Electrochemistry is one of the most misunderstood topics in Grade 12 Physical Science, and that is a problem — because the NSC exam tests it every single year in Paper 2, typically for 10 to 14 marks. Students lose marks not because the calculations are hard, but because they get confused about which electrode is the anode, which is the cathode, and which direction the electrons flow. Get those three things right and the marks follow. This post makes all of that clear.
✅ In this post you will learn:
- What a galvanic cell is and how to interpret a standard cell diagram
- What an electrolytic cell is and how it differs from a galvanic cell
- How to use the Standard Reduction Potential table to calculate cell voltage
- How to identify the anode and cathode in both cell types
- How to write half-reactions and the overall cell reaction
- How electrochemistry questions are structured in the NSC exam
⚡ The Big Picture: Two Types of Electrochemical Cells
Before you can answer exam questions, you need a firm mental model.
A galvanic cell converts chemical energy into electrical energy. The reaction happens spontaneously. A battery is a galvanic cell. The reaction wants to happen — you just connect the circuit and let it run.
An electrolytic cell uses electrical energy to force a chemical reaction that would not happen on its own. Electroplating and the production of aluminium from bauxite are examples. You have to push electricity into it.
Both cells have an anode and a cathode. Both involve oxidation and reduction. But the rules for which electrode does what are opposite between the two — and this is what trips students up every year.
📌 THE KEY RULE TO MEMORISE
OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)
In BOTH cell types: Oxidation always occurs at the ANODE. Reduction always occurs at the CATHODE.
The difference: In a galvanic cell the anode is NEGATIVE. In an electrolytic cell the anode is POSITIVE.
Galvanic Cells in Detail
In a galvanic cell, two different metals are placed in electrolyte solutions and connected by a wire (for electrons) and a salt bridge (for ions).
The more reactive metal oxidises — it loses electrons and dissolves into solution. This is the anode (negative electrode).
The less reactive metal gains electrons — metal ions from solution are deposited onto it. This is the cathode (positive electrode).
Example — The Zinc-Copper Cell (Daniell Cell):
Zn/Zn²⁺ is one half-cell. Cu/Cu²⁺ is the other.
From the Standard Reduction Potential table:
- Zn²⁺ + 2e⁻ → Zn : E° = −0.76 V
- Cu²⁺ + 2e⁻ → Cu : E° = +0.34 V
Zinc has the lower (more negative) reduction potential, so it oxidises. It is the anode.
Half-reactions: Anode (oxidation): Zn → Zn²⁺ + 2e⁻ Cathode (reduction): Cu²⁺ + 2e⁻ → Cu
Overall: Zn + Cu²⁺ → Zn²⁺ + Cu
Cell voltage: E°cell = E°cathode − E°anode E°cell = 0.34 − (−0.76) = +1.10 V
A positive E°cell means the reaction is spontaneous. This is how galvanic cells work.
The Standard Reduction Potential Table
The CAPS exam supplies this table in the data sheet. You do not memorise the values — but you must know how to read and use it.
The table lists half-reactions as reductions (gaining electrons). The values range from very negative (strong reducing agents, like Li) to very positive (strong oxidising agents, like F₂).
The rule for using the table:
| To find... | Do this... |
|---|---|
| Which substance oxidises | Find the lower E° value — it oxidises |
| Which substance reduces | Find the higher E° value — it reduces |
| Cell voltage | E°cell = E°cathode − E°anode |
| Spontaneous or not | Positive E°cell = spontaneous |
Important: When you reverse a half-reaction (from reduction to oxidation), you reverse the sign of E°. Do NOT reverse the sign when calculating E°cell with the formula above — the formula already accounts for it.
Electrolytic Cells in Detail
In an electrolytic cell, an external power source (battery) drives the reaction. The negative terminal of the battery connects to the cathode. The positive terminal connects to the anode.
Cathode: reduction occurs here. Positive ions from solution are attracted to this negative electrode and gain electrons.
Anode: oxidation occurs here. Negative ions or the electrode material itself loses electrons.
Example — Electrolysis of Copper(II) Sulfate solution with copper electrodes:
Cathode: Cu²⁺ + 2e⁻ → Cu (copper deposits onto the cathode) Anode: Cu → Cu²⁺ + 2e⁻ (copper dissolves from the anode)
This is exactly how electroplating works. The object to be plated is the cathode. The plating metal is the anode.
Example — Electrolysis of Water:
Cathode: 4H⁺ + 4e⁻ → 2H₂ (hydrogen gas produced) Anode: 2H₂O → O₂ + 4H⁺ + 4e⁻ (oxygen gas produced)
Volume of H₂ produced is always double the volume of O₂. The DBE tests this ratio regularly.
🔋 Galvanic vs Electrolytic — Side by Side
| Feature | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Energy conversion | Chemical → Electrical | Electrical → Chemical |
| Spontaneous? | Yes | No |
| Anode charge | Negative | Positive |
| Cathode charge | Positive | Negative |
| Power source needed? | No | Yes |
| Example | Battery, fuel cell | Electroplating, electrolysis |
Print this table. Stick it on your wall. It summarises everything.
⚠️ Common Mistakes Students Make
1. Flipping the anode/cathode charge between cell types In a galvanic cell the anode is negative. In an electrolytic cell the anode is positive. Students who do not learn this distinction will get the electrode charges wrong every time.
2. Reversing the E°cell formula E°cell = E°cathode minus E°anode — not the other way around. Writing it backwards gives you a negative value for a spontaneous reaction and you lose all the marks that follow.
3. Not balancing electrons before writing the overall equation If the half-reactions transfer different numbers of electrons, you must multiply one or both half-reactions so the electron count matches before combining them. The DBE deducts marks for unbalanced overall equations.
4. Confusing the salt bridge's function The salt bridge does not carry electrons — the wire does that. The salt bridge maintains electrical neutrality in each half-cell by allowing ions to flow between the solutions. If you say electrons flow through the salt bridge, that is incorrect.
5. Writing reduction at the anode Oxidation always occurs at the anode. Always. In both cell types. If you find yourself writing a reduction half-reaction at the anode, you have made an error somewhere.
📋 How This Topic Appears in the NSC Exam
Electrochemistry appears in Paper 2 of the NSC Physical Science exam.
It typically carries 10 to 14 marks and appears as Question 7 or Question 8.
The DBE structures these questions predictably:
- Part (a): identify the anode and cathode, state what occurs at each (3–4 marks)
- Part (b): write the half-reactions and the overall cell equation (4 marks)
- Part (c): calculate the standard cell potential E°cell using the data sheet (2–3 marks)
- Part (d): a comparison or application question — e.g. explain the role of the salt bridge, or describe what happens if the electrolyte concentration changes (2–3 marks)
In the 2023 NSC exam, Question 8 tested a galvanic cell based on iron and silver half-cells. Students were asked to write half-reactions, calculate E°cell, and explain the function of the salt bridge. The 2022 exam included an electrolytic cell question involving the electrolysis of brine, which is a classic CAPS application.
The pattern is consistent. If you can handle galvanic and electrolytic cells confidently, write balanced half-reactions, and use the E° table correctly, you will score well on this section.
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