redox/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 redox reaction to occur

disproportionation reaction

redox reaction whereby an element undergoes oxidation and reduction simultaneously

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 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⁺ (eg HCl) 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 (from standard 1.00M to 0.001M), by le chatelier’s principle, POE shifts left, extent of forward reaction decreases

E(Mg2+/Mg) becomes more negative

E_^cell = E^θ(Cu²⁺/Cu) – more negative 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. rate of forward reaction = rate of backward reaction.

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)

why is potassium manganate (VII) not a primary standard?

• difficult to obtain pure because it is reduced by substances in distilled water to form maganese (IV) dioxide

  • presence of maganese (IV) dioxide further catalyses auto-decomposition of the potassium manganate solution: 4MnO4- + 2H2O → 4MnO2 + 3O2 + 4OH-

  • manganate (VII) is inherently unstable in the presence of maganese (II) ions: 2MnO4- + 3Mn2+ + H2O → 5MnO2 + 4H+

• reaction is slow in acid but fast in neutral solution

sodium thiosulfate and iodine colour change

• added to iodine → brown iodine fades

• near endpoint, starch solution added → blue-black colour produced because starch-iodine complex formed

• with continued addition of sodium thiosulfate, blue-black colour disappears because iodine is reduced by sodium thiosulfate

  • I2 + 2S2O3 2- → 2I- + S4O6 2-

solution of FeSO4 titrated with acidified KMnO4 colour change

• Before endpoint:

  • All purple KMnO₄ added is immediately reduced to colourless Mn2+.

  • The solution turns from green (Fe²⁺) → yellow (Fe³⁺).

• At the endpoint:

  • All Fe²⁺ is oxidized to Fe³⁺.

  • The first drop of KMnO₄ that remains unreacted imparts a pale pink colour to the yellow solution: observed as orange

green → yellow → orange (yellow + pink)

balanced equation for solution of FeSO4 titrated with acidified KMnO4

5Fe2+ + MnO4- + 8H+ → 5Fe3+ + Mn2+ + 4H2O

activity series

• ranks metals according to the ease with which they undergo oxidation to form cations

• gives information of which single replacement reactions occur spontaneously, but not how fast they occur

reactions of metals with water

• metals above H in reactivity series displace H+ ions from cold water to form hydrogen gas and metal hydroxide: 2M (s) + 2H2O (l) → 2M+ (aq) + OH- (aq) + H2 (g)

• metals slightly lower but still above H will displace H+ ions from steam to form hydrogen gas and metal oxide: Mg (s) + H2O (g) → MgO (s) + H2 (g)

• metals below hydrogen do not undergo reaction

reactions of metals with acids

• metals above H in reactivity series will displace H+ ions from dilute acid to form hydrogen gas and metal salt: Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g)

• anions of some conc acids can act as oxidising agents for metals below H: Cu (s) + 4HNO3 (aq) → Cu(NO3)2 (aq) + 2NO2 (g) + 2H2O (l)

displacement reactions of metals with other cations

• metal: itself oxidised → reducing agent: higher in reactivity series, stronger RA since readily lose electrons

• metal cations: itself reduced → oxidising agent: lower in reactivity series, stronger OA since readily gain electrons

• metal (s) displaces ions of metal below it in reactivity series to form its own ions

  • solid displace ions in aq solution

  • solid displace ions in solid [thermite reaction: heat together, exothermic reaction, forms solid compound and molten metal]: Fe2O3 (s) + 2Al (s) → 2Fe (l) + Al2O3 (s)

reactions of metals

• in water

  • much higher than H: reacts with cold water → forms hydrogen gas and aqueous metal hydroxide

  • slightly higher than H: reacts with steam → forms hydrogen gas and solid metal oxide

• in acid

  • higher than H: forms hydrogen gas and salt

  • lower than H: oxidised by anion of HNO3 acid → forms NO2 gas, water and the aqueous ionic compound

• displacement

  • displace aqueous ions → forms solid metal and aqueous ion

  • displace solid ions [thermite reaction] → forms molten metal and solid compound

electrolysis

process of passing a direct current of electricity through an electrolyte to cause a non-spontaneous chemical reaction

electrolyte

molten compounds or aqueous solutions which can conduct electricity

voltaic vs electrolytic cell

• voltaic 2 redox reactions taking place simultaneously in separate locations, transfer of electrons through wire and transfer of ions through salt bridge VS electrolytic reactants are not separated

• voltaic is spontaneous redox reaction producing electricity VS electrolytic is supplying electricity so that non-spontaneous redox reaction occurs

factors determining which ion is preferentially discharged at each electrode

• standard electrode potentials of ions

  • cathode (reduction, forward reaction) → more positive preferred

  • anode (oxidation, backward reaction) → more negative/less positive preferred

• relative concentration: only when difference in E is small!

  • by LCP, conc of Cl- (on RHS of eqn) increase, POE shifts left, favour backward reaction (oxidation) more → E becomes more negative/less positive → if exceeds E of H2O, Cl- higher tendency for oxidation, preferentially discharged

  • higher concentration, higher tendency to be reduced/oxidised

• nature of electrodes (for anode only, since cathode undergoes reduction and metals undergo oxidation)

  • inert will not participate in reaction

  • reactive anode (eg Cu) will be oxidised. E_Cu²⁺/Cu is the least positvie/most negative, higher tendency for oxidation, preferentially discharged → application: purification

  • note: graphite inert but oxygen produced at anode will oxidise the carbon anode, forming CO2/CO which contaminates the oxygen

faraday’s 1st law

mass of any substance liberated or deposited at anode electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.

m ∝ Q

faraday’s 2nd law

no. of faradays F required to discharge one mole of an ion at an electrode is equal to the number of charges on the ion

Q = I x t = n_e x F

n_e is the no. of moles of electrons. use this and mole ratio to calculate no. of mol of product/reactant.

factors affecting mass of products in electrolysis

• time t

• magnitude of direct current I

• charge on ion of element → affects mole ratio and thus n_e