Chapter 16: Electrochemistry

Vocabulary flashcards covering key terms from Chapter 16: Electrochemistry.

Electrochemistry

Branch of chemistry that studies transformations between chemical and electrical energy.

Redox reaction

Reaction in which oxidation and reduction occur as coupled processes; overall involves transfer of electrons.

Reduction half-reaction

Half of a redox reaction in which a species gains electrons.

Oxidation half-reaction

Half of a redox reaction in which a species loses electrons.

Anode

Electrode where oxidation takes place (loss of electrons).

Cathode

Electrode where reduction takes place (gain of electrons).

Salt bridge

Ion-conducting link between half-cells that balances charge and prevents buildup of charge.

Electrochemical cell

Device that converts chemical energy to electrical energy or vice versa.

Voltaic (galvanic) cell

Spontaneous electrochemical cell that converts chemical energy to electrical energy.

Electrolytic cell

Electrochemical cell in which electrical energy drives a nonspontaneous reaction.

Cell diagram

Symbolic representation of a cell showing electrodes, junction, and electrolytes.

Standard reduction potential (E°)

Potential of a half-reaction at standard state (25 C, 1 M, 1 bar) with all species in standard states.

Standard cell potential (E°cell)

Cell potential under standard conditions; E°cell = E°cathode − E°anode.

Faraday constant (F)

Constant: approximately 9.65 × 10^4 C per mole of electrons.

Nernst equation

Relates ΔG and E to reaction quotient Q: ΔG = ΔG° + RT ln Q and Ecell = E°cell − (RT/nF) ln Q.

Gibbs free energy change for a cell (ΔGcell)

ΔGcell = − n F Ecell; relates cell energy to emf and electron moles.

Standard Hydrogen Electrode (SHE)

Reference electrode with E° = 0.00 V for the H+/H2 couple.

Copper couple E° (Cu2+/Cu)

Standard reduction potential for Cu2+ + 2e− → Cu is +0.342 V.

Zinc couple E° (Zn2+/Zn)

Standard reduction potential for Zn2+ + 2e− → Zn is −0.762 V.

Zinc-Air battery

Battery using Zn and O2 to form ZnO; E°cell around 1.65 V.

Lead-acid battery

Rechargeable battery based on Pb/PbO2: PbSO4 and H2SO4; E°cell ≈ 2.041 V.

Corrosion

Oxidation of a metal by environmental species; promoted by water, electrolytes, and dissimilar metals.

Cathodic protection

Corrosion inhibition by making a structure the cathode or using sacrificial anodes.

Sacrificial anode

A more easily oxidized metal placed to corrode instead of the protected metal.

Electrolysis

Process of using electrical energy to drive a nonspontaneous redox reaction.

Electroplating

Depositing a metal onto a surface via electrolytic process.

Fuel cell

Voltaic cell that uses a continuous fuel supply to produce electricity; net reaction often H2 + O2 → H2O.

Net fuel cell reaction and E°

Common net reaction: 2 H2 + O2 → 2 H2O with E° ≈ 1.229 V.

Nickel–metal hydride battery (NiMH)

Battery using NiO(OH)/Ni(OH)2 electrodes and H2O/OH− chemistry.

Lithium-ion battery

Rechargeable battery with Li+ intercalation; cathode typically cobalt oxide, anode graphite.

Equilibrium condition for Ecell and K

At equilibrium Ecell = 0 and Q = K; log K = n E°cell / 0.0592 at 298 K.

Electrical energy units

Current measured in amperes (A); 1 A = 1 C/s; Work = charge × voltage; 1 W = 1 J/s.

Discharge of a lead-acid battery

Process where Ecell decreases as the battery provides current; reactions involve Pb and PbO2 forming PbSO4.