ELECTROCHEMISTRY - Class 12 Tamil Nadu Board Chemistry

Electrochemistry: Core Concepts and Applications

Electrochemical Cells and Their Construction

Electrochemical cells convert chemical energy into electrical energy through spontaneous redox reactions. The cell consists of two half-cells, each with a metal electrode dipped in its respective salt solution, connected by a salt bridge and an external circuit.

• Galvanic/Voltaic Cell: Another name for the electrochemical cell.
• Example: Daniel cell setup uses a zinc rod in ZnSO$_4$ solution (anode) and a copper rod in CuSO$_4$ solution (cathode).
• Spontaneity: The cell reaction is spontaneous; $\Delta G < 0$.
• Electron Flow: Electrons move from anode (Zn) to…

Electrochemistry: Core Concepts and Applications

Electrochemical Cells and Their Construction

Electrochemical cells convert chemical energy into electrical energy through spontaneous redox reactions. The cell consists of two half-cells, each with a metal electrode dipped in its respective salt solution, connected by a salt bridge and an external circuit.

• Galvanic/Voltaic Cell: Another name for the electrochemical cell.
• Example: Daniel cell setup uses a zinc rod in ZnSO$_4$ solution (anode) and a copper rod in CuSO$_4$ solution (cathode).
• Spontaneity: The cell reaction is spontaneous; $\Delta G < 0$.
• Electron Flow: Electrons move from anode (Zn) to cathode (Cu).
• Conventional Current: Flows from cathode to anode (opposite to electron flow).
• Salt Bridge: Maintains electrical neutrality and completes the circuit.

Anode, Cathode, and Cell Representation

The anode is the electrode where oxidation occurs and carries a negative charge, while the cathode is where reduction occurs and carries a positive charge. The cell is represented with the anode on the left and cathode on the right, separated by a salt bridge.

• Anode: Oxidation, negative charge, left side, e.g., Zn $\rightarrow$ Zn$^{2+}$ + 2e$^-$.
• Cathode: Reduction, positive charge, right side, e.g., Cu$^{2+}$ + 2e$^-$ $\rightarrow$ Cu.
• Cell Representation: Zn(s) | Zn$^{2+}$(aq) || Cu$^{2+}$(aq) | Cu(s).
• Oxidation State Change: Increase at anode (oxidation), decrease at cathode (reduction).

Salt Bridge: Function and Application

The salt bridge is a mixture of inert electrolyte and agar-agar, forming a paste that allows ion migration to maintain charge balance. It prevents liquid junction potential and ensures the cell reaction proceeds smoothly.

• Inert Electrolytes: KCl, KNO$_3$, etc., with equal mobility for cation and anion.
• Cation Movement: K$^+$ moves toward the compartment with excess SO$_4^{2-}$ to neutralize charge.
• Anion Movement: Cl$^-$ moves toward the compartment with excess Zn$^{2+}$.
• Functions: Maintains electrical neutrality, completes the inner circuit, avoids liquid junction potential.

Standard Electrode Potential and EMF Calculation

Each electrode develops a potential under standard conditions (standard temperature, pressure, and 1 M concentration). The standard electrode potential is always considered as the standard reduction potential.

• Standard EMF Formula: $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$
• Only Reduction Potentials: Always use reduction potentials for both electrodes.
• Example Calculation: For a cell with given $E^\circ$ values, substitute directly into the formula.

Nernst Equation and Non-Standard Conditions

The Nernst equation is used to calculate cell potential under non-standard conditions. It relates the cell EMF to the concentrations of reactants and products.

• Nernst Equation: $E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.06}{n} \log Q_c$
• $n$: Number of electrons transferred.
• $Q_c$: Reaction quotient, written using concentrations of aqueous species (solids and liquids have activity 1).
• Application: Substitute given concentrations and calculate $E_{\text{cell}}$.

Electrochemical Series and Its Implications

The electrochemical series arranges elements/ions in order of increasing standard reduction potential. Hydrogen is assigned a value of zero.

• Order: Li, K, Ba, Sr, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, H, Cu, I, Ag, Hg, Br, Pt, Cl, Au, F.
• Above H: Negative $E^\circ$ values, strong reducing agents.
• Below H: Positive $E^\circ$ values, strong oxidizing agents.
• Oxidizing Power: Increases with higher $E^\circ$.
• Reducing Power: Increases with lower $E^\circ$.
• Relation: $E^\circ_{\text{oxidation}} = -E^\circ_{\text{reduction}}$.

Electrolytic Cells and Products of Electrolysis

Electrolytic cells convert electrical energy into chemical energy. The process involves inert electrodes (Pt or graphite) and an external battery.

• Anode: Positive, oxidation occurs.
• Cathode: Negative, reduction occurs.
• Molten Electrolyte Example: Electrolysis of molten NaCl yields Na at cathode and Cl$_2$ at anode.
• Aqueous Electrolyte: Water ions (H$^+$, OH$^-$) may compete with electrolyte ions; discharge depends on reduction potential.
• Discharge Order (Cations): Ion with higher reduction potential is reduced at cathode.
• Discharge Order (Anions): SO$_4^{2-}$ < NO$_3^-$ < OH$^-$ < Cl$^-$ < Br$^-$ < I$^-$.

Faraday’s Laws of Electrolysis

Faraday’s laws relate the amount of substance deposited or liberated at an electrode to the quantity of electricity passed.

• First Law: Mass deposited ($w$) is directly proportional to charge ($q$): $w \propto q$.
• $q = It$: Where $I$ is current, $t$ is time.
• Electrochemical Equivalent: $w = \frac{E \cdot q}{nF}$, where $E$ is equivalent weight, $n$ is number of electrons, $F$ is Faraday.
• Number of Moles: $n = \frac{q}{nF}$.
• Charge for 1 mol e$^-$: 1 Faraday.
• Example: To deposit 1 mol Al from Al$^{3+}$, 3 Faradays are required.

Resistance, Conductance, and Conductivity

Resistance ($R$) and conductance ($G$) are fundamental to understanding electrolytic conduction. Specific resistance (resistivity, $\rho$) and specific conductance (conductivity, $\kappa$) are also defined.

• Resistance: $R = \rho \frac{l}{a}$, unit: ohm ($\Omega$).
• Conductance: $G = \frac{1}{R}$, unit: siemens (S).
• Resistivity: $\rho$, unit: $\Omega$ cm.
• Conductivity: $\kappa = \frac{1}{\rho}$, unit: S cm$^{-1}$.
• Cell Constant: $l/a$, relates measured conductance to conductivity.

Molar and Equivalent Conductivity

Molar conductivity ($\Lambda_m$) and equivalent conductivity ($\Lambda_{eq}$) are measures of the conducting power of all ions produced by one mole or one equivalent of an electrolyte.

• Molar Conductivity: $\Lambda_m = \frac{1000 \kappa}{M}$, $M$ = molarity.
• Equivalent Conductivity: $\Lambda_{eq} = \frac{1000 \kappa}{N}$, $N$ = normality.
• Units: S cm$^2$ mol$^{-1}$ for molar, S cm$^2$ eq$^{-1}$ for equivalent.
• Variation with Concentration: For strong electrolytes, $\Lambda_m$ vs. $\sqrt{c}$ is a straight line with negative slope; for weak electrolytes, it is a curve.

Kohlrausch’s Law and Its Applications

Kohlrausch’s law states that at infinite dilution, the molar conductivity of an electrolyte is the sum of the individual contributions of its ions.

• General Form: For $AXBY$, $\Lambda_m^\circ (AXBY) = x \lambda_m^\circ (A^{y+}) + y \lambda_m^\circ (B^{x-})$.
• Application: Used to calculate degree of ionization ($\alpha$) and solubility ($S$) of sparingly soluble salts.
• Degree of Ionization: $\alpha = \frac{\Lambda_m}{\Lambda_m^\circ}$.
• Solubility Calculation: $K_{sp} = x^x y^y S^{x+y}$.

Batteries and Fuel Cells

Batteries are practical applications of electrochemical cells. They are classified as primary (non-rechargeable) and secondary (rechargeable).

• Primary Batteries: Leclanche cell (anode: Zn, cathode: MnO$_2$), Mercury cell (anode: Zn-Hg amalgam, cathode: HgO).
• Secondary Batteries: Lead storage battery (anode: Pb, cathode: PbO$_2$), Nickel-cadmium cell (anode: Cd, cathode: Ni(OH)$_3$).
• Fuel Cell: Hydrogen-oxygen fuel cell overall reaction: $2H_2 + O_2 \rightarrow 2H_2O$.

Corrosion

Corrosion is an electrochemical process where metals are oxidized, often leading to rust formation in iron.

• Anode Reaction: Fe $\rightarrow$ Fe$^{2+}$.
• Cathode Reaction: O$_2$ is reduced to H$_2$O.
• Further Oxidation: Fe$^{2+}$ can be oxidized to Fe$^{3+}$, forming Fe$_2$O$_3 \cdot xH_2O$ (rust).

Key Exam Points and Problem-Solving Strategies

For board exams, focus on direct application of formulas, clear identification of anode/cathode, and correct use of reduction potentials. Practice with book-back questions and ensure clarity in cell representation, calculation of EMF, and application of Faraday’s and Kohlrausch’s laws.

• Always use reduction potentials for EMF calculations.
• Identify anode/cathode based on oxidation/reduction and sign conventions.
• Apply Nernst equation for non-standard conditions.
• Use correct units and conversion factors in conductivity calculations.
• For electrolysis, determine which ion is discharged based on reduction potential order.
• In batteries, memorize anode/cathode materials and overall reactions.
• For corrosion, write both anodic and cathodic reactions and final rust formula.