Electrochemistry — Revision Notes

⚡ 30-Second Revision

Key formulas, constants, and definitions:\n* Redox: Oxidation (loss of $e^-$ at anode), Reduction (gain of $e^-$ at cathode).\n* Galvanic Cell: Spontaneous, chemical to electrical, anode (-), cathode (+).

\n* Electrolytic Cell: Non-spontaneous, electrical to chemical, anode (+), cathode (-).\n* Cell Potential: $E_{cell} = E_{cathode} - E_{anode}$.\n* Nernst Equation (298 K): $E_{cell} = E^\circ_{cell} - \frac{0.

0592}{n} \log Q$.\n* **Gibbs Free Energy:**$\Delta G = -nFE_{cell}$. For spontaneity,$\Delta G < 0$and$E_{cell} > 0$.\n* **Equilibrium Constant:**$E^\circ_{cell} = \frac{0.0592}{n} \log K_c$ (at 298 K).

\n* Faraday's First Law: $m = ZQ = \frac{E_w \times Q}{F}$.\n* Faraday's Constant: $F = 96485\ C/mol\ e^-$.\n* Molar Conductivity: $\Lambda_m = \frac{\kappa \times 1000}{C}$.\n* Kohlrausch's Law: $\Lambda^\circ_m = x\lambda^\circ_A + y\lambda^\circ_B$ (at infinite dilution).

2-Minute Revision

Electrochemistry is about the interconversion of chemical and electrical energy through redox reactions. Galvanic cells (like batteries) convert spontaneous chemical reactions into electricity, with oxidation at the negative anode and reduction at the positive cathode.

Electrolytic cells use external electricity to drive non-spontaneous reactions, with oxidation at the positive anode and reduction at the negative cathode. The cell potential ($E_{cell}$) is calculated from standard electrode potentials ($E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$) and adjusted for non-standard concentrations using the Nernst equation: $E_{cell} = E^\circ_{cell} - \frac{0.

0592}{n} \log Q$. Spontaneity is linked by$\Delta G = -nFE_{cell}$, where$\Delta G < 0$for spontaneous reactions, implying$E_{cell} > 0$. Faraday's laws quantify electrolysis: mass deposited is proportional to charge passed ($m = ZQ$), and for the same charge, masses are proportional to equivalent weights.

Conductivity of solutions depends on ion concentration and mobility, with molar conductivity ($\Lambda_m$) increasing with dilution for weak electrolytes due to increased dissociation, and slightly for strong electrolytes due to reduced interionic attraction.

Kohlrausch's law helps calculate $\Lambda^\circ_m$ for weak electrolytes. Corrosion is an electrochemical process of metal degradation.

5-Minute Revision

Electrochemistry fundamentally explores the relationship between chemical reactions and electrical energy, primarily through redox processes. We distinguish between two main types of electrochemical cells: galvanic (or voltaic) cells and electrolytic cells.

Galvanic cells harness spontaneous chemical reactions to produce electrical energy, exemplified by batteries. In these, oxidation occurs at the anode (negative electrode), and reduction at the cathode (positive electrode).

Electrons flow from anode to cathode through an external circuit, while a salt bridge maintains charge neutrality. The cell potential ($E_{cell}$) is a measure of the driving force, calculated as $E_{cell} = E_{cathode} - E_{anode}$.

For non-standard conditions, the Nernst equation, $E_{cell} = E^\circ_{cell} - \frac{0.0592}{n} \log Q$ (at 298 K), is used, where $Q$ is the reaction quotient and $n$ is the number of electrons transferred.

The spontaneity of a reaction is linked to cell potential by $\Delta G = -nFE_{cell}$, where a positive $E_{cell}$ corresponds to a negative $\Delta G$ (spontaneous). \n\nConversely, electrolytic cells use an external electrical energy source to drive non-spontaneous chemical reactions, such as electroplating or industrial production of elements.

Here, the anode is positive and the cathode is negative. The quantitative aspects of electrolysis are governed by Faraday's laws: the mass of substance deposited ($m$) is directly proportional to the charge passed ($Q$), given by $m = ZQ = \frac{E_w \times Q}{F}$, where $E_w$ is the equivalent weight and $F$ is Faraday's constant (96485 C/mol).

\n\nElectrolytic solutions conduct electricity due to ion movement. Their conductivity ($\kappa$) and molar conductivity ($\Lambda_m = \frac{\kappa \times 1000}{C}$) are crucial parameters. Molar conductivity for strong electrolytes increases slightly with dilution, while for weak electrolytes, it increases significantly due to enhanced dissociation.

Kohlrausch's law allows calculation of limiting molar conductivity (at infinite dilution) for weak electrolytes by combining values from strong electrolytes. Finally, corrosion, like rusting, is an electrochemical process where metals oxidize, and its prevention often involves electrochemical principles like cathodic protection.

Prelims Revision Notes

1
Redox Basics: — Oxidation = loss of electrons, increase in oxidation state. Reduction = gain of electrons, decrease in oxidation state. Always occur together.\n2. Electrochemical Cells:\n * Galvanic (Voltaic): Chemical $\rightarrow$ Electrical. Spontaneous ($\Delta G < 0$, $E_{cell} > 0$). Anode is negative, cathode is positive. Electrons flow anode to cathode. Salt bridge maintains neutrality.\n * Electrolytic: Electrical $\rightarrow$ Chemical. Non-spontaneous ($\Delta G > 0$, $E_{cell} < 0$). Anode is positive, cathode is negative. External power source required.\n3. **Electrode Potential ($E$):** Tendency of an electrode to gain/lose electrons. Standard Electrode Potential ($E^\circ$) at 1 M, 1 atm, 298 K. SHE ($E^\circ = 0\ V$) is reference.\n4. **Cell Potential ($E_{cell}$):** $E_{cell} = E_{cathode} - E_{anode}$. For standard: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$.\n5. Nernst Equation: $E_{cell} = E^\circ_{cell} - \frac{RT}{nF} \ln Q$. At 298 K: $E_{cell} = E^\circ_{cell} - \frac{0.0592}{n} \log Q$. $Q = \frac{[Products]^{coeff}}{[Reactants]^{coeff}}$ (exclude pure solids/liquids). Important for non-standard conditions.\n6. Thermodynamic Relations:\n * $\Delta G = -nFE_{cell}$.\n * $\Delta G^\circ = -nFE^\circ_{cell}$.\n * At equilibrium ($E_{cell}=0$, $\Delta G=0$, $Q=K_c$): $E^\circ_{cell} = \frac{RT}{nF} \ln K_c = \frac{0.0592}{n} \log K_c$ (at 298 K).\n7. Conductivity:\n * **Specific Conductivity ($\kappa$):** Reciprocal of resistivity. Unit: $S\ cm^{-1}$ or $S\ m^{-1}$.\n * **Molar Conductivity ($\Lambda_m$):** $\Lambda_m = \frac{\kappa \times 1000}{C}$ (if $\kappa$ in $S\ cm^{-1}$, $C$ in $mol/L$). Unit: $S\ cm^2\ mol^{-1}$.\n * Variation with Dilution: Strong electrolytes (slight increase due to reduced interionic attraction). Weak electrolytes (significant increase due to increased dissociation).\n8. Kohlrausch's Law: At infinite dilution, $\Lambda^\circ_m = x\lambda^\circ_A + y\lambda^\circ_B$. Useful for weak electrolytes (e.g., $\Lambda^\circ_m(CH_3COOH) = \Lambda^\circ_m(CH_3COONa) + \Lambda^\circ_m(HCl) - \Lambda^\circ_m(NaCl)$).\n9. Faraday's Laws of Electrolysis:\n * First Law: $m \propto Q \Rightarrow m = ZQ = \frac{E_w \times Q}{F}$. $Q = I \times t$. $F = 96485\ C/mol\ e^-$.\n * Second Law: For same $Q$, $\frac{m_1}{m_2} = \frac{E_{w1}}{E_{w2}}$.\n * **Equivalent Weight ($E_w$):** $E_w = \frac{\text{Molar Mass}}{\text{n-factor}}$ (n-factor = electrons transferred per formula unit).\n10. Batteries: Primary (non-rechargeable), Secondary (rechargeable), Fuel Cells (continuous supply of reactants, high efficiency, eco-friendly).\n11. Corrosion: Electrochemical process (e.g., rusting of iron: $Fe \rightarrow Fe^{2+}$ at anode, $O_2 \rightarrow H_2O$ at cathode). Prevention: coating, galvanization, cathodic protection.

Vyyuha Quick Recall

To remember the polarity of electrodes in different cells: \n'LEO the lion says GER' (Loss of Electrons is Oxidation, Gain of Electrons is Reduction) \n'AN OX' (ANode is OXidation) \n'RED CAT' (REDuction is at CAThode) \n\nFor polarity: \n'GALVANIC: A Negative Cat' (Anode is Negative, Cathode is Positive) \n'ELECTROLYTIC: A Positive Cat' (Anode is Positive, Cathode is Negative) \n\nThis helps distinguish the electrode polarities while keeping the LEO/GER and AN OX/RED CAT rules consistent for the processes occurring at them.