Electrochemistry — MHT-CET Chemistry

Formulas (3)

• Conductance, Conductivity and Their Units · Conductivity
  $$\kappa = \frac{1}{\rho} = G\cdot\frac{l}{a} = \frac{1}{R}\cdot\frac{l}{a}$$
• Cell Constant: l/a = κ × R · Cell constant
  $$\frac{l}{a} = \kappa \times R$$
• Conductivity From the Cell Constant and Resistance · Conductivity of the unknown
  $$\kappa = \frac{\text{cell constant}}{R}$$

Watch out for (3)

• Reading S cm² mol⁻¹ as the unit of conductivity
  → Conductance, Conductivity and Their Units
• Dividing when the cell constant is asked
  → Cell Constant: l/a = κ × R
• Trusting a printed exponent over the order of magnitude
  → Conductivity From the Cell Constant and Resistance

Formulas (3)

• Molar Conductivity: Λ = 1000κ/c and Back · Molar conductivity
  $$\Lambda = \frac{1000\,\kappa}{c},\qquad \kappa = \frac{\Lambda\,c}{1000}$$
• Kohlrausch's Law: Λ₀ From the Ions · Kohlrausch's law
  $$\Lambda_0 = \nu_+\,\lambda_+^0 + \nu_-\,\lambda_-^0$$
• Degree of Dissociation: α = Λ/Λ₀ · Degree of dissociation
  $$\alpha = \frac{\Lambda_c}{\Lambda_0}$$

Watch out for (3)

• Dropping the 1000
  → Molar Conductivity: Λ = 1000κ/c and Back
• Adding all three values
  → Kohlrausch's Law: Λ₀ From the Ions
• Inverting the ratio
  → Degree of Dissociation: α = Λ/Λ₀

Formulas (3)

• What Forms at Each Electrode: Molten Versus Aqueous NaCl · Electrode reactions
  $$\text{cathode: } M^{n+} + n e^- \to M;\qquad \text{anode: } 2X^- \to X_2 + 2e^-$$
• Faraday's First Law: W = ItM/nF · Faraday's first law
  $$W = \frac{I\,t\,M}{n\,F},\qquad Q = I\,t,\qquad F = 96500\ \text{C mol}^{-1}$$
• Charge for a Redox Change, and Cells in Series · Charge for n electrons; series cells
  $$Q = \text{mol} \times n \times F,\qquad \frac{W_1}{W_2} = \frac{M_1/n_1}{M_2/n_2}$$

Watch out for (3)

• Sodium at the cathode from BRINE
  → What Forms at Each Electrode: Molten Versus Aqueous NaCl
• Using n = 1 for a divalent metal
  → Faraday's First Law: W = ItM/nF
• n = 3 for dichromate
  → Charge for a Redox Change, and Cells in Series

Formulas (6)

• Reading a Cell: Anode Left, Cathode Right · Cell notation
  $$\text{anode (–)}\ \big|\ \text{anode ion}\ \big\|\ \text{cathode ion}\ \big|\ \text{cathode (+)}$$
• E°cell = E°cathode − E°anode · Standard emf
  $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$
• The Electrochemical Series: Who Reduces, Who Oxidises, Who Deposits · Spontaneity test
  $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} > 0 \iff \text{spontaneous}$$
• The Nernst Equation: E = E° − (0.0592/n) log Q · Nernst equation (298 K)
  $$E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0592}{n}\log_{10} Q$$
• Potential of One Electrode at a Given Concentration · Oxidation electrode potential
  $$E_{\text{ox}} = -E^\circ_{\text{red}} - \frac{0.0592}{n}\log[\text{M}^{n+}]$$
• ΔG° = −nFE° and the Bridge to K · Gibbs energy and K
  $$\Delta G^\circ = -nFE^\circ,\qquad E^\circ = \frac{0.0592}{n}\log_{10} K$$

Watch out for (6)

• Anode = positive
  → Reading a Cell: Anode Left, Cathode Right
• Adding the two potentials
  → E°cell = E°cathode − E°anode
• Picking F⁻ as the strongest oxidising agent
  → The Electrochemical Series: Who Reduces, Who Oxidises, Who Deposits
• Forgetting the square on [Ag⁺]
  → The Nernst Equation: E = E° − (0.0592/n) log Q
• Doubling E° when the equation is doubled
  → Potential of One Electrode at a Given Concentration
• Losing the minus sign
  → ΔG° = −nFE° and the Bridge to K

Formulas (1)

• The Hydrogen–Oxygen Fuel Cell · Fuel cell net reaction
  $$2\text{H}_2(g) + \text{O}_2(g) \to 2\text{H}_2\text{O}(l),\qquad E^\circ = 1.23\ \text{V}$$

Reference tables (1)

Watch out for (2)

• Swapping discharge and recharge at the positive plate
  → Dry Cell, Lead Accumulator and Nickel–Cadmium Cell
• Calling H₂ the oxidising agent because it 'burns'
  → The Hydrogen–Oxygen Fuel Cell