1. Introduction to Electrochemistry
Electrochemistry is the branch of chemistry that studies the relationship between electricity and
chemical reactions, especially those involving electron transfer (redox reactions). It explains how
chemical energy can be converted to electrical energy and vice versa.
2. Important Terms
Electrolyte: Substance that produces ions and conducts electricity in molten or aqueous
state (e.g., NaCl, H₂SO₄).
Electrode: Conductor through which current enters or leaves an electrolyte.
Anode: Electrode where oxidation occurs (loss of electrons).
Cathode: Electrode where reduction occurs (gain of electrons).
Electrochemical Cell: Device that converts chemical energy into electrical energy or
vice versa.
3. Redox Reactions in Electrochemistry
Oxidation: Loss of electrons.
Reduction: Gain of electrons.
Redox reaction: Simultaneous oxidation and reduction.
Example:
Zn (s)→Zn2++2e−(oxidation at anode)\text{Zn (s)} \to \text{Zn}^{2+} + 2e^- \quad
(\text{oxidation at anode})Zn (s)→Zn2++2e−(oxidation at anode)
Cu2++2e−→Cu (s)(reduction at cathode)\text{Cu}^{2+} + 2e^- \to \text{Cu (s)} \quad
(\text{reduction at cathode})Cu2++2e−→Cu (s)(reduction at cathode)
4. Types of Electrochemical Cells
a) Galvanic (Voltaic) Cell
Generates electricity spontaneously from redox reactions.
Anode is negative; cathode is positive.
Electrons flow through external circuit from anode to cathode.
Example: Daniell Cell
, Anode: Zinc electrode in ZnSO₄ solution.
Cathode: Copper electrode in CuSO₄ solution.
Reactions:
Anode (oxidation):
Zn (s)→Zn2++2e−\text{Zn (s)} \to \text{Zn}^{2+} + 2e^-Zn (s)→Zn2++2e−
Cathode (reduction):
Cu2++2e−→Cu (s)\text{Cu}^{2+} + 2e^- \to \text{Cu (s)}Cu2++2e−→Cu (s)
b) Electrolytic Cell
Uses external electrical energy to cause a non-spontaneous chemical reaction.
Anode is positive; cathode is negative.
Example: Electrolysis of molten NaCl produces Na metal and Cl₂ gas.
5. Electrode Potential and Standard Electrode Potential
Electrode potential is the potential difference developed at the interface of electrode and
electrolyte.
Measured against standard hydrogen electrode (SHE), which is assigned 0 V.
Standard electrode potential (E°) is measured under standard conditions (1 M, 1 atm,
25°C).
6. Electrochemical Series
List of elements arranged by their standard electrode potentials.
Elements with more negative E° oxidize easily (strong reducing agents).
Elements with more positive E° are less reactive (strong oxidizing agents).
7. Nernst Equation
Calculates electrode potential under non-standard conditions:
E=E∘−RTnFlnQE = E^\circ - \frac{RT}{nF} \ln QE=E∘−nFRTlnQ
Electrochemistry is the branch of chemistry that studies the relationship between electricity and
chemical reactions, especially those involving electron transfer (redox reactions). It explains how
chemical energy can be converted to electrical energy and vice versa.
2. Important Terms
Electrolyte: Substance that produces ions and conducts electricity in molten or aqueous
state (e.g., NaCl, H₂SO₄).
Electrode: Conductor through which current enters or leaves an electrolyte.
Anode: Electrode where oxidation occurs (loss of electrons).
Cathode: Electrode where reduction occurs (gain of electrons).
Electrochemical Cell: Device that converts chemical energy into electrical energy or
vice versa.
3. Redox Reactions in Electrochemistry
Oxidation: Loss of electrons.
Reduction: Gain of electrons.
Redox reaction: Simultaneous oxidation and reduction.
Example:
Zn (s)→Zn2++2e−(oxidation at anode)\text{Zn (s)} \to \text{Zn}^{2+} + 2e^- \quad
(\text{oxidation at anode})Zn (s)→Zn2++2e−(oxidation at anode)
Cu2++2e−→Cu (s)(reduction at cathode)\text{Cu}^{2+} + 2e^- \to \text{Cu (s)} \quad
(\text{reduction at cathode})Cu2++2e−→Cu (s)(reduction at cathode)
4. Types of Electrochemical Cells
a) Galvanic (Voltaic) Cell
Generates electricity spontaneously from redox reactions.
Anode is negative; cathode is positive.
Electrons flow through external circuit from anode to cathode.
Example: Daniell Cell
, Anode: Zinc electrode in ZnSO₄ solution.
Cathode: Copper electrode in CuSO₄ solution.
Reactions:
Anode (oxidation):
Zn (s)→Zn2++2e−\text{Zn (s)} \to \text{Zn}^{2+} + 2e^-Zn (s)→Zn2++2e−
Cathode (reduction):
Cu2++2e−→Cu (s)\text{Cu}^{2+} + 2e^- \to \text{Cu (s)}Cu2++2e−→Cu (s)
b) Electrolytic Cell
Uses external electrical energy to cause a non-spontaneous chemical reaction.
Anode is positive; cathode is negative.
Example: Electrolysis of molten NaCl produces Na metal and Cl₂ gas.
5. Electrode Potential and Standard Electrode Potential
Electrode potential is the potential difference developed at the interface of electrode and
electrolyte.
Measured against standard hydrogen electrode (SHE), which is assigned 0 V.
Standard electrode potential (E°) is measured under standard conditions (1 M, 1 atm,
25°C).
6. Electrochemical Series
List of elements arranged by their standard electrode potentials.
Elements with more negative E° oxidize easily (strong reducing agents).
Elements with more positive E° are less reactive (strong oxidizing agents).
7. Nernst Equation
Calculates electrode potential under non-standard conditions:
E=E∘−RTnFlnQE = E^\circ - \frac{RT}{nF} \ln QE=E∘−nFRTlnQ
