BI1014 - redox reactions

oxidation

Loss of electrons (↑ oxidation number)

reduction

Gain of electrons (↓ oxidation number)

oxidising agent

- Gains electrons (is reduced)

- Causes oxidation

- Becomes more negative

reducing agent

- Loses electrons (is oxidised)

- Causes reduction

- Becomes more positive

general rules of oxidation numbers

1. Free elements = 0 (e.g. Na, Cl₂)

2. Monoatomic ions = their charge (e.g. Ca²⁺ = +2)

3. Molecular compounds:

- More electronegative element gets negative oxidation number

- Bonds between same element don't count

Total = 0 (or ion charge for polyatomic ions)

common assignments

- Hydrogen = +1

Exception: hydrides (e.g. LiAlH₄) = -1

- Oxygen = -2

Exception: F₂O = +2

group number rule

Max oxidation number = group number

Min oxidation number = group number − 8

anode

Reaction -> Oxidation

Charge -> Positive

Electron flow -> Out

Attracts -> Anions

cathode

Reaction -> Reduction

Charge -> Negative

Electron flow -> In

Attracts -> cations

standard electrode potentials

To calculate ΔE°:

- Identify half-reactions

- Place the more negative E° on top (to reverse)

- Add both E° values

If ΔE° > 0, the redox reaction is spontaneous

discharging

spontaneous redox (ΔE° > 0)

charging

reverse reaction with external voltage

nernst equation

R = 8.314 J/mol·K

T = temperature (K)

n = number of electrons

F = 96,485 C/mol (Faraday constant)

Q = reaction quotient

types of electrode

1st kind

2nd kind

pH electrode

1st kind

Metal electrode in solution of its own ions

2nd kind

Inert metal (e.g. Pt) in contact with redox pair

pH electrode

Needs calibration due to unknown cell constant