Electrochemistry

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Electrochemistry — Key Facts for MDCAT

Galvanic (Voltaic) Cell: Spontaneous redox reaction → generates electrical energy Electrolytic Cell: Electrical energy → drives non-spontaneous reaction

Key Definitions:

• Anode: Electrode where oxidation occurs (electrons are lost); in galvanic cells, it’s the negative electrode
• Cathode: Electrode where reduction occurs (electrons are gained); in galvanic cells, it’s the positive electrode
• Salt Bridge: Maintains electrical neutrality by allowing ion flow (KNO$_3$ or NH$_4$NO$_3$ are common)

Cell EMF: $$E_{cell}^\circ = E_{cathode}^\circ - E_{anode}^\circ = E_{reduction}^\circ - E_{oxidation}^\circ$$ Standard conditions: 1 M concentration, 1 atm pressure, 25°C (298 K)

Standard Reduction Potentials (selected):

Half-reaction	$E^\circ$ (V)
F$_2$ + 2e$^-$ → 2F$^-$	+2.87
O$_2$ + 4H$^+$ + 4e$^-$ → 2H$_2$O	+1.23
Ag$^+$ + e$^-$ → Ag	+0.80
Fe$^{3+}$ + e$^-$ → Fe$^{2+}$	+0.77
O$_2$ + 2H$_2$O + 4e$^-$ → 4OH$^-$	+0.40
Cu$^{2+}$ + 2e$^-$ → Cu	+0.34
2H$^+$ + 2e$^-$ → H$_2$	0.00 (reference)
Fe$^{2+}$ + 2e$^-$ → Fe	-0.44
Zn$^{2+}$ + 2e$^-$ → Zn	-0.76
Al$^{3+}$ + 3e$^-$ → Al	-1.66

More positive $E^\circ$ = stronger oxidising agent (more readily reduced).

⚡ Exam tip: In a galvanic cell, the electrode with HIGHER (more positive) reduction potential becomes the cathode. Zinc ($E^\circ = -0.76$ V) is oxidised (anode) when coupled with copper ($E^\circ = +0.34$ V) in a Daniell cell. $E_{cell} = +0.34 - (-0.76) = +1.10$ V. Always subtract the anode potential from the cathode.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students who want genuine understanding.

Electrochemistry — Complete Study Guide

Nernst Equation: For non-standard conditions: $$E = E^\circ - \frac{RT}{nF}\ln Q$$ At 25°C: $E = E^\circ - \frac{0.0592}{n}\log_{10} Q$

Where $n$ = number of electrons transferred, $F$ = Faraday constant (96,485 C/mol), $Q$ = reaction quotient.

For the cell reaction: $aA + bB \rightarrow cC + dD$ $$Q = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$

Relationship between $E$, $\Delta G$, and $K_{eq}$: $$\Delta G^\circ = -nFE^\circ$$ $$\Delta G^\circ = -RT\ln K_{eq}$$ $$\therefore E^\circ = \frac{RT}{nF}\ln K_{eq} = \frac{0.0592}{n}\log_{10}K_{eq}$$

When $E_{cell} = 0$ (equilibrium): $Q = K_{eq}$

Faraday’s Laws of Electrolysis:

1. First law: Mass deposited $m = \frac{Q \times M}{n \times F} = \frac{It \times M}{n \times F}$
2. Second law: Equal quantities of electricity deposit equal equivalents of different substances

Where: $Q$ = charge (coulombs), $I$ = current (amperes), $t$ = time (seconds), $M$ = molar mass, $n$ = electrons per ion.

Electrolytic Refining of Copper: Impure copper is the anode, pure copper is the cathode, CuSO$_4$ solution is the electrolyte. At anode: Cu → Cu$^{2+}$ + 2e$^-$ (impurities fall as sludge). At cathode: Cu$^{2+}$ + 2e$^-$ → Cu (pure)

Concentration Cells: Same electrodes and redox couple but different concentrations. $E_{cell} = E^\circ - \frac{0.0592}{n}\log\frac{[\text{low}]}{[\text{high}]} = \frac{0.0592}{n}\log\frac{[\text{high}]}{[\text{low}]}$

⚡ Common mistakes: Forgetting that $n$ in the Nernst equation is the total electrons transferred in the balanced equation (not just per half-reaction). Confusing the sign of $E_{cell}$ — a positive $E_{cell}$ means the reaction is spontaneous as written. For concentration cells, the half-cell with lower concentration is the anode.

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Electrochemistry — Advanced Notes

Types of Electrodes:

1. Metal-metal ion electrode: $M^{n+}(aq) + ne^- \rightleftharpoons M(s)$
2. Gas-gas ion electrode: $2H^+(aq) + 2e^- \rightleftharpoons H_2(g)$ (standard hydrogen electrode, SHE)
3. Quinhydrone electrode: For measuring pH
4. Glass electrode: For pH measurement (used in pH meters)

Reference Electrodes:

• Standard Hydrogen Electrode (SHE): $E^\circ = 0.00$ V by definition. Pt electrode with H$_2$ gas at 1 atm and H$^+$ at 1 M.
• Calomel electrode: Hg$_2$Cl$_2$/Hg electrode with KCl salt bridge. $E = +0.24$ V (saturated calomel, SCE).
• Silver-silver chloride electrode: $E = +0.22$ V (saturated KCl). More stable and reproducible than calomel.

Decomposition Potential: Minimum voltage needed to start electrolysis. For NaCl brine (industrial chloro-alkali process): $$2NaCl + 2H_2O \rightarrow 2NaOH + H_2 + Cl_2$$ Required voltage ~2.2 V (vs theoretical ~1.9 V). The excess is due to overvoltage (kinetic barrier).

Overvoltage (Overpotential): The extra voltage beyond the theoretical value needed to initiate/bypass the kinetic barriers of electrode reactions. Overvoltage is higher for gas evolution (hydrogen and oxygen bubbles must form).

Commercial Cells:

Cell	Reaction	Application
Daniell cell	Zn + CuSO$_4$ → ZnSO$_4$ + Cu	Early battery
Leclanche cell	NH$_4$Cl + Zn → ZnCl$_2$ + NH$_3$ + H$_2$	Flashlight battery
Alkaline cell	KOH + Zn → K$_2$ZnO$_2$ + H$_2$	Longer lasting
Lead-acid battery	Pb + PbO$_2$ + H$_2$SO$_4$ ⇌ PbSO$_4$	Car battery (rechargeable)
Fuel cell	H$_2$ + ½O$_2$ → H$_2$O	Clean energy (Apollo missions)

Fuel Cells: Galvanic cell where reactants (fuel + oxidant) are supplied continuously from external reservoirs. Hydrogen-oxygen fuel cell:

• Anode: $2H_2 + 4OH^- \rightarrow 4H_2O + 4e^-$ (in alkaline medium)
• Cathode: $O_2 + 2H_2O + 4e^- \rightarrow 4OH^-$
• Overall: $2H_2 + O_2 \rightarrow 2H_2O$, $E^\circ = +1.23$ V
• Efficiency: ~60–70% (vs ~40% for internal combustion engines)

Corrosion: Iron rusting is an electrochemical process:

• Anode (Fe): $Fe \rightarrow Fe^{2+} + 2e^-$
• Cathode (in contact with Fe): $2H^+ + \frac{1}{2}O_2 \rightarrow H_2O$ (in acidic); or $O_2 + 2H_2O + 4e^- \rightarrow 4OH^-$ (in neutral/alkaline)
• Rust = $Fe_2O_3 \cdot xH_2O$

Prevention: Cathodic protection (sacrificial anode, e.g., Mg or Zn attached to iron pipeline), painting, galvanising (zinc coating), electroplating.

MDCAT Question Patterns: MDCAT Pakistan electrochemistry questions frequently test: (1) calculating cell potential using $E_{cell} = E_{cathode} - E_{anode}$, (2) Nernst equation calculations, (3) Faraday’s laws of electrolysis (mass deposited calculations), (4) identifying cathode and anode in both galvanic and electrolytic cells, (5) relationship $\Delta G = -nFE$. 2–3 questions per paper.

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