Electrochemistry Explained: Matric Physical Sciences Paper 2
Master electrochemistry for Matric Physical Sciences Paper 2 — galvanic cells, electrolytic cells, standard electrode potentials, EMF calculations, and common NSC question patterns.
By Tania Galant in Subject Guides · 8 min read
Key Takeaways
Electrochemistry is worth approximately 20-28 marks in Paper 2 and combines theory with calculations
Galvanic cells convert chemical energy to electrical energy while electrolytic cells do the reverse
The Table of Standard Electrode Potentials (Table 4B) is provided in the exam and is essential for EMF calculations
Understanding the differences between galvanic and electrolytic cells is one of the most frequently tested concepts
# Electrochemistry Explained: Matric Physical Sciences Paper 2
Electrochemistry connects chemistry to electricity and is one of the most important topics in Matric Physical Sciences Paper 2. Worth approximately 20-28 marks, it tests your understanding of galvanic cells, electrolytic cells, and the use of the Table of Standard Electrode Potentials (Table 4B). The good news is that electrochemistry follows very clear patterns, and once you understand the principles, the questions become predictable.
This guide covers everything you need to know: how galvanic and electrolytic cells work, how to calculate EMF, how to use Table 4B, and the types of questions the NSC examiners ask. For your full Physical Sciences study plan, see our [physical sciences guide](/blog/matric-physical-sciences-past-papers-and-exam-guide-your-complete-study-companion).
## Galvanic (Voltaic) Cells
> **Read more:** For a comprehensive overview, see our [physical sciences exam guide](/blog/matric-physical-sciences-past-papers--exam-guide).
A galvanic cell converts **chemical energy into electrical energy**. It produces electricity through a spontaneous redox reaction.
### Components of a Galvanic Cell
| Component | Function |
|---|---|
| Two half-cells | Each contains an electrode in an electrolyte solution |
| Anode | The electrode where oxidation occurs (negative terminal) |
| Cathode | The electrode where reduction occurs (positive terminal) |
| Salt bridge | Completes the circuit by allowing ion flow; maintains electrical neutrality |
| External circuit | Wires connecting the two electrodes; electrons flow through here |
### How a Galvanic Cell Works
1. At the **anode**, the more reactive metal is oxidised (loses electrons).
- The metal atoms go into solution as ions.
- Electrons are released into the external circuit.
2. Electrons flow through the external circuit from anode to cathode.
3. At the **cathode**, the less reactive metal ions in solution are reduced (gain electrons).
- Metal ions are deposited on the cathode as solid metal.
4. The **salt bridge** allows ions to flow between the two half-cells to maintain charge balance.
- Negative ions move towards the anode half-cell.
- Positive ions move towards the cathode half-cell.
### Identifying the Anode and Cathode
Using Table 4B (the Table of Standard Electrode Potentials):
- The metal **higher up** in the table (more negative E° value) is the **anode** (oxidised).
- The metal **lower down** in the table (more positive E° value) is the **cathode** (reduced).
**Memory aid:** "An Ox" and "Red Cat" — Anode = Oxidation, Reduction = Cathode.
### Writing Half-Reactions
- **Anode (oxidation):** Write the half-reaction from Table 4B in reverse.
Example: Zn → Zn²⁺ + 2e⁻
- **Cathode (reduction):** Write the half-reaction exactly as it appears in Table 4B.
Example: Cu²⁺ + 2e⁻ → Cu
### Cell Notation
The standard way to represent a galvanic cell:
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
- Single line (|) represents a phase boundary.
- Double line (||) represents the salt bridge.
- Anode is written on the left, cathode on the right.
## EMF Calculations
The EMF (electromotive force) of a cell is the potential difference between the two half-cells.
### Formula
**EMF = E°(cathode) - E°(anode)**
Where E° values are the standard electrode potentials from Table 4B.
### Example
For a zinc-copper galvanic cell:
- E°(Zn²⁺/Zn) = -0.76 V (anode)
- E°(Cu²⁺/Cu) = +0.34 V (cathode)
EMF = (+0.34) - (-0.76) = +0.34 + 0.76 = **+1.10 V**
### Key Points About EMF
- A **positive EMF** means the reaction is spontaneous (galvanic cell will work).
- A **negative EMF** means the reaction is non-spontaneous (will not happen on its own — needs an external power source, i.e., electrolytic cell).
- The EMF does **not** depend on the number of moles or the coefficients in the balanced equation.
- EMF values are for standard conditions (1 mol/dm³ concentration, 25°C, 1 atm).
### Factors That Affect EMF
While standard EMF is calculated from Table 4B, the actual EMF can change:
- **Concentration:** Increasing the concentration of the cathode electrolyte increases EMF. Increasing the concentration of the anode electrolyte decreases EMF.
- **Temperature:** Changes in temperature affect EMF (but you do not need to calculate this).
## Electrolytic Cells
An electrolytic cell converts **electrical energy into chemical energy**. It uses an external power source to drive a non-spontaneous reaction.
### Components of an Electrolytic Cell
| Component | Function |
|---|---|
| External power source (battery) | Provides the energy to drive the non-spontaneous reaction |
| Anode | Connected to the positive terminal of the battery; oxidation occurs here |
| Cathode | Connected to the negative terminal of the battery; reduction occurs here |
| Electrolyte | The solution or molten substance that contains the ions |
| Electrodes | May be inert (e.g., carbon, platinum) or active |
### How an Electrolytic Cell Works
1. The external power source pushes electrons from the anode to the cathode through the external circuit.
2. At the **anode** (positive electrode), anions are oxidised (lose electrons).
3. At the **cathode** (negative electrode), cations are reduced (gain electrons).
4. Ions in the electrolyte migrate to the appropriate electrode.
### Important Differences from Galvanic Cells
| Feature | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Energy conversion | Chemical → Electrical | Electrical → Chemical |
| Reaction type | Spontaneous | Non-spontaneous |
| EMF | Positive | Negative (needs external energy) |
| Anode charge | Negative | Positive |
| Cathode charge | Positive | Negative |
| Salt bridge | Yes (two separate solutions) | No (one electrolyte) |
| External power source | No | Yes |
| Electron flow | Anode → Cathode (through wire) | Anode → Cathode (through wire) |
**Critical note:** In BOTH types of cells, oxidation occurs at the anode and reduction occurs at the cathode. This never changes. What changes is the charge of the electrodes.
### Products at Electrodes in Electrolytic Cells
For the NSC exam, you need to know the products of electrolysis for common electrolytes:
**Electrolysis of concentrated NaCl (brine):**
- Cathode: Na⁺ is not reduced (water is reduced instead) → H₂ gas
- Anode: Cl⁻ is oxidised → Cl₂ gas
**Electrolysis of CuSO₄ with copper electrodes:**
- Cathode: Cu²⁺ + 2e⁻ → Cu (copper deposited)
- Anode: Cu → Cu²⁺ + 2e⁻ (copper dissolves)
- This is the basis of copper refining.
**Electrolysis of water (acidified):**
- Cathode: H⁺ ions reduced → H₂ gas
- Anode: OH⁻ ions oxidised → O₂ gas
## Using Table 4B Effectively
The Table of Standard Electrode Potentials (Table 4B) is provided in the exam. Here is how to use it:
### Reading the Table
- Half-reactions are written as reduction reactions (gaining electrons).
- The table is arranged from most negative E° (top) to most positive E° (bottom).
- Species at the **top left** are strong reducing agents (easily oxidised).
- Species at the **bottom right** are strong oxidising agents (easily reduced).
### Predicting Spontaneous Reactions
A reaction is spontaneous when:
- The reducing agent (species being oxidised) is **higher** in the table.
- The oxidising agent (species being reduced) is **lower** in the table.
**Think of it as a "Z" pattern:** The oxidation half-reaction is read from right to left at the top, and the reduction half-reaction is read from left to right at the bottom.
### Writing Balanced Equations
1. Write the oxidation half-reaction (reverse the Table 4B entry for the anode).
2. Write the reduction half-reaction (as it appears in Table 4B for the cathode).
3. Balance electrons: multiply half-reactions so the electrons cancel.
4. Add the two half-reactions.
5. Simplify (cancel species that appear on both sides).
### Example
Balanced equation for Zn-Cu cell:
Oxidation: Zn → Zn²⁺ + 2e⁻
Reduction: Cu²⁺ + 2e⁻ → Cu
Electrons are already balanced (2e⁻ each), so add:
**Zn + Cu²⁺ → Zn²⁺ + Cu**
## Common NSC Question Patterns
### Pattern 1: Identify the Anode and Cathode
Given a galvanic cell description or diagram, identify which electrode is the anode and which is the cathode. Use Table 4B — the more reactive metal (higher in the table) is the anode.
### Pattern 2: Calculate EMF
Use the formula: EMF = E°(cathode) - E°(anode). Always check that your answer is positive for a galvanic cell.
### Pattern 3: Write Half-Reactions and Balanced Equations
Write the oxidation and reduction half-reactions, balance electrons, and combine into a net cell reaction.
### Pattern 4: Compare Galvanic and Electrolytic Cells
Describe the differences in a table or paragraph format. This is a very common question.
### Pattern 5: Predict Products of Electrolysis
Given an electrolyte and electrode material, predict what is produced at each electrode.
### Pattern 6: Effect of Concentration Changes on EMF
Explain how increasing or decreasing the concentration of a particular electrolyte affects the cell's EMF.
## Practice Strategy
| Week | Focus | Activity |
|---|---|---|
| 1 | Galvanic cells — setup, half-reactions, EMF | Study theory, label diagrams, calculate EMF |
| 2 | Electrolytic cells — setup, products, applications | Study theory, compare with galvanic cells |
| 3 | Table 4B practice — predicting reactions, balancing | Work through multiple examples |
| 4 | Past paper questions | Full electrochemistry sections under timed conditions |
Download past papers from our [past papers page](/past-papers) and use our [past papers guide](/blog/the-complete-guide-to-matric-past-papers-everything-you-need-to-know) for effective practice.
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## Related Resources
- [Matric Physical Sciences Past Papers & Exam Guide: Your Complete Study Companion](/blog/matric-physical-sciences-past-papers-exam-guide-your-complete-study-companion)
- [Browse All Matric Past Papers](/past-papers)
- [Exam Preparation Guide](/exam-preparation)
- [Matric Mathematics Paper 1 vs Paper 2: Key Differences and How to Prepare for Each](/blog/matric-mathematics-paper-1-vs-paper-2-key-differences-and-how-to-prepare-for-each)
- [Euclidean Geometry Proofs: A Complete Guide for Matric Mathematics](/blog/euclidean-geometry-proofs-a-complete-guide-for-matric-mathematics)
- [Newton's Laws Made Simple: Matric Physical Sciences Paper 1 Guide](/blog/newtons-laws-made-simple-matric-physical-sciences-paper-1-guide)
- [Start Practising Free on LearningLoop](/auth?tab=register)
## Frequently Asked Questions
### How many marks is electrochemistry worth in Paper 2?
Electrochemistry is typically worth 20-28 marks in Paper 2.
### Which is the anode and which is the cathode?
The anode is where oxidation occurs. The cathode is where reduction occurs. In a galvanic cell, the anode is negative. In an electrolytic cell, the anode is positive. But oxidation always happens at the anode regardless of cell type.
### Do I need to memorise Table 4B?
No. Table 4B is provided in the exam paper. However, you must know how to read and use it correctly.
### What is the difference between EMF and potential difference?
EMF is the maximum potential difference a cell can provide (under standard conditions with no current flowing). Potential difference is the actual voltage when current flows, which is usually slightly less than the EMF due to internal resistance.
### Why does the salt bridge matter?
Without a salt bridge, the circuit is incomplete. The salt bridge allows ions to flow between the half-cells, maintaining electrical neutrality. Without it, charge would build up in each half-cell and the reaction would stop.
### How do I know which half-reaction to reverse?
The species being oxidised (the one higher in Table 4B) has its half-reaction reversed. The species being reduced (lower in the table) keeps its half-reaction as written in the table.
### Can EMF be negative?
For a galvanic cell, EMF should be positive (spontaneous reaction). If you calculate a negative EMF, it means the reaction is non-spontaneous and would require an electrolytic cell.
### What are the practical applications of electrolytic cells?
Electroplating (coating objects with a thin layer of metal), refining of metals (purifying copper), and the production of chemicals (chlorine gas from brine).
Explore more [Physical Sciences past papers](/subjects/physical-sciences) on our [subjects page](/subjects).