AQA GCSE Mastery: Chemical Changes (Topic 4) – Reactivity, Acids, and Electrolysis (Dubai Guide)
Chemical change is the engine of industry and the essence of chemistry. Here in the UAE, we see this on a massive scale. Consider the transformation of raw materials into the high-grade aluminum produced at the vast smelters in Jebel Ali and Al Taweelah—a process requiring immense electrical energy. Contrast this high-intensity industrial process with the slow, relentless chemical changes occurring on the coast, where the humid, salty air corrodes unprotected metal infrastructure.
These are both examples of chemical changes, driven by the fundamental principles of reactivity and electron transfer.
For students in Dubai studying the AQA GCSE Combined Science: Trilogy (8464) specification, Topic 4 (Chemical Changes) is a cornerstone of Chemistry Paper 1. It introduces the study of reactivity, the behavior of acids and bases, and the powerful process of electrolysis.
This guide provides a definitive, syllabus-aligned resource for mastering these concepts, including the required practicals, and connects the theory directly to the world around us in the UAE.
Part 1: The Reactivity of Metals and Redox
The Reactivity Series
The reactivity series is a list of metals arranged in order of their reactivity, from highest to lowest. You must memorize this order, including the non-metals Carbon (C) and Hydrogen (H), as they serve as crucial benchmarks.
Potassium (K) – Most Reactive
Sodium (Na)
Lithium (Li)
Calcium (Ca)
Magnesium (Mg)
(Carbon)
Zinc (Zn)
Iron (Fe)
(Hydrogen)
Copper (Cu)
Silver (Ag)
Gold (Au) – Least Reactive
The "Why" of Reactivity
The reactivity of a metal is determined by how easily its atoms lose their outer shell electrons to form positive ions. The easier it is to lose electrons, the more reactive the metal. Metals like Potassium (Group 1) readily lose their single outer electron, whereas metals like Copper require significantly more energy to form an ion.
Displacement Reactions
The reactivity series allows us to predict reactions. The fundamental principle of displacement reactions is:
A more reactive metal will displace a less reactive metal from its compound.
For example, if you place a piece of zinc (more reactive) into a solution of copper sulfate (less reactive), the zinc will dissolve, and copper metal will be produced.
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Oxidation and Reduction (Redox)
Displacement reactions are examples of Redox reactions, where electrons are transferred. We use the mnemonic OIL RIG to define the two key processes:
Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)
In the example above:
Zinc loses 2 electrons to become a Zn²⁺ ion (Oxidation).
The Copper ion (Cu²⁺) gains 2 electrons to become a Copper atom (Reduction)
Writing Ionic Equations (Higher Tier Only)
For Higher Tier, you must be able to represent redox reactions using ionic equations, which only show the species that change during the reaction. Spectator ions (ions that remain unchanged) are omitted.
Let’s look at the reaction: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Separate the aqueous ionic compounds into their ions: Zn(s) + Cu²⁺(aq) + SO₄²⁻(aq) → Zn²⁺(aq) + SO₄²⁻(aq) + Cu(s)
Cancel the spectator ions (SO₄²⁻ appears on both sides).
The resulting ionic equation is: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Extraction of Metals
Most metals are found naturally in the Earth’s crust as compounds (ores). The extraction method depends on the metal’s reactivity.
Metals below Carbon: Extracted by heating the ore with carbon (e.g., Iron in a blast furnace). Carbon displaces the metal. This is called reduction by carbon.
Metals above Carbon: Are too reactive to be extracted by carbon. They must be extracted using a more powerful method: electrolysis.
Dubai Context 1: Corrosion and the Coastal Environment Corrosion (rusting, in the case of iron) is a redox reaction where the metal is oxidized by oxygen in the presence of water. Dubai’s coastal environment, characterized by high humidity and salinity (salt in the air), significantly accelerates this corrosion. Protecting the infrastructure of our city requires methods like galvanizing (coating iron with a layer of more reactive zinc) or sacrificial protection, where a more reactive metal is attached and corrodes (is oxidized) in place of the structural metal.
Part 2: Acids, Bases, and Neutralization
The pH Scale and Definitions
In aqueous solutions (dissolved in water):
Acids produce Hydrogen ions (H⁺).
Alkalis (soluble bases) produce Hydroxide ions (OH⁻).
The pH scale (ranging from 0 to 14) is a measure of the concentration of H⁺ ions.
pH < 7 is Acidic
pH = 7 is Neutral
pH > 7 is Alkaline
(Higher Tier Only): The pH scale is logarithmic. As the pH decreases by one unit, the concentration of H⁺ ions increases by a factor of 10.
Strong vs. Weak Acids (Higher Tier Only)
These consist of small, discrete molecules (e.g., H₂O, O₂, CH₄).
The Crucial Distinctioa
It is crucial to understand the difference between the strength of an acid and its concentration. Concentration refers to how much acid is dissolved in the water. Strength refers to how the acid behaves when dissolved.
Strong Acids (e.g., Hydrochloric, Sulfuric, Nitric) completely ionize (or dissociate) in water. Every molecule releases its H⁺ ion.
Weak Acids (e.g., Ethanoic, Citric, Carbonic) only partially ionize. Only a small fraction of the molecules release their H⁺ ions.
Reactions of Acids (The "Big Three")
Acids react in predictable patterns. You must learn these general equations:
1. Acid + Metal → Salt + Hydrogen
Example: Magnesium + Hydrochloric Acid → Magnesium Chloride + Hydrogen Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) (Note: This is also a redox reaction, provided the metal is above hydrogen in the reactivity series.)
2. Acid + Base → Salt + Water
Bases include metal oxides and metal hydroxides. This is a neutralization reaction. Example: Sulfuric Acid + Copper Oxide → Copper Sulfate + Water H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)
3. Acid + Metal Carbonate → Salt + Water + Carbon Dioxide
Example: Nitric Acid + Calcium Carbonate → Calcium Nitrate + Water + Carbon Dioxide 2HNO₃(aq) + CaCO₃(s) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)
The Ionic Equation for Neutralization
In any neutralization reaction between an acid and an alkali, the essential chemical change is:
H⁺(aq) + OH⁻(aq) → H₂O(l)
Skills Focus: Naming Salts
The name of the salt produced depends on the metal reactant and the acid used:
Hydrochloric Acid produces Chlorides.
Sulfuric Acid produces Sulfates.
Nitric Acid produces Nitrates.
Dubai Context 2: pH Adjustment in Water Treatment Dubai relies heavily on desalination. The water produced immediately after the reverse osmosis stage is often slightly acidic and lacks essential minerals. DEWA (Dubai Electricity and Water Authority) must carefully treat this water before distribution. This involves using neutralization processes, often adding alkaline substances, to adjust the pH to a safe, neutral level, protecting both the consumer and the water distribution network.
Required Practical 1: Preparation of a Soluble Salt
You need to know how to prepare a pure, dry sample of a soluble salt from an insoluble reactant (e.g., a metal oxide).
Methodology (Example: Preparing Copper Sulfate from Copper Oxide and Sulfuric Acid):
Reaction: Gently warm the dilute sulfuric acid. Add the copper oxide (the base) in excess (until no more dissolves). This ensures all the acid has reacted.
Filtration: Filter the mixture. This removes the unreacted (excess) copper oxide. The filtrate is the copper sulfate solution.
Crystallization: Gently heat the solution to evaporate some of the water, making the solution saturated.
Cooling: Leave the solution to cool. Crystals of copper sulfate will form. Filter the crystals and dry them.
Part 3: Electrolysis (The Major Topic)
Electrolysis is a sophisticated process with significant industrial applications.
The Fundamentals of Electrolysis
Electrolysis is the decomposition (breaking down) of an ionic compound using electrical energy.
Key Terminology
Electrolyte: The ionic compound being broken down. It must be molten (liquid) or dissolved (aqueous) so the ions are free to move and carry the charge.
Electrodes: Conductors (usually carbon/graphite or inert metals) that connect the electrolyte to the power supply.
Anode: The positive electrode. Negative ions (anions) are attracted here. Oxidation occurs at the anode. (Mnemonic: AN OX).
Cathode: The negative electrode. Positive ions (cations) are attracted here. Reduction occurs at the cathode. (Mnemonic: RED CAT).
Electrolysis of Molten Ionic Compounds
This is the simplest form of electrolysis because the electrolyte only contains two types of ions.
Example: Electrolysis of Molten Lead(II) Bromide (PbBr₂)
Ions present: Pb²⁺ and Br⁻.
At the Cathode (-): Pb²⁺ ions are attracted. They gain electrons (reduction) and form molten lead metal.
At the Anode (+): Br⁻ ions are attracted. They lose electrons (oxidation) and form bromine gas (Br₂).
Writing Half Equations (Higher Tier Only)
Half equations show the reactions occurring at each electrode, including the movement of electrons (represented as e⁻).
Using the Lead(II) Bromide example:
Cathode (Reduction): Pb²⁺ + 2e⁻ → Pb
Anode (Oxidation): 2Br⁻ → Br₂ + 2e⁻
Mastering the construction of balanced half-equations requires practice and a clear understanding of ionic theory and electron transfer. If you find these concepts challenging, personalized instruction from our expert AQA chemistry tutors in Dubai can provide the strategies needed to secure top marks in this complex area.
Industrial Application: Extraction of Aluminium
Aluminium is highly reactive (above carbon), so it must be extracted via electrolysis. This is an energy-intensive process.
The Process:
The Electrolyte: Aluminium oxide (Al₂O₃, purified from bauxite ore) has a very high melting point (over 2000°C). To save energy, it is dissolved in molten cryolite. Cryolite lowers the operating temperature to around 950°C.
The Reactions:
At the Cathode (-): Aluminium ions (Al³⁺) are reduced to form molten aluminium metal. Al³⁺ + 3e⁻ → Al
At the Anode (+): Oxide ions (O²⁻) are oxidized to form oxygen gas. 2O²⁻ → O₂ + 4e⁻
The Complication: The electrodes are made of carbon (graphite). Because the temperature is so high, the oxygen produced at the anode immediately reacts with the carbon electrodes, forming carbon dioxide (C + O₂ → CO₂). This means the anodes burn away and must be replaced regularly.
Dubai Context 3 (Crucial): Emirates Global Aluminium (EGA) This process is the foundation of one of the UAE’s largest heavy industries. Emirates Global Aluminium (EGA), with massive smelters in Jebel Ali (Dubai) and Al Taweelah (Abu Dhabi), is one of the world’s largest producers of premium aluminum. The entire operation relies on the large-scale, continuous electrolysis of aluminum oxide, consuming vast amounts of electricity to produce this vital metal.
Electrolysis of Aqueous Solutions (The Rules)
Electrolysis of aqueous solutions (dissolved compounds) is more complex because water itself dissociates slightly, producing H⁺ and OH⁻ ions. This means there is competition at the electrodes.
You must learn the rules to predict the products:
Rule at the Cathode (-)
The less reactive element is produced.
If the metal is more reactive than Hydrogen (e.g., Na, Mg, Al), Hydrogen gas is produced.
If the metal is less reactive than Hydrogen (e.g., Cu, Ag, Au), the Metal is produced.
Rule at the Anode (+)
If halide ions (Cl⁻, Br⁻, I⁻) are present, the Halogen gas is produced (e.g., Chlorine gas).
If no halide ions are present (e.g., Sulfate SO₄²⁻, Nitrate NO₃⁻), Oxygen gas is produced.
Worked Example: Electrolysis of Sodium Chloride Solution (Brine)
Ions Present: Na⁺, Cl⁻ (from salt), H⁺, OH⁻ (from water).
Cathode (-): Competition between Na⁺ and H⁺. Sodium is more reactive than hydrogen. Hydrogen gas is produced.
Anode (+): Competition between Cl⁻ and OH⁻. Chloride is a halide ion. Chlorine gas is produced.
Left in Solution: The remaining ions (Na⁺ and OH⁻) form Sodium Hydroxide (NaOH).
Dubai Context 4: Electroplating Electroplating uses electrolysis to coat an object with a thin layer of a desired metal. The object to be coated is used as the cathode. This technique is widely used in Dubai, from adding protective and decorative finishes to customized luxury cars to the gold-plating and silver-plating of intricate jewelry found in the Gold Souk.
Required Practical 3: Electrolysis of Aqueous Solutions
You must be able to investigate what happens when aqueous solutions are electrolyzed using inert electrodes.
Setup: A simple electrolytic cell, often collecting gases in inverted test tubes above the electrodes.
Observations: Identify the products using standard chemical tests:
Hydrogen (H₂): Burns with a ‘squeaky pop’ when a lit splint is applied.
Oxygen (O₂): Relights a glowing splint.
Chlorine (Cl₂): Bleaches damp litmus paper white.
Conclusion: Mastering Chemical Transformation
Topic 4 covers the fundamental ways in which substances are transformed. The key themes are interconnected: the reactivity of a metal determines how it reacts with acids and how it is extracted; neutralization governs the behavior of acids and bases; and electrolysis provides a powerful tool for decomposition and extraction.
Underpinning almost all of these processes is the concept of Redox—the transfer of electrons. Whether it’s a metal corroding in the Dubai humidity or being formed at the cathode during electrolysis, understanding where the electrons are going is the key to mastering chemical change and achieving success in your AQA GCSE exams