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3.1.11 Electrode Potential (Alkaline hydrogen-oxygen fuel cell (used to…
3.1.11 Electrode Potential
Half cells:
when a rod of metal is dipped into a solution of its own ions; equilibrium forms
e.g. zinc half cell: Zn(s) ⇌ Zn²⁺(aq) + 2e⁻
two half cells joined to make an electrical cell
electrical potential (how easily electrons are released by the metal/ how good reducing agent it is) cannot be measured directly; can measure potential difference
connect two half cells with a salt bridge (filter paper soaked in a salt solution) and measure potential difference between them with a voltmeter
more negative reading on voltmeter indicates that the metal more easily loses its electrons
e.g. zinc and copper(redox reaction):
zinc dissolves
Zn(s) → Zn²⁺(aq) + 2e⁻
electrons flow through wire and combine with copper ions
in solution to form copper metal deposited on the copper
electrode
Cu²⁺(aq) + 2e⁻ → Cu(s)
overall reaction: (electrons cancel out)
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
representing cells:
vertical solid line: phase boundary e.g. between solid and solution
double vertical solid line: salt bridge
species with highest oxidation state is written next to salt bridge
E° cell value of right electrode must be written with a polarity (+ / -)
working out: E° cell = E° negative cell + E° positive cell
or: emf = E(R) - E(L)
if emf is positive, the reaction is feasible (but may be too slow to work irl)
signs of electrodes tell direction of redox reactions; electrons travel from negative to positive
if a redox system only involves metal ions but no metals (e.g. Fe²⁺/Fe³⁺, or gaseous ions) all relevant ions are put into a beaker with a platinum electrode to make electrical contact, connected to a SHE
e.g. Pt | H₂(g) | 2H+(aq) || Fe³⁺(aq) , Fe²⁺(aq) | Pt
Standard electrode potentials:
standard hydrogen electrode
potential difference of cell is defined as zero
if connected to another electrode, the measured voltage (emf) is the electrode potential of the cell
electrodes with negative reading when connected to SHE are better reducing agents, and better at releasing electrons
hydrogen gas bubbled into a solution of H⁺ ions; electrical contact is made via a piece of unreactive platinum coated with finely divided platinum black to increase SA as the hydrogen does not conduct on its own; under standard conditions
standard conditions:
298K
100kPa
1.00moldm⁻³ ion solution
changing the conditions will change the electrode potential of the cell
electrochemical series:
the more negative the E° value, the better the metal is at releasing electrons; better reducing agent
the more positive the E° value, the better the metal is at accepting electrons; better oxidising agent
equilibria are written as reduction reactions (electrons on left of arrow)
most negative values at the top
Non-rechargeable cells:
zinc/copper cells
contained liquids so not suitable for portable devices
provide electricity for old fashioned telegraphs
zinc/carbon cells
positive carbon electrode: 2NH₄⁺(aq) + 2e⁻ → 2NH₃(g) + H₂(g)
negative zinc rod: Zn(s) → Zn²⁺(aq) + 2e⁻
overall reaction: 2NH₄⁺(aq) + Zn(s) → 2NH₃(g) + H₂(g) + Zn²⁺(aq)
emf ~1.5V
as cell discharges, the zinc is used up and the walls of the zinc canister become thin and the paste leaks out
irreversible reactions
waste issues
Rechargeable batteries
lead-acid batteries
positive lead-dioxide-coated electrode: PbO₂(s) + 4H⁺(aq) + 2SO₄²⁻(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l)
negative lead electrode: Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻
overall equation: PbO₂(s) + 4H⁺(aq) + 2SO₄²⁻(aq) + Pb(s) → 2PbSO₄(s) + 2H₂O(l)
emf ~ 2V
car batteries: reactions are reversed as the battery is charged up and electrons flow in the reverse direction, driven by the car's generator
nickel/cadmium
Cd to Ni on discharge
Ni to Cd on recharge
negative cadmium electrode: Cd(OH)₂(s) + 2e- ⇌ Cd(s) + 2OH⁻(aq)
positive nickel electrode: NiO(OH)(s) + H₂O(l ) + e- ⇌ Ni(OH)₂(s) + OH⁻(aq)
overall equation: 2NiO(OH)(s) + Cd(s) + 2H₂O(l) ⇌ 2Ni(OH)₂(s) + Cd(OH)₂(s)
emf ~ +1.2V
more expensive but can be recharged up to 500 times
lithium ion
positive lithium cobalt oxide electrode: Li⁺ + CoO₂ + e⁻ → Li⁺[CoO₂]⁻
emf = +1V
negative carbon electrode: Li -→ Li⁺ + e⁻
emf = -3V
used in laptops, tablets, smartphones; light as lithium is the least dense letal
solid electrolyte so it doesn't leak
Alkaline hydrogen-oxygen fuel cell
used to generate an electrical current and do not need to be electrically recharged
positive electrode: O₂(g) + 2H₂O(l) +4e⁻ → 4OH⁻(aq)
emf = +0.4V
negative electrode: 2H₂(g) + 4OH⁻(aq) → 4H₂O(l) + 4e⁻
emf = -0.83V
overall: 2H₂(g) + O₂(g) → 2H₂O(l)
emf: 1.23V
two electrodes of a porous platinum based material, separated by a semi permeable membrane; electrolyte is NaOH solution
used in some cars and spacecraft; only by-product is H₂O which is not a pollutant; gives out three times as much energy as petrol of the same mass
hydrogen is sourced from crude oil (not renewable); hydrogen is very flammable; takes up a lot of space
Sammer Sheikh