anode = Oxidation, Reduction = Cathode.

A simple electrolysis setup. Positive cations migrate to the negative cathode, while negative anions migrate to the positive anode.

Electrolysis of Molten Compounds

For a simple molten binary ionic compound (one containing only two elements), the products are straightforward:

The metal is always produced at the cathode.
The non-metal is always produced at the anode.

Example: lead(II) bromide

PbBr₂(l) → Pb(l) + Br₂(g)

At the cathode: Pb²⁺ ions gain electrons and form molten lead.

Pb²⁺(l) + 2e⁻ → Pb(l)

At the anode: Br⁻ ions lose electrons and form bromine gas (brown vapour).

2Br⁻(l) → Br₂(g) + 2e⁻

When writing half-equations for electrolysis, check that the electrons are on the correct side. Cathode = electrons on the left (gained). Anode = electrons on the right (lost).

Extraction of Aluminium

Aluminium is too reactive to be extracted by reduction with carbon. It must be extracted by electrolysis of aluminium oxide (Al₂O₃), purified from bauxite ore.

Why is cryolite used?

Pure aluminium oxide has an extremely high melting point (over 2000°C). Melting it directly would be hugely expensive. Instead, the aluminium oxide is dissolved in molten cryolite (Na₃AlF₆), which lowers the operating temperature to about 950°C. This significantly reduces energy costs.

At the electrodes

Both electrodes are made from carbon (graphite).

Cathode: Al³⁺ + 3e⁻ → Al (molten aluminium sinks to the bottom and is tapped off).
Anode: 2O²⁻ → O₂ + 4e⁻ (oxygen gas is produced).

The carbon anodes must be replaced regularly. At the high operating temperatures, the oxygen produced reacts with the carbon anodes to form CO₂, causing them to burn away. This is a major ongoing cost of the process.

A common exam question: "Why must the carbon anodes be replaced?" Answer: because the oxygen produced at the anode reacts with the hot carbon, forming carbon dioxide, which gradually wears the anodes away.

Electrolysis of Aqueous Solutions Required Practical

When an ionic compound is dissolved in water, there is an added complication: water itself partially ionises, introducing extra H⁺ and OH⁻ ions into the solution. This means there are competing ions at each electrode.

Rules for the cathode (negative electrode)

Both the metal cations and H⁺ ions from water migrate to the cathode. Which one is discharged depends on reactivity:

If the metal is more reactive than hydrogen (e.g. Sodium, calcium, aluminium), hydrogen gas is produced at the cathode.
If the metal is less reactive than hydrogen (e.g. Copper, silver), the metal is deposited at the cathode.

Rules for the anode (positive electrode)

Both the non-metal anions and OH⁻ ions from water migrate to the anode:

If halide ions (Cl⁻, Br⁻, I⁻) are present, the halogen is produced (e.g. Chlorine gas, bromine).
If no halide ions are present, oxygen gas is produced (from the discharge of OH⁻ ions).

This is Required Practical 3. You need to be able to predict the products at each electrode for any given aqueous solution, and describe the observations (gas bubbles, metal plating, colour changes).

Predicting products: electrolysis of copper sulfate solution

Cathode: Cu²⁺ and H⁺ are present. Copper is less reactive than hydrogen, so copper metal is deposited. Observation: pink/brown solid coats the cathode.

Anode: SO₄²⁻ and OH⁻ are present. No halide ions are present, so oxygen gas is produced. Observation: bubbles at the anode.

Predicting products: electrolysis of sodium chloride solution

Cathode: Na⁺ and H⁺ are present. Sodium is more reactive than hydrogen, so hydrogen gas is produced. Observation: bubbles; squeaky pop with a lighted splint.

Anode: Cl⁻ and OH⁻ are present. Halide ions are present, so chlorine gas is produced. Observation: bubbles; bleaches damp litmus paper.

Predicting products: electrolysis of copper(II) bromide solution

Cathode: Cu²⁺ and H⁺ are present. Copper is less reactive than hydrogen, so copper metal is deposited. Observation: brown/pink solid coats the cathode.

Anode: Br⁻ and OH⁻ are present. Bromide is a halide ion, so bromine is produced. Observation: orange/brown colour near the anode.

Practice: write the half-equations for the electrolysis of copper(II) bromide solution

Cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)

Anode (oxidation): 2Br⁻(aq) → Br₂(aq) + 2e⁻

Check: 2 electrons lost at the anode = 2 electrons gained at the cathode. ✔

Electrolysis of Brine

Brine is a concentrated solution of sodium chloride (NaCl). Its electrolysis is an important industrial process because it produces three useful products:

Chlorine gas (Cl₂) at the anode - used in bleach, PVC plastics, and water purification/disinfection.
Hydrogen gas (H₂) at the cathode - used as a fuel and in the manufacture of margarine (hardening vegetable oils).
Sodium hydroxide solution (NaOH) left in the solution - used in soap, paper and ceramics manufacturing, and oven cleaners.

Half-equations

2Cl⁻(aq) → Cl₂(g) + 2e⁻ (anode)

2H⁺(aq) + 2e⁻ → H₂(g) (cathode)

Testing the products

Chlorine: bleaches damp litmus paper (turns it white).
Hydrogen: squeaky pop with a lighted splint.
Sodium hydroxide: turns universal indicator blue/purple (alkaline).

The sodium ions (Na⁺) and hydroxide ions (OH⁻) remain in solution because sodium is too reactive to be discharged. They combine to form NaOH, which is why the remaining solution is alkaline.

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