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Summary - Unit 14 - Redox II (9CH0) £2.99
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  2.  
  3. PEARSON (PEARSON)
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  5. Chemistry 2015
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  7. Unit 14 - Redox II (9CH0)

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Summary - Unit 14 - Redox II (9CH0)
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A summary of topic 14, organised so the notes are easy to understand. The notes are on slides, so they can be printed out and used as revision cards or posters, for revision on the go. The notes cross-reference the specification so it is easy to see where each bit of information has come from. They...

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  • August 5, 2024
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Standard electrode potential 2 half cells = battery
The position of equilibrium and therefore the electrode potential depends - Need to choose a salt bridge
on factors such as temperature, pressure of gases & conc of reagents
∴ to compare the electrode potentials they are measured against a Can't use KCl because Cl- ions
standard present in solution
= the p.d. (voltage) produced between an electrode and a solution in a half
cell, relative to the standard hydrogen electrode [conditions = 1M, 298k, 1
atm]
Standard hydrogen electrode (SHE)
When the SHE is connected to another Can't use KCl because Cl- ions
half-cell, the standard electrode potential
= for measuring electrode potentials against of that half-cell can be read off a will react w/ Ag to form AgCl ppt
⤷ 1 mol/dm^3 of H+ voltmeter
⤷ 100kPa or 1atm pressure of H2
⤷ inert platinum electrode - Salt bridge completes a circuit
⤷ at 298k H+(aq) + e- - Allows ions to flow between 2 half calls
→ 1/2H2(g) - A HIGH RESISTANCE voltmeter is used to measure voltage no
(acid) current flows and the maximum potential difference is achieved
Half cells What does E° (E cell) mean?
= a piece of metal dipped into a solution of its own ions, an eqm is set up E° = a measure of the tendency
of a chemical reaction to occur
spontaneously in an
electrochemical cell
The further the eqm lies on More +ve = ↑ reactive
the left, the ↑ tendency of the metal to lose its electrons
More reactive = more readily forms ions
Types of half cells
1. Metal electrode + metal 2. Metal ion + metal ion:
ions:




3. Non-metal electrode
& non-metal ion:

, Electrochemical cells Feasibility
= produces electricity from chemical reaction - 2 half cells E° values = how easily it can be oxidised or reduced
connected together The more +ve the value = easier to reduce the species on the left of
E.g. Zn & Cu the half-equation
e- flow from more -ve side ⤷The reaction will tend to the forward direction
to more +ve side The less +ve the value, the easier it is to oxidise the species on the
right of the half-equation
Zn → Cu
⤷The reaction will tend to proceed in the backward direction
Feasible when the E° is positive
Entropy Change & ln K
Cell potential is related to entropy (S) & the eqm constant (K)
A larger cell potential = bigger change in total entropy
Because eqm lies on the LHS for Zn: Therefore, we can say that cell potential is directly proportional
to total entropy change
Expressing electrochemical cells: E° ∝ ΔStotal
Electrode | ion(s) || ion(s) | electrode SHE = always on LHS The use of two equations for Gibbs free energy change can also show
- More -ve electrode goes 1st that:
- || = salt bridge E° ∝ lnK
- | = phase change ΔG = –nF E°
- , = same phase ΔG = –RT ln K
half cell with the greatest negative potential is Limitations of Standard Electrode Potential Predictions
written on the left of the salt bridge, so 1. As standard electrode potentials are measured using solutions,
E°cell = E°right – E°left we have to consider the le Châtelier's effect on concentration
Examples: using non-standard conditions
- Zn & Cu If conc ↑on one side of the equation, the eqm will shift the the other
Zn(s) | Zn^2+(aq) || Cu^2+(aq) | Cu(s) side.
- H & Cr If this is towards the reduction reaction (forward), the electrode
Pt(s) | H2(g) | H+(aq) || Cr2O7^2- + H+ , Cr^3+ | Pt(s) potential will be more +ve & more feasible
If this is towards the oxidation reaction (backwards), the electrode
Calculating E°
potential will be less +ve & less feasible
E°cell = (more +ve E°) - (less +ve E°)
2. Reaction kinetics can also affect the prediction
The reaction may have a high Ea which inhibits the reaction. A
Feasible = +ve
reaction may be thermodynamically feasible but does not occur as the
number activation energy is so high which means that the reactants are
Not feasible = -ve kinetically stable
3. Many redox reactions are not aqueous

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