Electrochemistry: Terms, Definitions, Tables, and Formulas (Chapter 20 Study Notes)

Electrochemistry and Redox Reactions

Lightning and Batteries

Electrochemistry is the study of chemical processes that cause electrons to move, which is the basis for batteries and lightning. These processes involve oxidation-reduction (redox) reactions, where electrons are transferred between substances.

• Oxidation: Loss of electrons by a substance.

• Reduction: Gain of electrons by a substance.

• Redox Reaction: A chemical reaction involving the transfer of electrons.

• Batteries: Devices that convert chemical energy into electrical energy via redox reactions.

Balancing Oxidation-Reduction Equations

Balancing redox reactions is essential for understanding electrochemical cells. The half-reaction method is commonly used:

1. Assign oxidation states to all atoms and identify the substance being oxidized and reduced.

2. Write separate half-reactions for oxidation and reduction.

3. Balance all elements except H and O in each half-reaction.

4. Balance O by adding H2O; balance H by adding H+.

5. Balance charge by adding electrons.

6. Multiply half-reactions by appropriate factors so electrons cancel.

7. Add the half-reactions and simplify.

Example: Balancing the reaction between Zn and Cu2+ in a battery.

Electrochemical Cells

Voltaic Cells: Generating Electricity from Spontaneous Chemical Reactions

Voltaic (or galvanic) cells use spontaneous redox reactions to generate electricity. They consist of two half-cells connected by a salt bridge.

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

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

• Salt Bridge: Allows ion flow to maintain charge balance.

• Electrolyte: Solution containing ions that conduct electricity.

Example: Daniell cell (Zn/Cu cell).

Cell Notation and Diagrams

Cell notation is a shorthand for representing electrochemical cells:

• Anode is written on the left, cathode on the right.

• Single vertical line (|) separates phases; double vertical line (||) represents the salt bridge.

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

Standard Electrode Potentials

Standard electrode potentials (Eo) measure the tendency of a half-cell to gain or lose electrons under standard conditions (1 M, 1 atm, 25°C).

• Standard Hydrogen Electrode (SHE): Reference electrode, Eo = 0 V.

• Reduction potential: Tendency to gain electrons.

• Oxidation potential: Tendency to lose electrons (negative of reduction potential).

Half-Cell	Standard Reduction Potential (Eo, V)
Zn2+ + 2e- → Zn(s)	-0.76
Cu2+ + 2e- → Cu(s)	+0.34
Standard Hydrogen Electrode	0.00

Cell Potential, Free Energy, and the Equilibrium Constant

The cell potential (Ecell) is related to the free energy change (ΔG) and the equilibrium constant (K) of the reaction:

• Cell potential (Ecell): The voltage produced by the cell.

• Free energy (ΔG): Determines spontaneity of the reaction.

• Equilibrium constant (K): Relates to the extent of the reaction.

Key equations:

• $\Delta G = -nFE_{cell}$

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

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

Example: Calculating cell potential for a Zn/Cu cell.

Batteries and Electrolysis

Batteries: Using Chemistry to Generate Electricity

Batteries are practical applications of electrochemical cells. Different types include:

• Lead-acid battery: Used in cars; consists of lead and lead dioxide electrodes in sulfuric acid.

• Nickel-cadmium (NiCd) battery: Uses solid cadmium anode and NiO(OH) cathode.

• Nickel-metal hydride (NiMH) battery: Similar to NiCd but uses a hydrogen-absorbing alloy instead of cadmium.

Comparison Table:

Battery Type	Anode	Cathode	Electrolyte	Application
Lead-acid	Pb	PbO2	H2SO4	Automobiles
NiCd	Cd	NiO(OH)	KOH	Rechargeable devices
NiMH	Hydrogen alloy	NiO(OH)	KOH	Rechargeable devices

Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity

Electrolysis uses electrical energy to drive nonspontaneous chemical reactions. It is used in processes such as electroplating and the production of chemicals.

• Electrolytic cell: Consumes electrical current to drive a nonspontaneous reaction.

• Anode: Positive electrode (oxidation occurs).

• Cathode: Negative electrode (reduction occurs).

Example: Electrolysis of molten NaCl to produce Na and Cl2.

Corrosion and Prevention

Corrosion: Undesirable Redox Reactions

Corrosion is the degradation of metals due to redox reactions with environmental agents, such as oxygen and water. Rusting of iron is a common example.

• Oxidation: Iron loses electrons and forms Fe2+ or Fe3+.

• Reduction: Oxygen gains electrons and forms water.

Prevention of rust:

1. Keep iron dry.

2. Coat iron with a substance impermeable to water.

3. Place a more reactive metal in contact with iron (sacrificial anode).

4. Coating iron with a metal that forms such protective films.

Key Calculations and Concepts

Essential Skills for Electrochemistry

• Balance redox equations using the half-reaction method.

• Assign oxidation numbers.

• Construct and interpret cell diagrams.

• Calculate cell potentials using standard reduction potentials.

• Relate cell potential to free energy and equilibrium constant.

• Predict spontaneity of reactions.

• Perform calculations involving AG, Ecell, and K for electrochemical reactions.

• Analyze electrolysis reactions and predict products at electrodes.

Example: Predicting whether a redox reaction is spontaneous using Ecell.

Additional info: These notes cover the main topics of Chapter 20: Electrochemistry, including batteries, balancing redox reactions, cell potentials, electrolysis, and corrosion. All equations are provided in LaTeX format for clarity.