Electrode potentials

Question 1

Draw electrochemical cell

zinc and copper

Answer 1

Question 2

Electrochemical cells

description

Answer 2

• A cell has two half–cells
• The two half cells have to be connected with a salt bridge
• Simple half cells will consist of a metal (acts an electrode) and a solution of a compound containing that metal (eg Cu and CuSO4 )
• These two half cells will produce a small voltage if connected into a circuit. (i.e. become a Battery or cell).

Question 3

Why does a voltage form?

zinc and copper cell

Answer 3

When connected together the zinc half-cell has more of a tendency to oxidize to the Zn2+ ion and release electrons than the copper half-cell. (Zn Zn2+ + 2e-)

More electrons will, therefore, build upon the zinc electrode than the copper electrode. A potential difference is created between the two electrodes. The zinc strip is the negative terminal and the copper strip is the positive terminal. This potential difference is measured with a high resistance voltmeter, and is given the symbol E. The E for the above cell is E= +1.1V

Question 4

Why use a high resistance voltmeter?

Answer 4

The voltmeter needs to be of very high resistance to stop the current from flowing in the circuit. In this state it is possible to measure the maximum possible potential difference (E). The reactions will not be occurring because the very high resistance voltmeter stops the current from flowing.

Question 5

Salt Bridge

Answer 5

The salt bridge is used to connect up the circuit. The free-moving ions conduct the charge

A salt bridge is usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate

The salt should be unreactive with the electrodes and electrode solutions. E.g. potassium chloride would not be suitable for copper systems because chloride ions can form complexes with copper ions. A wire is not used because the metal wire would set up its own electrode system with the solutions

Question 6

What happens if the current is allowed to flow?

Answer 6

If the voltmeter is removed and replaced with a bulb or if the circuit is short-circuited, a current flows

The reactions will then occur separately at each electrode.

The voltage will fall to zero as the reactants are used up.

The most positive electrode will always undergo reduction.

Cu2+ (aq) + 2e- —- Cu(s)

(positive as electrons are used up)

The most negative electrode will always undergo oxidation.

Zn(s) —- Zn2+ (aq) + 2e-

(negative as electrons are given off)

Question 7

Cell Diagrams

Answer 7

• The solid vertical line represents the boundary between phases e.g. solid (electrode) and solution (electrolyte)
• The double line represents the salt bridge between the two half cells
• the voltage produced is indicated
• the more positive half cell is written on the right if possible (but this is not essential)

Question 8

Systems that do not include metals.

what if…

Answer 8

If a system does not include a metal that can act as an electrode, then a platinum electrode must be used and included in the cell diagram. It provides a conducting surface for electron transfer. A platinum electrode is used because it is unreactive and can conduct electricity

Question 9

Measuring the electrode potential of a cell

Answer 9

• It is not possible to measure the absolute potential of a half electrode on its own. It is only possible to measure the potential difference between two electrodes.
• To measure it, it has to be connected to another half-cell of known potential, and the potential difference between the two half-cells measured.
• by convention we can assign a relative potential to each electrode by linking it to a reference electrode (hydrogen electrode), which is given a potential of zero Volts

Question 10

Draw SHE

Answer 10

Question 11

Describe SHE

Answer 11

The potential of all electrodes are measured by comparing their potential to that of the standard hydrogen electrode. The standard hydrogen electrode (SHE) is assigned the potential of 0 volts.

The hydrogen electrode equilibrium is:

H2 (g) —–2H+ (aq) + 2e

In a cell diagram the hydrogen electrode is represented by:

Pt |H2 (g) | H+ (aq)

Question 12

Components of a standard hydrogen electrode

4

Answer 12

1. Hydrogen gas at a pressure of 100kPa
2. A solution containing the hydrogen ion at 1.0 mol dm-3 (the solution is usually 1 mol dm-3 HCl)
3. The temperature at 298K
4. Platinum electrode

Question 13

Secondary standards

Answer 13

The standard hydrogen electrode is difficult to use, so often a different standard is used which is easier to use. These other standards are themselves calibrated against the SHE. This is known as using a secondary standard - i.e. a standard electrode that has been calibrated against the primary standard.

Question 14

Standard Electrode Potentials

Answer 14

When an electrode system is connected to the hydrogen electrode system, and standard conditions apply the potential difference measured is called the standard electrode potential,

Question 15

Calculating the EMF of a cell

Answer 15

Ecell = Ered – Eox

Question 16

Using series of standard electrode potentials

Answer 16

Question 17

Example 1

Use electrode data to explain why fluorine reacts with water. Write an equation for the reaction that occurs.

Answer 17

First, apply the idea that more positive Eo will reduce (go forward) and more negative Eo will oxidize (go backward)

Explanation to write As Eo F2 /F- > Eo O2 /H2O, F2 will oxidize H2O to O2

Can also work out Ecell and quote it as part of your answer Ecell = Ered - Eox = 2.87-1.23 =1.64V

Question 18

Use the half-equations to explain in terms of oxidation states what happens to hydrogen peroxide when it is reduced

Answer 18

As Eo H2O2 /H2O > Eo O2 /H2O2 , H2O2 disproportionates from -1 oxidation state to 0 in O2 and -2 in H2O

Question 19

Effect of conditions on Cell voltage Ecell

Answer 19

Use La chateliers principle

Question 20

Effect of concentration on Ecell

Answer 20

Looking at cell reactions is a straight forward application of le Chatelier. So increasing concentration of ‘reactants’ would increase Ecell and decreasing them would cause Ecell to decrease

Question 21

Effect of temperature on Ecell

Answer 21

Most cells are exothermic in the spontaneous direction so applying Le Chatelier to a temperature rise to these would result in a decrease in Ecell because the equilibrium reactions would shift backwards

If the Ecell positive it indicates a reaction might occur. There is still a possibility, however, that the reaction will not occur or will occur so slowly that effectively it does not happen. If the reaction has a high activation energy the reaction will not occur

Question 22

Cells are non-rechargeable when

Answer 22

the reactions that occur with in them are non-reversible.

Question 23

Lithium-ion cell

Answer 23

Li+ + CoO2 + e- —–Li+ [CoO2] - E=+0.6V

Li+ + e- —- Li E=-3.0V

Lithium-ion cells are used to power cameras and mobile phones

Overall discharge Li + CoO2 ——-LiCoO2 E=3.6V

Li | Li+ || Li+ , CoO2 | LiCoO2 | Pt

The overall reaction would be reversed in the recharging state

The reagents in the cell are absorbed onto powdered graphite that acts as a support medium. The support medium allows the ions to react in the absence of a solvent such as water. Water would not be good as a solvent as it would react with the lithium metal

Question 24

Fuel cells

Answer 24

A fuel cell uses the energy from the reaction of a fuel with oxygen to create a voltage

Question 25

Hydrogen fuel cell (potassium hydroxide electrolyte)

Alkaline conditions

draw + describe

Answer 25

4e- + 4H2O —–2H2 +4OH- E=-0.83V

4e- + 2H2O +O2 —- 4OH- E=+0.4V

Overall reaction

2H2 + O2 —— 2H2O E=1.23V

Fuel cells will maintain a constant voltage over time because they are continuously fed with fresh O2 and H2 so maintaining constant concentration of reactants. This differs from ordinary cells where the voltage drops over time as the reactant concentrations drop

Higher temperatures are therefore used to increase rate but the reaction is exothermic so by applying le Chatelier would mean the E cell falls. A higher pressure can help counteract this.

Question 26

Advantages of Fuel cells

Answer 26

Advantages of Fuel cells over conventional petrol or diesel-powered vehicles

(i) less pollution and less CO2 . (Pure hydrogen emits only water whilst hydrogen-rich fuels produce only small amounts of air pollutants and CO2 ).
(ii) greater efficiency

Question 27

Limitations of hydrogen fuel cells

4

Answer 27

(i) expensive

Hydrogen is readily available by the electrolysis of water, but this is expensive. To be a green fuel the electricity needed would need to be produced from renewable resources

(ii) storing and transporting hydrogen, in terms of safety, the feasibility of a pressurized liquid and a limited life cycle of a solid ‘adsorber’ or ‘absorber’
(iii) limited lifetime (requiring regular replacement and disposal) and high production costs,
(iv) use of toxic chemicals in their production

Question 28

Hydrogen can be stored in fuel cells as

3

Answer 28

(i) as a liquid under pressure,
(ii) adsorbed on the surface of solid material
(iii) absorbed within a solid material

Question 29

Ethanol fuel cells

Answer 29

Ethanol fuel cells have also been developed. Compared to hydrogen fuel cells they have certain advantages including. Ethanol can be made from renewable sources in a carbon-neutral way. Raw materials to produce ethanol by fermentation are abundant. Ethanol is less explosive and easier to store than hydrogen. New petrol stations would not be required as ethanol is a liquid fuel.

Equation that occurs at oxygen electrode

4e- + 4H+ +O2 —–2H2O E=1.23V

Equation that occurs at ethanol electrode

C2H5OH + 3H2O → 2CO2 + 12H+ + 12e-

Overall equation

C2H5OH + 3O2 → 2CO2 + 3H2O