Electrochemistry links chemistry and electricity. Spontaneous redox reactions in galvanic cells release electrical energy, while electrolytic cells use electrical energy to drive non-spontaneous reactions. The standard electrode potential table allows us to predict which reactions will occur and to calculate cell voltages.
12.1 Galvanic Cells
Definition
Galvanic cell
An electrochemical cell in which a spontaneous redox reaction produces electrical energy. The anode (negative electrode) undergoes oxidation; the cathode (positive electrode) undergoes reduction.
In a galvanic (voltaic) cell, two different metals are placed in separate electrolyte solutions connected by a salt bridge. The more reactive metal (e.g. zinc) loses electrons at the anode (oxidation). Electrons flow through the external circuit to the cathode (e.g. copper), where reduction occurs. The salt bridge maintains electrical neutrality by allowing ion movement between the two half-cells.
Zn-Cu galvanic cell half-reactions
- Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻
- Cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)
- Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Exam Tip
Memory aid: OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). In a galvanic cell: ANODE = oxidation = negative; CATHODE = reduction = positive. In an electrolytic cell the signs flip — see Section 12.2.
Worked Example
A galvanic cell uses a magnesium anode and a silver cathode. Write (a) the half-reactions and (b) the overall cell equation.
Given
- Anode metal: Mg
- Cathode: Ag⁺/Ag
Find
Half-reactions and overall equation
Solution
- 1Anode (oxidation): Mg(s) → Mg²⁺(aq) + 2e⁻
- 2Cathode (reduction): Ag⁺(aq) + e⁻ → Ag(s) — balance electrons: multiply by 2
- 3Balanced cathode: 2Ag⁺(aq) + 2e⁻ → 2Ag(s)
- 4Overall: Mg(s) + 2Ag⁺(aq) → Mg²⁺(aq) + 2Ag(s)
Practice Question
A galvanic cell is constructed with a zinc electrode in ZnSO₄ and an iron electrode in FeSO₄. Using the EMF series (E°Zn²⁺/Zn = −0.76 V; E°Fe²⁺/Fe = −0.44 V), identify which electrode is the anode, which is the cathode, and write the half-reactions.
(6 marks)
Watch Out
Common mistake: writing the anode half-reaction as a reduction. The anode ALWAYS undergoes oxidation (electrons are produced). The cathode ALWAYS undergoes reduction (electrons are consumed).
12.2 Electrolytic Cells and Standard Electrode Potentials
Definition
Electrolytic cell
An electrochemical cell in which an external electrical source drives a non-spontaneous redox reaction.
Definition
Standard electrode potential
The potential difference measured under standard conditions (1 mol·dm⁻³, 25°C, 100 kPa) relative to the standard hydrogen electrode.
In an electrolytic cell, the external power supply forces current in a direction opposite to that which would occur spontaneously. The electrode connected to the positive terminal of the supply is the anode (oxidation occurs). The electrode connected to the negative terminal is the cathode (reduction occurs). Note: in electrolytic cells the anode is positive and the cathode is negative — the opposite of a galvanic cell.
Galvanic Cell vs Electrolytic Cell
| Property | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Energy conversion | Chemical → electrical | Electrical → chemical |
| Reaction type | Spontaneous redox | Non-spontaneous redox |
| Anode sign | Negative (−) | Positive (+) |
| Cathode sign | Positive (+) | Negative (−) |
| External supply | Not required — cell produces voltage | Required — external battery drives reaction |
| E°cell | Positive (spontaneous) | Negative (non-spontaneous) |
| Example | Zn-Cu Daniell cell, car battery | Electroplating, water electrolysis |
Formula
Standard cell emf
E°cell = standard cell emf (V), E°cathode = standard reduction potential of cathode (V), E°anode = standard reduction potential of anode (V)
SI unit: V
Exam Tip
Using the EMF table: the electrode with the higher (more positive) standard reduction potential is the cathode. The electrode with the lower (more negative) standard reduction potential is the anode. A positive E°cell means the reaction is spontaneous (galvanic). A negative E°cell means it is non-spontaneous (electrolytic).
Worked Example
Calculate the standard cell emf for a galvanic cell with a zinc anode and a copper cathode. (E°Zn²⁺/Zn = −0.76 V; E°Cu²⁺/Cu = +0.34 V). Is the reaction spontaneous?
Given
- E°cathode (Cu²⁺/Cu) = +0.34 V
- E°anode (Zn²⁺/Zn) = −0.76 V
Find
E°cell
Solution
- 1E°cell = E°cathode − E°anode = (+0.34) − (−0.76) = 0.34 + 0.76 = 1.10 V
- 2E°cell = +1.10 V (positive → spontaneous reaction ✓)
Practice Question
Using the standard electrode potentials: E°(Pb²⁺/Pb) = −0.13 V and E°(Sn²⁺/Sn) = −0.14 V. (a) Identify which is the anode and which is the cathode in a galvanic cell. (b) Calculate E°cell. (c) Write the overall cell equation.
(6 marks)
Watch Out
Common mistake: reversing the formula as E°cell = E°anode − E°cathode. Always use E°cell = E°cathode − E°anode. The cathode value comes first. If you get a negative answer for what should be a galvanic cell, you have the electrodes the wrong way around.
12.3 Practical Applications: Electroplating and Rechargeable Cells
Electroplating is a practical application of electrolysis. An electrolytic cell is used to deposit a thin layer of one metal onto another. The object to be plated is made the cathode (reduction occurs, metal deposits). The plating metal (e.g. copper, silver, gold) is made the anode — it dissolves to replenish the electrolyte. The electrolyte contains ions of the plating metal.
Electroplating setup (copper plating example)
- Cathode: the object to be plated (e.g. iron spoon) — connected to − terminal
- Anode: pure copper plate — connected to + terminal, dissolves slowly
- Electrolyte: copper sulfate solution (CuSO₄)
- Cathode half-reaction: Cu²⁺(aq) + 2e⁻ → Cu(s) — copper deposits on object
- Anode half-reaction: Cu(s) → Cu²⁺(aq) + 2e⁻ — copper anode dissolves
Real World
Electroplating applications: chrome plating on car bumpers (corrosion resistance), gold plating on jewellery and electronics (conductivity and aesthetics), silver plating on cutlery. Galvanising (coating steel with zinc) uses a similar process to protect steel from rusting.
Rechargeable cells (secondary cells) can be recharged by reversing the current. The lead-acid battery in a car is a classic example. During discharge (galvanic mode), lead is oxidised at the anode and lead dioxide is reduced at the cathode. During charging (electrolytic mode), the reactions are reversed. Lithium-ion batteries in phones and laptops work on the same principle — intercalation of lithium ions.
Primary vs Secondary (Rechargeable) Cells
| Property | Primary Cell | Secondary Cell |
|---|---|---|
| Rechargeability | Cannot be recharged — discarded when flat | Can be recharged many times |
| Mode | Galvanic only | Galvanic (discharge) and electrolytic (charge) |
| Examples | Alkaline AA battery, zinc-carbon cell | Lead-acid car battery, Li-ion phone battery |
| Cost | Lower initial cost | Higher initial cost but cheaper long-term |
| Environmental impact | More waste (disposable) | Less waste (reusable) |
Worked Example
During the charging of a lead-acid battery the following reaction occurs at the cathode: PbSO₄(s) + 2e⁻ → Pb(s) + SO₄²⁻(aq). (a) Is this cell now acting as a galvanic or electrolytic cell? (b) Is the cathode connected to the positive or negative terminal of the external charger?
Given
- Cathode reaction (charging): PbSO₄ + 2e⁻ → Pb + SO₄²⁻
Find
Cell type and terminal polarity
Solution
- 1During charging, an external electrical source drives the reaction → electrolytic cell.
- 2Reduction occurs at the cathode in an electrolytic cell.
- 3In an electrolytic cell, the cathode is connected to the negative (−) terminal of the external supply.
Practice Question
Describe the process of silver electroplating of a copper spoon. In your answer identify the anode, cathode, and electrolyte, and write the cathode half-reaction.
(5 marks)
Practice Question
A galvanic cell produces a current of 0.5 A for 2 hours. Calculate the charge transferred. (Q = It)
(3 marks)
Exam Tip
Exam tip: If a question asks you to compare a galvanic cell and an electrolytic cell, always address these four points: (1) energy conversion direction, (2) spontaneity of reaction, (3) sign of anode, (4) sign of cathode. Four points = four marks.
Watch Out
Do not confuse the salt bridge in a galvanic cell with the electrolyte in an electrolytic cell. They serve different purposes: the salt bridge maintains charge balance between two separate half-cells; the electrolyte in an electrolytic cell is a single solution through which ions carry the current.