Electrochemistry: Redox and Cells

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your NECO SSCE Paper II or III.

Electrochemistry links electrical energy to redox (oxidation–reduction) reactions. Oxidation is loss of electrons (oxidation number rises); reduction is gain of electrons (oxidation number falls). Both happen together in any redox process. The anode is where oxidation occurs; the cathode is where reduction occurs — but electrode polarity flips between galvanic and electrolytic cells.

Must-know relationships:

E°cell = E°cathode − E°anode (V); a positive value means the reaction is spontaneous.
ΔG° = −nFE°cell, where n = moles of e⁻ transferred and F = 96 500 C mol⁻¹.
m = (I × t × M) / (n × F) — mass of metal deposited during electrolysis.

Quick exam pointers:

The standard hydrogen electrode (SHE) is the 0.00 V reference.
In a galvanic cell, the more reactive metal (more negative E°) is the anode (negative); the less reactive is the cathode (positive).
NECO typically tests balancing half-equations, identifying oxidising/reducing agents, and one Faraday calculation.

🟡 Standard — Regular Study (2d–2mo)

Standard content for students preparing weeks ahead of the NECO SSCE.

Assigning Oxidation Numbers

Oxidation numbers let you track electron movement without writing half-equations. The rules below cover 90% of NECO items.

Rule	Example
Free elements = 0	O₂, Fe, Cl₂ all = 0
Combined O = −2 (except peroxides = −1)	H₂O₂: O = −1
Combined H = +1 (except metal hydrides = −1)	NaH: H = −1
Combined F is always −1	HF, NaF
Sum in a neutral compound = 0; sum in an ion = ion charge	SO₄²⁻: S + 4(−2) = −2

Balancing Redox by the Ion–Electron Method

Write oxidation and reduction half-equations separately, balance atoms other than O and H, balance O with H₂O, balance H with H⁺ (acidic medium) or OH⁻ (alkaline), then balance charge with electrons. Multiply the half-equations so electrons cancel, then add them.

Tip: In alkaline medium, after balancing with H⁺, add the same number of OH⁻ to both sides to convert leftover H⁺ into water.

Cell Potential and Spontaneity

Standard electrode potentials are measured against the SHE (E° = 0.00 V). The cell potential tells you whether a redox reaction is feasible.

E°cell = E°cathode − E°anode
If E°cell > 0 → spontaneous (galvanic cell discharges)
If E°cell < 0 → non-spontaneous; requires electrolysis

For a Zn–Cu cell: E°cell = (+0.34) − (−0.76) = +1.10 V, so Zn is oxidised and Cu²⁺ is reduced.

Faraday’s Laws of Electrolysis

First law: mass deposited m ∝ Q, where Q = I × t.
Second law: m ∝ equivalent weight (M/n).

The working equation is m = (I × t × M) / (n × F), with F = 96 500 C mol⁻¹.

Symbol	Meaning	Unit
m	Mass deposited	g
I	Current	A
t	Time	s
M	Molar mass	g mol⁻¹
n	Electrons per ion	—
F	Faraday constant	C mol⁻¹

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for long-term mastery and adjacent-topic links.

Worked Example: Faraday Calculation

A current of 2.5 A passes through a solution of CuSO₄ for 1 hour 12 minutes (4 320 s). Find the mass of copper deposited (Cu = 64, n = 2, F = 96 500 C mol⁻¹).

m = (I × t × M) / (n × F) = (2.5 × 4 320 × 64) / (2 × 96 500) = 691 200 / 193 000 = 3.58 g.

A common trap is to omit n or to divide by M instead of multiply — always state the half-equation first (Cu²⁺ + 2e⁻ → Cu) so n is unambiguous.

Electrode Polarity: Galvanic vs Electrolytic

Feature	Galvanic (voltaic) cell	Electrolytic cell
Energy conversion	Chemical → electrical	Electrical → chemical
Anode sign	Negative	Positive
Cathode sign	Positive	Negative
Reaction at anode	Oxidation	Oxidation
Reaction at cathode	Reduction	Reduction
Spontaneity	E°cell > 0	Driven by external source

Students lose marks by assuming anode is always positive — it is positive only in electrolysis.

Edge Cases and Common Mistakes

Cl₂ is an oxidising agent but contains no oxygen — oxidising/reducing ability is defined by electron transfer, not by oxygen content.
Peroxides (H₂O₂, Na₂O₂) assign O = −1, not −2; fluorine combined with anything stays at −1.
Reversing E° subtraction gives a negative E°cell for spontaneous reactions — always subtract anode from cathode.
Nernst equation Ecell = E°cell − (0.0592/n) log₁₀ Q at 298 K is needed only when concentrations deviate from 1 M; standard NECO items usually stay at standard conditions.

Connections and Applications

Electroplating uses Faraday’s laws to deposit a thin metal coating (e.g. Cr on steel) for corrosion resistance.
Electrorefining of copper purifies blister copper using impure copper as anode and pure copper as cathode.
Extraction of reactive metals (Na, Al, Mg) needs electrolysis of fused salts because aqueous solution would discharge H⁺ first.
Rusting of iron is an undesired galvanic process: Fe acts as anode, O₂ at the cathode, with water as electrolyte.

Exam strategy: NECO Paper III structured questions on electrochemistry usually carry 8–12 marks split across (a) half-equation balancing, (b) identifying oxidising/reducing agents, and (c) a Faraday calculation. Allocate about 4 minutes and always write units in the final answer.

Practice Prompts

Balance the reaction between MnO₄⁻ and Fe²⁺ in acidic solution and identify the oxidising agent.
Calculate the time required to deposit 5.4 g of aluminium (Al = 27, n = 3) using a 3 A current.

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Electrochemistry: Redox and Cells