CHEM1011 · Chemistry 1a
Electrochemistry
Week 9 covers redox: assigning oxidation numbers, balancing half-equations in acidic and basic solution, and building galvanic cells whose standard potential E°cell = E°(cathode) − E°(anode) predicts spontaneity. The Nernst equation for non-standard conditions, the links to ΔG° and K, and corrosion and electrolysis are Mastery content. This is the topic examined in In-Term Test 2, so cell potential and Nernst calculations are near-certain items.
What this chapter covers
- 01Oxidation (loss of electrons, oxidation number up) and reduction (gain, number down); OIL RIG; oxidising vs reducing agent
- 02Oxidation-number rules and balancing redox half-equations in acidic (H₂O, H⁺, e⁻) and basic (add OH⁻) solution
- 03Galvanic cells: anode = oxidation (−), cathode = reduction (+); electron flow anode → cathode; the salt bridge
- 04Cell notation anode | anode solution ‖ cathode solution | cathode; the standard hydrogen electrode (E° = 0)
- 05Standard cell potential E°cell = E°(cathode) − E°(anode); positive E°cell → spontaneous
- 06Thermodynamic links ΔG° = −nFE°cell and E°cell = (RT/nF) ln K (Mastery)
- 07Nernst equation E = E°cell − (0.0592/n)·log Q at 298 K; concentration cells (Mastery)
- 08Electrolysis (Faraday's law m = ItM/nF) and corrosion / cathodic protection (Mastery)
Cell potential under non-standard conditions with the Nernst equation
- +1Silver has the higher reduction potential, so Ag⁺/Ag is the cathode and Cu²⁺/Cu the anode. E°cell = E°(cathode) − E°(anode) = 0.80 − 0.34 = +0.46 V, with n = 2 electrons transferred.
- +1Write Q for Cu(s) + 2Ag⁺ → Cu²⁺ + 2Ag (solids omitted): Q = [Cu²⁺]/[Ag⁺]² = 0.10/(0.010)² = 0.10/1.0 × 10⁻⁴ = 1.0 × 10³.
- +1Apply the Nernst equation: E = E°cell − (0.0592/n)·log Q = 0.46 − (0.0592/2)·log(1.0 × 10³) = 0.46 − (0.0296)(3).
- +1E = 0.46 − 0.089 = +0.37 V. The potential is still positive (spontaneous), just lower than E°cell because the product ion Cu²⁺ is more concentrated than the reactant Ag⁺.
Key terms
- Oxidation and reduction
- Oxidation is loss of electrons (oxidation number increases); reduction is gain (number decreases) — OIL RIG. The oxidising agent is itself reduced; the reducing agent is itself oxidised.
- Galvanic (voltaic) cell
- A cell in which a spontaneous redox reaction produces electricity. The anode is the site of oxidation (negative terminal) and the cathode of reduction (positive); electrons flow anode → cathode and a salt bridge keeps the solutions electroneutral.
- Standard cell potential (E°cell)
- E°cell = E°(cathode) − E°(anode), both as standard reduction potentials. A positive E°cell means a spontaneous cell reaction; it links to free energy by ΔG° = −nFE°cell.
- Standard hydrogen electrode (SHE)
- The reference half-cell (Pt, H₂ at 1 bar, [H⁺] = 1 M) assigned E° = 0.00 V, against which all other standard reduction potentials are measured.
- Nernst equation
- E = E°cell − (RT/nF) ln Q, or at 298 K E = E°cell − (0.0592/n)·log Q; it gives the cell potential under non-standard concentrations (Mastery).
- Faraday's law
- In electrolysis the mass of product is m = (I·t·M)/(nF), where charge Q = I·t, F = 96485 C mol⁻¹ and n is the electrons per formula unit (Mastery).
Electrochemistry FAQ
How do I balance a redox reaction by half-equations?
Split the reaction into an oxidation and a reduction half-reaction. Balance the atoms other than O and H, then balance O with H₂O, balance H with H⁺ (in acidic solution), and balance charge with electrons. Multiply the two halves so the electrons cancel, then add them. In basic solution, add enough OH⁻ to both sides to neutralise the H⁺ to water. The electrons must cancel exactly — that is your check.
How do I find E°cell and know if a cell is spontaneous?
Identify the two reduction potentials. The couple with the higher (more positive) E° is reduced and becomes the cathode; the lower is oxidised and becomes the anode. Then E°cell = E°(cathode) − E°(anode), using both as reduction potentials — the formula's minus sign already handles the reversal. A positive E°cell means the reaction is spontaneous, matching a negative ΔG° through ΔG° = −nFE°cell.
When do I use the Nernst equation instead of E°cell?
Use E°cell only at standard conditions (all species at 1 M / 1 bar). As soon as concentrations differ from standard, the potential changes and you need the Nernst equation, E = E°cell − (0.0592/n)·log Q at 298 K. It also handles concentration cells (same couple both sides, E°cell = 0, potential purely from the concentration difference), which underlie pH electrodes and biological membrane potentials.
Can Sia help me with redox balancing and cell calculations?
Yes. Sia can assign oxidation numbers, balance half-equations in acidic or basic solution, build the cell notation, compute E°cell and ΔG°, and run the Nernst equation for a non-standard or concentration cell. It explains the method and checks your working; it does not do graded assessment, and UNSW academic-integrity rules apply.
Exam move
Since electrochemistry is examined in In-Term Test 2 and again in the final, split it into skills and drill each. Make oxidation-number assignment and half-equation balancing mechanical, including the extra OH⁻ step for basic solution and the electrons-must-cancel check. For cells, always identify cathode/anode from the reduction potentials first, then apply E°cell = E°(cathode) − E°(anode) without hand-flipping signs, and be able to draw the labelled cell (anode, cathode, salt bridge, electron flow) and write the line notation. Treat ΔG° = −nFE°cell as the bridge that ties this topic to thermochemistry. For the Mastery layer, drill the Nernst equation with careful attention to Q (correct exponents from the balanced equation) and n, and practise the concentration-cell and 'flat cell' (E = 0, Q = K) variants. Keep Faraday's law (m = ItM/nF) and the corrosion/cathodic-protection logic ready for short-answer items. When a Nernst or balancing step slips, ask Sia to redo it step by step.
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