Understanding points
Reactivity 3.2.12—The hydrogen half-cell H+(aq) + e− ⇌ 12H2(g) is assigned a standard electrode
potential of zero by convention. It is used in the measurement of standard electrode potential, E⦵. (AHL)
Reactivity 3.2.13—Standard cell potential, E⦵cell, can be calculated from standard electrode
potentials. E⦵cell has a positive value for a spontaneous reaction. (AHL)
Reactivity 3.2.14—The equation ΔG⦵ = − nFE⦵ cell shows the relationship between standard
change in Gibbs energy and standard cell potential for a reaction. (AHL)
Reactivity 3.2.15—During electrolysis of aqueous solutions, competing reactions can occur at the anode and cathode, including the oxidation and reduction of water. (AHL)
Reactivity 3.2.16—Electroplating involves the electrolytic coating of an object with a metallic thin
Layer. (AHL)
Standard hydrogen electrode (SHE)
•
H+(aq) + e- → 12H2(g) E⊖ = 0 V
•
Provides reference point for electrode potential comparison of different half-cells
•
When the half cell contains a metal above hydrogen in the reactivity series → electrons flow from half cell to SHE → electrode potential is negative
•
When the half cell contains a metal below hydrogen in the reactivity series → electrons flow from SHE to half cell → electrode potential is positive
1.
Gaseous element is phase separated with water dissolved ion
2.
Inert Pt does not ionise
•
Electrode potential & redox strength comparison
•
Relative oxidising and reducing strengths:
◦
More -ve E⊖→ stronger reducing agent, likely to be oxidised
◦
More +ve E⊖ → stronger oxidising agent, likely to be reduced
Reactivity Series (↑) | Electrode Potential | Reducing Potential | Oxidising Potential |
K
Na
Ca
Mg
Al
Zn
Fe
Pb
Cu
Ag
Au
Pt | Positive
↑
↓
Negative | Strongest (... oxidised)
↑
↓
Weakest (unreactive) | Weak (... not reduced)
↑
↓
Weakest (unreactive) |
Electromotive force (EMF): electrical action produced by a non-electric source
•
The energy per unit electric charge imparted by an energy source
•
EMF in a voltaic cell results in the movement of e- from anode (-ve electrode) to the cathode (+ve electrode)
•
Standard = 1M conc./ 298K/ Pt electrode (gas)/ 100kPa (gas)
•
E⊖ = Standard electrode potential = Standard reduction potential
Calculating cell potential using standard electrode potentials
•
Cell potential = E⊖cell = E⊖cathode - E⊖anode at 298K, 1 atm, 1mol dm-3
•
Remember:
•
Do not change the +/- sign of E⊖
•
Do not multiply/scale up E⊖ value; it is not a per mole value
•
Likely to be +ve but not always
Determining the spontaneity/feasibility of a reaction using standard electrode potentials
1.
Write out half equations as they appear in the overall reaction
2.
Calculate E⊖cell of reaction
a.
E⊖cell = E⊖reduction - E⊖oxidation
3.
∆Go = -nFE⊖ where n= # mole of e-, and F= Faraday constant
a.
E⊖cell > 0 → ΔGº<0 → spontaneous reaction
b.
E⊖cell < 0 → ΔGº>0 → non-spontaneous reaction
Electrolysis of aqueous solutions
•
Water competes with the anions and cations at the anode and cathode
•
If the standard electrode potential of the cation is more negative than −0.83V, then water will be preferentially reduced at the cathode
◦
H2O(l) + e- → 12H2(g) + OH-(aq) E⦵=−0.83V
•
If the oxidation potential of the anion is more negative than −1.23V, then water will be preferentially oxidized at the anode
◦
H2O(l) → 12O2(g) + 2H+(aq) + 2e- E⦵= −1.23V
•
If one of the ions is much more concentrated than the other, it will be preferentially discharged
◦
Dilute NaCl → oxygen evolved
◦
Concentrated NaCl → Cl2 evolved
•
If the electrode is not inert, it may take part in the redox reaction
Factors affecting the amount of product formed in electrolysis
1.
Current (e- flow)
a.
Higher current means more e- availability for redox rxns → more product
2.
Time (duration)
a.
Longer duration of electrolysis means more e- flow overall for redox rxns → more product
3.
Charge (z) of ions
a.
Higher charge ions require more e- → less product formed
Electroplating
•
Coating one metal with a layer of another metal
•
The object to be electroplated is placed at the cathode and placed in a soln with the ions of the metal to plate the object, with the metal to plate the object as the anode
•
If the anode is made of impure metal and the cathode of pure metal, this process can be used to purify / refine impure metal
◦
e.g. pure copper has lower electrical resistance
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