QCE Chemistry · Unit 3
QCE Chemistry Unit 3 Topic 2: Oxidation & Reduction — Flashcards & Quiz
QCE Chemistry Unit 3 Topic 2 examines electron transfer reactions and their applications in electrochemistry. These free flashcards and true/false questions cover oxidation and reduction definitions, oxidation states, balancing redox equations, the activity series, galvanic cells, standard electrode potentials, electrolytic cells, Faraday's laws, and industrial applications including aluminium smelting and electroplating. Every card is aligned to the QCAA Senior Chemistry syllabus.
Key Terms
- Oxidation
- The loss of electrons by a species, resulting in an increase in oxidation state. QCAA Chemistry Unit 3 Topic 2 EA questions require students to identify the oxidised species, write the oxidation half-equation and name the reducing agent in a given redox reaction.
- Reduction
- The gain of electrons by a species, resulting in a decrease in oxidation state. QCAA assessments pair reduction identification with the oxidising agent — the substance that accepts electrons and is itself reduced. Students must distinguish between the process and the agent.
- Galvanic cell
- An electrochemical cell that converts chemical energy into electrical energy through spontaneous redox reactions, with oxidation at the anode and reduction at the cathode. QCAA Chemistry Unit 3 Topic 2 EA diagram questions require labelling of electrodes, electron flow direction, ion migration and the salt bridge.
- Standard electrode potential (E)
- The voltage measured for a half-cell under standard conditions (1 M, 25 degrees C, 1 atm) relative to the standard hydrogen electrode. QCAA EA questions require students to use the electrochemical series to predict reaction spontaneity and calculate cell voltage using E(cell) = E(cathode) minus E(anode).
- Electrolytic cell
- An electrochemical cell that uses external electrical energy to drive a non-spontaneous redox reaction, with oxidation at the anode and reduction at the cathode. QCAA Unit 3 Topic 2 assessments contrast electrolytic cells with galvanic cells and may ask students to predict discharge products during electrolysis.
- Salt bridge
- A device connecting two half-cells in a galvanic cell, allowing ion migration to maintain electrical neutrality while preventing direct mixing of solutions. QCAA Chemistry EA questions test whether students understand that the salt bridge completes the circuit by balancing charge, not by transferring electrons.
Sample Flashcards
Q1: Define oxidation and reduction in terms of electron transfer.
Oxidation is the LOSS of electrons (OIL). Reduction is the GAIN of electrons (RIG). In every redox reaction, one species is oxidised and another is reduced. The total electrons lost must equal the total electrons gained. The species that causes oxidation is the oxidising agent (it is itself reduced).
Q2: Explain the rules for assigning oxidation states.
Rules: 1) Elements in their standard state = 0. 2) Monatomic ions = charge. 3) Oxygen = −2 (except peroxides = −1). 4) Hydrogen = +1 (except metal hydrides = −1). 5) Fluorine = −1 always. 6) Sum of oxidation states = overall charge. An increase in oxidation state = oxidation; a decrease = reduction.
Q3: Write and balance a redox half-equation.
Steps: 1) Write the unbalanced half-equation. 2) Balance atoms other than O and H. 3) Balance O by adding H₂O. 4) Balance H by adding H⁺. 5) Balance charge by adding electrons. For alkaline conditions, add OH⁻ to both sides to neutralise H⁺.
Q4: Describe the metal activity series and its use in predicting reactions.
The activity series ranks metals by their tendency to lose electrons (be oxidised). More active metals are stronger reducing agents. A more active metal will displace a less active metal from solution. The series from most to least active includes: K, Na, Ca, Mg, Al, Zn, Fe, Sn, Pb, H, Cu, Ag, Au.
Q5: Describe the structure and function of a galvanic (voltaic) cell.
A galvanic cell converts chemical energy to electrical energy through spontaneous redox reactions. Components: two half-cells (electrodes in electrolyte solutions), a salt bridge (maintains electrical neutrality), and an external circuit. The anode (oxidation, negative terminal) loses mass; the cathode (reduction, positive terminal) gains mass.
Q6: Explain the role of the salt bridge in a galvanic cell.
The salt bridge: 1) Completes the electrical circuit by allowing ion flow between half-cells. 2) Maintains electrical neutrality — cations migrate toward the cathode half-cell, anions migrate toward the anode half-cell. Without it, the reaction stops almost immediately as charge builds up.
Q7: Explain standard electrode potentials (E°) and how to calculate cell EMF.
E° values are measured relative to the Standard Hydrogen Electrode (SHE, E° = 0.00 V) under standard conditions (25°C, 1 mol/L, 1 atm). More positive E° = stronger oxidising agent. E°cell = E°cathode − E°anode. A positive E°cell means the reaction is spontaneous.
Q8: Describe the structure and function of an electrolytic cell.
An electrolytic cell uses electrical energy to drive a non-spontaneous redox reaction. Components: external power source, two electrodes in an electrolyte. Anode is positive (oxidation occurs), cathode is negative (reduction occurs). Unlike galvanic cells, the polarity is reversed.
Sample Quiz Questions
Q1: Oxidation involves the gain of electrons.
Answer: FALSE
Oxidation is the LOSS of electrons (OIL RIG). Reduction is the gain of electrons.
Q2: The oxidising agent is the species that is reduced in a redox reaction.
Answer: TRUE
The oxidising agent accepts electrons (causes oxidation) and is itself reduced.
Q3: The oxidation state of oxygen is always −2 in all compounds.
Answer: FALSE
Oxygen is −2 in most compounds but is −1 in peroxides (e.g. H₂O₂) and 0 in elemental form (O₂).
Q4: A more active metal can displace a less active metal from solution.
Answer: TRUE
The more active metal is more easily oxidised and donates electrons to reduce the less active metal's ions.
Q5: Copper can displace zinc from a zinc sulfate solution.
Answer: FALSE
Copper is LESS active than zinc. Only a MORE active metal can displace a less active one.
Why It Matters
Oxidation and reduction is a topic where conceptual understanding and practical application converge in QCE Chemistry Unit 3. The QCAA external exam tests your ability to assign oxidation states, balance redox half-equations, predict cell voltages using standard electrode potentials and explain the operation of galvanic and electrolytic cells. Many students find this topic challenging because it requires simultaneous tracking of electron transfer, ion movement and electrode processes. Mastering redox gives you a significant advantage in the exam and connects directly to the student experiment, where electrochemistry practicals are commonly assessed. This topic builds on the equilibrium concepts from Topic 1, as galvanic and electrolytic cells involve competing forward and reverse reactions governed by electrode potentials. QCAA exam questions frequently require you to use a standard electrode potential table to predict whether a reaction is spontaneous, calculate cell voltage, and identify which electrode is the anode and which is the cathode.
Key Concepts
Oxidation States and Redox Identification
Assign oxidation states systematically using the standard rules. Identify which species is oxidised (loses electrons, oxidation state increases) and which is reduced (gains electrons, oxidation state decreases). Practise identifying the oxidising agent and reducing agent in complex reactions — QCAA frequently tests this distinction.
Half-Equations and Balancing Redox Reactions
Use the half-equation method to balance redox reactions in acidic and basic solutions. Write separate oxidation and reduction half-equations, balance atoms and charges, then combine. QCAA exams frequently present unbalanced equations and expect you to apply this systematic approach.
Galvanic Cells and Standard Electrode Potentials
Explain how galvanic cells convert chemical energy to electrical energy. Draw and label a cell showing anode (oxidation), cathode (reduction), salt bridge, electron flow and ion movement. Use the electrochemical series to predict which metal is the anode and calculate standard cell potential using E = E(cathode) - E(anode).
Electrolytic Cells and Industrial Applications
Contrast electrolytic cells (electrical energy drives non-spontaneous reactions) with galvanic cells. Understand electrolysis of molten salts and aqueous solutions, predicting which ions discharge at each electrode. Apply this knowledge to industrial processes like aluminium smelting (Hall-Heroult) and electroplating.
Common Mistakes to Avoid
- Confusing the oxidising agent with the oxidised species — the oxidising agent is the substance that GAINS electrons (is reduced), while the oxidised species LOSES electrons. QCAA Chemistry Unit 3 Topic 2 EA marking rubrics specifically test this distinction.
- Drawing electron flow through the salt bridge instead of through the external wire — electrons flow through the external circuit from anode to cathode, while ions migrate through the salt bridge to balance charge. QCAA diagram questions deduct marks for this error.
- Calculating cell voltage by adding half-cell potentials instead of subtracting — the correct formula is E(cell) = E(cathode) minus E(anode). QCAA EA calculation questions require consistent use of this convention with the electrochemical series.
- Forgetting to reverse the sign of the anode half-equation when combining half-equations — when writing the overall equation, the oxidation half-equation must be reversed from the electrochemical series table. QCAA assessments test systematic half-equation balancing.
- Stating that the anode is always positive — in a galvanic cell the anode is negative, but in an electrolytic cell the anode is positive. QCAA Chemistry EA questions on electrochemistry require students to specify the cell type when discussing electrode polarity.
Study Tips
- Practise assigning oxidation states to every atom in ten different compounds and ions until the rules become automatic.
- Draw and label both galvanic and electrolytic cells from memory, including electron flow direction, ion migration and electrode reactions.
- Use the electrochemical series to predict five different cell voltages, writing the half-equations and calculating EMF each time.
- Balance three redox equations using the half-equation method in acidic solution, then repeat in basic solution to build confidence with both approaches.
- Use flashcards with spaced repetition to drill oxidation state rules, electrochemical series rankings and cell diagram conventions — instant recall of terms like oxidising agent, reducing agent, anode and cathode is essential for exam success.
- Before your exam, work through the practice questions in this set at least twice using spaced repetition. Testing yourself repeatedly is the most effective revision strategy for long-term retention.
Related Topics
Frequently Asked Questions
What does QCE Chemistry Unit 3 Topic 2 cover?
Unit 3 Topic 2 covers oxidation and reduction, oxidation states, balancing redox half-equations, the activity series, galvanic cells, electrode potentials, electrolytic cells and industrial electrochemistry.
How many flashcards are in this set?
This free set contains 20 flashcards and 20 true/false quiz questions aligned to the QCAA Senior Chemistry syllabus.
Are these aligned to the QCE syllabus?
Yes — every card maps to QCAA syllabus objectives for QCE Chemistry Unit 3 Topic 2: Oxidation and Reduction.
Last updated: March 2026 · 20 flashcards · 20 quiz questions · Content aligned to the QCAA Syllabus