Chapter 20: Electrochemistry

Chemistry 1062: Principles of Chemistry II

Andy Aspaas, Instructor

Oxidation-Reduction reactions

• Oxidation-reduction (redox) reaction: transfer of electrons from one species to another

• H3O+ becomes simply H+ when dealing with redox reactions to simplify balancing

– (still the same species, just different notation)• Skeleton oxidation-reduction equation: involves only the

species being oxidized and reduced. – Write oxidation numbers above each species. – No spectator ions, no balancing

• Half reaction: shows only one oxidation OR one reduction– Most redox reactions are split into an oxidation half-

reaction and a reduction half-reaction– LEO, GER

Balancing redox equations in acidic solutions

• For each half reaction…

– Balance everything except H or O

– Balance O by adding H2O to one side

– Balance H by adding H+ to one side– Balance charge by adding e- to one side

• Multiply each half reaction by a factor so that the electrons cancel when the two half reactions are added together (e- cannot appear in the final equation)

• Add the reactions, cancel anything that appears on the left and right, and simplify the coefficients to the smallest integers

Practice balancing acidic redox reactions

• Balance I2(s) + NO3-(aq) IO3

-(aq) + NO2(g) in acidic solution

• Half reactions

• Cancel electrons

• Add half-reactions

• Simplify

Voltaic cells

• A voltaic cell consists of two half-cells

– Each half-cell contains a metal rod dipped in a solution containing that metal ion

– Anode: a species is being oxidized– Cathode: a species is being reduced

• Cell reaction: redox reaction for entire voltaic cell

• Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Zn(s) Zn2+(aq): oxidation half-reaction, anode

Cu2+(aq) Cu(s): reduction half-reaction, cathode

Cell notation

• Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

• Cell notation: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

– Anode || Cathode

• Write half reactions and cell reactions for the following cell:

Tl(s) | Tl+(s) || Sn2+(aq) | Sn(s)

emf, Standard Cell emf, Standard electrode potential

• Electromotive force, emf, Ecell = electrical pressure across the conductors of an electrochemical cell

– Unit: Volt, V– Measure of the driving force of a cell reaction

• Standard cell emf, Eocell = solutes are 1 M, gases are 1 atm,

temperature is 25 oC• Standard electrode potenital

– By convention, the standard hydrogen electrode has an emf of 0 V

– All reactions shown as reductions• Ecell = Ecathode – Eanode

• Ecell is positive for spontaneous reactions as written

Practice calculating Ecell

• Using standard potentials, calculate Ecell for Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

• Standard cell potentials are an intensive property

– Do not depend on quantity!– If you have to multiply a half-reaction to cancel

electrons, do not multiply the Eo for that half-reaction

Free energy and K from Ecell

• ΔGo = -nFEcell

n = moles of electrons transferred

F = Faraday’s constant, 96,500 C/mol e-

• This gives an answer in J, since 1 J = 1 C·V

• Convert to kJ since that’s what ΔGo is usually expressed in

• Ecell = (0.0592 / n) log K (Nernst equation)

Practice with ΔGo and K

• Calculate ΔGo and K for the following cell: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

• ΔGo = -nFEcell

• Ecell = (0.0592 / n) log K