Redox Equilibria

Question 1

How do electrochemical cells work?
Zn + Cu as an example.

Answer 1

• Electrochemical cells are made from two different metals dipped in salt solutions (e.g. zinc into zinc sulfate) of their own ions (setting up an equilibrium) and connected by a wire; the external circuit.
• It is a redox process; one metal is being oxidised , the other reduced.
• Zinc is oxidised more easily than copper; in the half-cell on the left, zinc (from the zinc electrode) is oxidised to form Zn2+ ions, releasing electrons into the external circuit.
• In the other half-cell, the same number of electrons are taken from the external circuit, reducing the Cu2+ ions to Cu atoms.
• Electrons frow through the wire from the most reactive metal (most easily oxidised; loses electrons easiest) to the least.
• A voltmeter can measure the potential difference on the external circuit; the e.m.f., E_cell.
• The following two half reactions occur:

Zn(s) → Zn2+(aq) + 2e-
Cu2+(aq) + 2e- → Cu(s)

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

• Half reactions are usually reversible, but in this direction as zinc loses its electrons more readily than copper; it is a better reducing agent; as shown with a more negative electrode potential; zinc is oxidised, whilst copper is reduced (positive electrode potential).
• A metal that’s easily oxidsed has a very negative electrode potential; one that’s harder to oxidise has a less negative/positive electrode potenial.

1. Zinc dissolves to form Zn2+(aq), increasing the conc. of Zn2+(aq).
2. The electrons flow through the wire to the copper rod where they would combine with Cu2+(aq) ions (from the copper sulfate solution), depositing fresh copper on the rod and decreasing Cu2+ conc.

Question 2

What is the conventional representation of drawing electrochemical cells?
Zn = -0.76
Cu = +0.34

What is the cell potential/e.m.f, E_cell?

Answer 2

• A vertical solid line indicates a phase boundary; between a solid and a solution.
  |
• A double vertical line shows a salt bridge. ||
• The hal-cell with the more negative potential goes on the left; Zn(s).
• The oxidised forms go in the centre like intermediaries; Zn2+(aq) & Cu2+(aq).
• Most positive (weakest reducing agent) goes on the right; Cu(s).

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

• E_cell = (E^o_RHS - E^o_LHS)
  Ecell = Eoright - Eoleft
  Ecell = +0.34 - (-0.76) = 1.10V
• The e.m.f. (Ecell) is always positive as the more negative E^o value is being subtracted from the more positive E^o value.
• The individual cell potential may be negative.

Question 3

How are cell used to measure electode potentials?
(standard hydrogen electrode)

What are standard conditions?

Answer 3

• Measure the electrode potential of a half-cell against a standard hydrogen electrode.
• **The standard electrode potential of a half-cell is the voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode.

Standard conditions**
: 298K, 100kPa, 1.00 mol dm^-3 solution of ions.

• The standard hydrogen electrode is always shown on the left; doesn’t matter if the other half-cell has a more positive value.
• The standard hydrogen electrode half-cell has an electrode potential of 0.00V.
• Measured e.m.f. = E^o_right - E^o_left
• Eoleft (hydrogen electrode) = 0.00V thus the measured value is the electrode potenial of that cell.
• Voltage reading = E^o_right

Question 4

What is the importance of the conditions when measuring an electrode potential?

Answer 4

• Conditions affect the value of the electrode potential; half-cell reactions are reversible, thus the equilibrium position is affected by changes in temperature, pressure and concentration.
• Changing the equilibrium changes the cell potential.
• Standard conditions are henceforth used.

Question 5

How can you predict if a reaciton will occur?
Will the following occur?
Fe3+(aq) + Cl-(aq) → Fe2+(aq) + 1/2Cl2(aq)

Fe = +0.77
Cl = +1.36

Answer 5

• Write both half-equations down, with the more negative standard electrode on top.
• Oxidised state goes on the left, reduced state on the right.
• The electrons should flow from the more negative electrode to the less negative; know which way around your half-equations go.
• Draw anti-clockwise arrows; shows direction of half-equation.
• Combine half-equations
• If combined equation is the same; feasible. If not, it is not feasible.

Question 6

What are electrochemical cells used for?

Answer 6

• Can be used as a commercial source of electrical energy; used as batteries, some rechargeable, some not.

Question 7

What types of electrochemical cells are used commercially?

Answer 7

Non-rechargeable (irreversible)

• E.g. dry alkaline battery; used for items that don’t use a lot of power or are only used for short periods of time.
• Half-equations have non-reversible arrows as not practical to reverse them in a battery; they can be made to run backwards under the right conditions, but trying to do this may result in leakage/explosion.
• Also cannot be recharged as the ammonium ions produce hydrogen gas which escapes from the battery; ammonium ions cannot be reformed.

Rechargeable (reversible)

• Two types; NiCad (nickel-cadium) and Li-ion (lithium ion).
• To recharge these, a current is supplied to force electrons to flow in the opposite direction around the circuit and reverse the reactions.
• Possible as none of the substances in a rechargeable battery escape or are used up.

Fuel cells

• In most cells the chemicals that generate the electricity are contained in the electrodes and the electrolyte that form the cell; in a fuel cell, the chemicals are stored separately outside the cell and fed in when the electricity is required; e.g. hydrogen-oxygen fuel cell.

Question 8

Outline the electrode reactions of a hydrogen-oxygen fuel cell and appreciate that a fuel cell does not need to be electrically recharged.

Answer 8

• Hydrogen and oxygen gases are fed into two separate platinum-containing electrodes.
• The electrodes are separated by an ion-exchange membrane that allows protons (H+ ions) to pass through it, but stops e^-.
• Hydrogen is fed to the negative electrode:

2H₂ → 4H⁺ + 4e^-

• The electrons flow from the negative electrode through an external circuit to the postive electrode.
• The H+ ions pass through the ion-exchange membrane towards the positive electrode.
• Oxygen is fed to the positive electrode:

O₂ + H⁺ +4e^- → 2H₂O

• The overall effect is that H₂ and O₂ react to make water:

2H₂ + O₂ → 2H₂O

Question 9

What are the advantages and disadvantages to society with the use of these cells?

(rechargeable, non-rechargeable, fuel cells)

Answer 9

Non-rechargeable + Rechargeable cells

1. Cost: non-rechargeable batteries are cheaper than rechargeables, however they have to be replaced every time they run out, so rechargeables are more cost-effective in the long run.
2. Lifetime: a non-rechargeable battery works longer than a rechargeable battery; but once a rechargeable battery rusn out, you can recharge it, whereas non-rechargeables have to be diposed of.
3. Power: non-rechargeable batteries can’t supply as much power as rechargeables, so are no use in devices that use a lot of power (mobile, laptop etc.).
4. Use of resources and waste: more non-rechargeable batteries are produced as they can only be used once, which uses more resources and means they create more waste then rechargeables. Both types can be recycled and the metals recovered to be used again, but often they are binned and landfill’d.
5. Toxicity: non-rechargeable batteries are less likely to contain toxic metals lead and cadium (though they may contain mercury), so are less hazardous in landfill in the event of a leak and pollutes water sources.

Fuel Cells

• Main advantage of fuel cells over batteries is that they don’t need electrical recharging; as long as hydrogen and oxygen are supplied, the cell wil continue to produce electricity.
• The only waste product is water, so no toxic chemicals to dispose of, no CO2 emissions directly from cell.
• However you need energy to produce a supply of oxygen and hydrogen; produced from the electrolysis of water i.e. reusing the waste product of the fuel cell, but this requires copious electricity; normally gathered by burning fossil fuels.
• Thus whole process isn’t usally carbon neutral.
• Hydrogen is also highly flammable and prone to explosion; needs to be handled carefully when stored or transported.