5.2.3 Redox and electrode potentials

Definitions

Term	Definition
Oxidising agent	A reagent that accepts / takes in electrons / oxidises another species / is reduced
Reducing agent	A reagent that donates / gives out electrons / reduces another species / is oxidised
Standard electrode potential \(E^\ominus\)	The emf of a half-cell compared to a hydrogen electrode under standard conditions
Standard half cell	The pure metal in contact with a 1 molar solution of its ions at a temperature of 298 K

Acidic condition
- Add 1 \(H_2O\) on the side with less oxygen for every missing oxygen
- Add 2 \(H^+\) on the other side for every \(H_2O\) added
Alkaline condition
- Add 2 \(OH^-\) on the side with less oxygen for every missing oxygen
- Add 1 \(H_2O\) on the other side for every pair of \(OH^-\) added
Balance changes in oxidation numbers with electrons

Manganate reduced from \(Mn^{7+}\) to \(Mn^{2+}\)
\(KMnO_4\) is used as the oxidation agent and added to burette
The reducing agent is added to the conical flask
Excess dilute \(H_2SO_4\) added to provide \(H^+\) for reduction of \(MnO_4^-\)
Manganate in the burette is deep purple \(\rightarrow\) reacts and become colourless when added to the flask
Once all reducing agent has reacted there will be nothing to decolourise the manganate
End point = when permanent pink colour appears in flask
\(KMnO_4\) is self indicating so no indicator needed

\(Na_2S_2O_3\) (sodium thiosulfate) reduces iodine to iodide ions and forms \(Na_2S_4O_6\) (sodium tetrathionate)
- \(2S_2O_3^{2-}(aq) + I_2(aq) \rightarrow 2I^-(aq) + S_4O_6^{2-}(aq)\)
Standard solution of \(Na_2S_2O_3\) added to burette
An excess of iodide is added to conical flask as iodine is almost insoluble in water but very soluble in iodide solution
Solution of oxidising agent being analysed is added to conical flask with excess \(KI\)
Oxidising agent reacts with iodide ions to produce iodine \(\rightarrow\) turn solution yellow-brown
Iodine reduced back to \(I^-\) ions \(\rightarrow\) brown colour fades gradually
Indicator is starch: added towards the end when most of the iodine has reacted and gives a blue-black colour
End point: blue black disappears when iodine is used up and reaction mixture becomes colourless

More positive value: greater tendency to gain electrons + undergo reduction
More negative value: greater tendency to lose electrons + undergo oxidation

Contains the species present in a redox half equation (forward reaction shows reduction)
Consists of a piece of pure metal in contact with a solution of its ions
Electron can be transferred between the metal and its ions

Standard conditions must be used
- Metal ion solutions must have a concentration of \(1 \ mol \cdot dm^{-3}\)
- Gas half cells must be at 100kPa pressure + use inert electrode
- Temperature must be 298K
Half cell measured connected by wire with a standard hydrogen half cell
Two solutions connected with a salt bridge to complete the electrical circuit
- Typically contains a concentrated solution of an electrolyte that does not react with either solution (e.g. \(KNO_3\))
Connect the half-cells to voltmeter (high resistance)

Strongest oxidising agent = most positive \(E^\ominus\) + most likely to be reduced (will be on right of equation if written as standard form)
Strongest reducing agent = most negative \(E^\ominus\) + most likely to be oxidised
(\(X\) system is more negative than \(Y\) system \(\rightarrow\) \(X\) shifts left and \(Y\) shifts right / \(X\) reduces \(Y\))

Reactions can have very high activation energies \(\rightarrow\) very slow reaction rates
Actual conditions may not be standard
- Value of electrode potential will be different from standard value if concentration isn’t \(1 \ mol \cdot dm^{-3}\)
  - Increasing ion concentration: equilibrium shifts right, removes electrons, \(E^\ominus\) less negative
  - Reducing ion concentration: equilibrium shifts left, increases electrons in system, \(E^\ominus\) more negative
- Standard potentials are for aqueous equilibria: many reactions aren’t aqueous
There may be an alternative reaction that is more favourable

3 main types
Primary (non-rechargeable) cells
- Electrical energy produced by an irreversible reaction
- The reactants will be used up eventually and the reaction becomes too slow to create enough voltage
- Used for low current, long-storage devices e.g. clocks or smoke detectors
Secondary (rechargeable) cells
- Electrical energy produced by a reversible reaction
- The cell can be recharged by reversing cell polarities and forcing electrical energy through the cell
- Cell reaction is reversed during recharging

Can be regular shape or a flexible solid polymer
Charging / discharging: \(Li^+\) ions move between electrodes, electrons move through connecting wires
Negative electrode: graphite coated with lithium metal, \(Li \rightarrow Li^+ + e^-\)
Positive electrode: metal oxide (typically \(CoO_2\)), \(Li^+ + CoO_2 + e^- \rightarrow LiCoO_2\)

Use energy from reaction of fuel with oxygen to create voltage
A continuous supply of fuel and an oxidising agent (usually oxygen) to the cell
Fuel supplied to one electrode; oxidising agent to the other electrode
Can operate continuously if fuel and oxygen are continuously supplied - don’t have to be recharged

Produce only water as combustion product, no \(CO_2\) produced
Fuel cells using hydrogen rich fuels e.g. methanol are being developed
- Will produce carbon dioxide alongside water
- May produce more pollution but liquid fuel is much easier to store than gaseous ones
Hydrogen supplied to anode
- Oxidation in anode / negative electrode = oxidation
Oxygen supplied to cathode
- Reduction in cathode / positive electrode = reduction

Alkaline conditions
- Negative electrode / anode: \(2H_2(g) + 4OH^-(aq) \rightarrow 4H_2O(l) + 4e^-\)
- Positive electrode / cathode: \(O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq)\)
Acidic conditions
- Negative electrode / anode: \(2H_2(g) \rightarrow 4H^+(aq) + 4e^-\)
- Positive electrode / cathode: \(O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)\)