Electrochemistry: Study Notes for General Chemistry

Electrochemistry

Standard Reduction Potentials

Standard reduction potentials are a measure of the tendency of a chemical species to gain electrons and be reduced. These values are measured under standard conditions (25°C, 1 M, 1 atm) and are referenced against the standard hydrogen electrode (SHE), which is assigned a potential of 0.00 V.

• Reduction: Gain of electrons by a species.

• Oxidation: Loss of electrons by a species.

• Standard Reduction Potential (E0): Indicates how easily a species is reduced compared to SHE.

Half-Reaction	E0 (V)
F2 + 2e- → 2F-	+2.87
Ag+ + e- → Ag	+0.80
Cu2+ + 2e- → Cu	+0.34
Zn2+ + 2e- → Zn	-0.76
Li+ + e- → Li	-3.04

Example: The species with the lowest (most negative) E0 value is the strongest reducing agent and most likely to donate electrons.

Electrochemical Cells

An electrochemical cell is a device that converts chemical energy into electrical energy (or vice versa) via redox reactions. It consists of two half-cells connected by a conductive wire and a salt bridge.

• Half-Cell: Contains an electrode and an electrolyte where either oxidation or reduction occurs.

• Salt Bridge: Maintains electrical neutrality by allowing ion flow between half-cells.

Example: The cell with the largest positive Ecell will produce the largest quantity of electricity.

Galvanic (Voltaic) Cells

Galvanic cells are spontaneous electrochemical cells that generate electricity from spontaneous redox reactions. They are commonly used in batteries.

• Anode: Electrode where oxidation occurs (loss of electrons).

• Cathode: Electrode where reduction occurs (gain of electrons).

• Electrons flow from anode to cathode through the external circuit.

• The salt bridge allows ions to flow and maintain charge balance.

Example: In a cell with Mg and Cu electrodes, Mg is oxidized at the anode and Cu2+ is reduced at the cathode.

Galvanic Cell Electrodes and Electron Flow

• At the anode, mass decreases as metal atoms lose electrons and enter solution.

• At the cathode, mass increases as metal ions gain electrons and deposit as solid metal.

• Electrons always flow from anode to cathode.

Electrolytic Cells

Electrolytic cells use electrical energy to drive nonspontaneous chemical reactions (electrolysis). They are used for processes such as electroplating and decomposition of compounds.

• Anode: Positive electrode (oxidation occurs).

• Cathode: Negative electrode (reduction occurs).

• Requires an external power source.

Key Differences from Galvanic Cells:

• Electrolytic cells require energy input; galvanic cells produce energy.

• Electron flow is still from anode to cathode, but the anode is positive in electrolytic cells.

Cell Potential (Ecell)

The cell potential is the voltage difference between the two electrodes of an electrochemical cell.

• Standard Cell Potential (Ecell0): Calculated using standard reduction potentials:

$E_{cell}^{0} = E_{cathode}^{0} - E_{anode}^{0}$

If Ecell > 0, the reaction is spontaneous (galvanic cell). If Ecell < 0, the reaction is nonspontaneous (electrolytic cell).

The Nernst Equation

The Nernst equation allows calculation of cell potential under nonstandard conditions:

$E_{cell} = E_{cell}^{0} - \frac{0.0592}{n} \log Q$

• Q: Reaction quotient (ratio of product and reactant concentrations).

• n: Number of moles of electrons transferred.

Relationship Between Ecell, ΔG, and K

The cell potential, Gibbs free energy, and equilibrium constant are related as follows:

• $\Delta G^{0} = -nFE_{cell}^{0}$

• $\Delta G^{0} = -RT \ln K$

• $E_{cell}^{0} = \frac{0.0592}{n} \log K$

Where F is Faraday's constant (96,485 C/mol e-), R is the gas constant, and T is temperature in Kelvin.

Cell Notation

Cell notation is a shorthand way to represent electrochemical cells:

• Single vertical line (|) separates different phases.

• Double vertical line (||) represents the salt bridge.

• Anode (oxidation) is written on the left; cathode (reduction) on the right.

Example: Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s)

Electroplating and Electrochemical Stoichiometry

Electroplating uses electrical current to deposit a metal onto a surface. The amount of substance deposited or dissolved can be calculated using Faraday's laws:

• $Q = I \times t$

• $\text{mol e}^- = \frac{Q}{F}$

• $\text{mol M} = \frac{\text{mol e}^-}{n}$

Where Q is charge (Coulombs), I is current (Amperes), t is time (seconds), F is Faraday's constant, and n is the number of electrons per mole of metal.

Summary Table: Key Electrochemistry Relationships

Quantity	Formula	Units
Cell Potential	$E_{cell}^{0} = E_{cathode}^{0} - E_{anode}^{0}$	Volts (V)
Gibbs Free Energy	$\Delta G^{0} = -nFE_{cell}^{0}$	Joules (J)
Equilibrium Constant	$E_{cell}^{0} = \frac{0.0592}{n} \log K$	—
Nernst Equation	$E_{cell} = E_{cell}^{0} - \frac{0.0592}{n} \log Q$	Volts (V)
Charge	$Q = I \times t$	Coulombs (C)

Additional info:

• Practice problems and examples are included throughout to reinforce concepts and calculations.

• Tables and diagrams are used to summarize standard reduction potentials and cell components.