electrochemistry

electrochemical processes

redox reactions where energy released by spontaneous reaction is converted to electricity
OR electrical energy is used to cause a non-spontaneous reaction to occur

anode

electrode where oxidation occurs

cathode

electrode where reduction occurs

how is electricity generated through a spontaneous redox reaction?

• oxidation reaction and reduction reaction take place simultaneously in separate locations

• transfer of electrons through external metal wire

• transfer of ions through salt bridge

• reaction progresses, sets up constant flow of electrons, generates electricity

purpose of salt bridge

• complete the electrical circuit: allowing the movement of ions between half-cells

• maintain electrical neutrality: prevent build-up of positive charges at anode/build-up of negative charges at cathode → this would prevent cell from operating

types of half-cells

• ion-ion

  • platinum (inert) electrode submerged in solution that is mixture of two metal ions

• metal-ion

  • metal (reactive) electrode submerged in solution of metal ions

• gas-ion

  • gas bubbling over platinum (inert) electrode submerged in gaseous ion solution

standard reduction potential

aka standard electrode potential
the electromotive force between a half-cell when connected to standard hydrogen electrode under standard conditions (1 mol dm^-3, 298K, 1 bar)

by comparing half-cell against standard hydrogen electrode, it determines the relative extent of/tendency for the forward reduction reaction (see data booklet) under standard conditions

conditions for measuring standard reduction potential

• 298K

• 1 bar pressure for gases

• 1 mol dm^-3 for ions

• use platinum electrode if half-cell doesn’t involve metal

• salt bridge completes circuit

factors affecting electrode potential

• temperature and pressure

• nature of metal

• concentration of ions

• medium where reaction takes place (eg acidic conditions? → see redox table)

standard hydrogen electrode

• 1 bar hydrogen gas bubbling over platinum electrode immersed in 1.00 mol dm^-3 solution of H⁺ under standard conditions (298K, 1 bar)

• standard reduction potential = 0V because at equilibrium

IA9 Q4: what is the expected cell potential for a voltaic cell made from standard copper half-cell connected to magnesium half cell where concentration of Mg²⁺ ions was 0.001 mol dm^-3?

Mg²⁺ + 2e ⇌ Mg where E^θ(Mg²⁺/Mg) = -2.37V (for 1.00 mol dm^-3)

when concentration of Mg²⁺ decreases, by le chatelier’s principle, POE shifts left, extent of forward reaction decreases

E_^cell = E^θ(Cu²⁺/Cu) – E^θ(Mg²⁺/Mg) therefore E_^cell increases

the more positive the E^θ value,

• POE lies on right

• forward reaction (reduction) favoured

• oxidising agent (LHS) higher tendency to gain e^-, stronger oxidising agent/itself higher tendency to be reduced

the more negative the E^θ value,

• POE lies on left

• backward reaction (oxidation) favoured

• reducing agent (RHS) higher tendency to lose e^-, stronger reducing agent/itself higher tendency to be oxidised

standard cell potential

maximum potential difference between electrodes. the tendency of electrons to flow through the external circuit of a voltaic cell under standard conditions (1 mol dm^-3, 298K, 1 bar).

E^θ_cell = E^θ_{cathode/reduction} – E^θ_{anode/oxidation}

note: E^θ_cell is always positive for a battery / DO NOT CHANGE THE SIGNS OF ANY VALUES FROM DATA BOOKLET WHEN USING EQUATION

E^θ_cell > 0

ΔG < 0 → spontaneous forward redox reaction, non-spontaneous backward redox reaction

E^θ_cell < 0

ΔG > 0 → spontaneous backward redox reaction, non-spontaneous forward redox reaction

E^θ_cell = 0

ΔG = 0 → equilibrium

uses of E^θ_cell

predict feasibility of reaction (based on ΔG)

however, might not be accurate if:

• very high E_a needed for reaction → kinetically slow

• not standard conditions

uses of E^θ

• determine strength of OA/RA

• predict direction of electron flow (based on above)