Concise Electrochemistry Notes

Electrochemical Potential

• Electrochemistry studies the relationship between electricity and chemical reactions.
• It's an interdisciplinary science applicable in engineering, chemistry, biology, and physics.

Electrochemical Cells

• Systems that incorporate a redox reaction to produce or utilize electrical energy.
• Two types:
  • Voltaic/Galvanic Cell: Releases free energy from a spontaneous reaction to produce electricity ($W’ = ∆G$).
  • Electrolytic Cell: Absorbs free energy from a source of electricity to drive a nonspontaneous reaction.

Voltaic/Galvanic Cell Construction

• Based on the spontaneity of a redox reaction where \Delta G < 0.
• Requires electrodes (anode and cathode), electrolytes (ionic salts), and half-cell separators (membranes, salt bridges).

Cell Potential ($E_{cell}$)

• The difference in electrical potential between the two electrodes.
• Standard cell potential ($E^o_{cell}$) is measured at specified conditions (298 K, 1 atm, 1 M).

Batteries

• Primary: Lithium ion batteries – types, working principles and advantages
• Secondary batteries

Standard Electrode Potential

• Measured using a reference half-cell, typically the Standard Hydrogen Electrode (SHE).
• SHE: $E^o = 0 \text{ V}$
• By convention, always refers to the half-reaction written as a reduction.

Calculation of Standard Cell Potential from Standard Half-Cell Potentials

• $E^o<em>{cell} = E^o</em>{\text{red. half-cell}} – E^o_{\text{ox. half-cell}}$
  • Example: $E<em>{zinc} = -0.76 V$ and $E</em>{copper} = 0.34 V$ resulting in $E_{ocell} = 1.1 V$
• Half-cell with a smaller or more negative $E^o$ acts as the anode; the half-cell with a larger or more positive $E^o$ acts as the cathode.

Electrochemical Series

• Lists elements and molecules based on their ability to oxidize or reduce.
• Used to predict whether a metal will react with a solution of another metal's ions.

Non-Standard State Cell Potential

• Nernst Equation: Used to calculate $E_{cell}$ under non-standard conditions.
• $\Delta G= \Delta G^o + RT \ln Q$
• $\Delta G= -nFE<em>{cell} \text{ and } \Delta G^o = - nFE^o</em>{cell}$
• Relates Gibbs energy to reaction quotient, showing how $E_{cell}$ varies with concentration.

Concentration Cells

• Galvanic cells with two equivalent half-cells but different ion concentrations.
• $E^o<em>{Cell} = 0$ * $E</em>{cell}$ depends only on the difference in concentration of ions.

Reversible vs Standard Hydrogen Electrode

• Concentration of $H^+$ remains constant in SHE
• Concentration of $H^+$ can change in RHE

Electrolytic Cells

• Use electrical energy to drive a nonspontaneous reaction (\Delta G > 0).
• Involve electroplating and recovering metals from ores.
• Faraday's Laws of Electrolysis:
  • Amount of chemical change is proportional to the quantity of electricity used.
  • Amounts of chemical changes are proportional to their equivalent weights.
• $F = 96485 \text{ C mol}^{-1}$
  • $F = N_A \times e = 6.023 \times 10^{23} \text{ mol}^{-1} \times 1.602 \times 10^{-19} \text{ C} = 96485 \text{ C mol}^{-1}$
• Key difference from voltaic cells: Anode is positive, and cathode is negative.