JEE Chemistry Electrochemistry Notes Set 01

Electrolytic Conductance

1. Conductance G is the reciprocal of resistance. Specific conductance (conductivity) \( \kappa=G\times l/A \), where \( l/A \) is the cell constant. Molar conductivity is \( \Lambda_m=\kappa\times1000/M \) and equivalent conductivity is \( \Lambda_{eq}=\kappa\times1000/N \).
2. On dilution the specific conductance falls, while molar and equivalent conductivities rise because ionisation and ionic mobility increase. Conductance of an electrolyte rises with temperature, unlike a metal.
3. Strong electrolytes show a linear fall of \( \Lambda_m \) with \( \sqrt{c} \) (Debye-Huckel-Onsager); weak electrolytes need extrapolation and their \( \Lambda^\infty \) is found from Kohlrausch's law.

Kohlrausch Law

1. At infinite dilution every ion contributes independently: \( \Lambda^\infty=\lambda^\infty_++\lambda^\infty_- \). It is used to find \( \Lambda^\infty \) of weak electrolytes, for example \( \Lambda^\infty_{CH_3COOH}=\Lambda^\infty_{CH_3COONa}+\Lambda^\infty_{HCl}-\Lambda^\infty_{NaCl} \).
2. Degree of dissociation \( \alpha=\Lambda_m/\Lambda^\infty \). For a sparingly soluble salt, solubility \( S=\kappa\times1000/\Lambda^\infty \).

Galvanic Cells and Electrode Potential

1. In a galvanic cell oxidation occurs at the anode (negative) and reduction at the cathode (positive). The cell is written anode on the left, for example Zn | Zn2+ || Cu2+ | Cu, and \( E^\circ_{cell}=E^\circ_{cathode}-E^\circ_{anode} \).
2. The standard hydrogen electrode (Pt, H2 at 1 bar, 1 M H+) is assigned 0 V. A more negative E° means a stronger reducing agent; a more positive E° means a stronger oxidising agent.
3. A metal with lower reduction potential displaces a metal of higher potential from its salt solution, and metals above hydrogen liberate H2 from dilute acids. E° is an intensive property and is not multiplied by coefficients, while \( \Delta G^\circ=-nFE^\circ \) is extensive.

Nernst Equation and Equilibrium

1. At 298 K, \( E=E^\circ-\dfrac{0.0591}{n}\log Q \). For a hydrogen electrode \( E=-0.059\,pH \) and for a concentration cell \( E=\dfrac{0.0591}{n}\log\dfrac{c_2}{c_1} \).
2. At equilibrium E = 0 and \( E^\circ=\dfrac{0.0591}{n}\log K_c \). A reaction is feasible when E° is positive.
3. When half reactions with different numbers of electrons are combined, add the \( \Delta G^\circ \) values (nFE°) and not the potentials.

Electrolysis and Faraday's Laws

1. First law: mass deposited \( w=ZIt=\dfrac{E\,I\,t}{96500} \). Second law: for the same charge, masses deposited are proportional to equivalent weights.
2. One faraday (96500 C) liberates one gram equivalent of any substance. Aqueous Na2SO4 gives H2 at the cathode and O2 at the anode; fused salts and hydrides give the metal or gas from the ions present.
3. At the anode of an electrolytic cell oxidation occurs (positive electrode), and at the cathode reduction occurs (negative electrode).