(1) Electrode Potentials

Question 1

What is an electrode/half cell?

Answer 1

A strip of metal dipped into a solution of its own ions.

An equilibrium is set up

Question 2

How can 2 half cells be connected?

Answer 2

• two metal rods are connected with wires and a high resistance voltmeter
• the two beakers of electrolyte are connected with a salt bridge to complete the circuit.
• salt bridge is soaked in potassium nitrate solution

Question 3

Why is KNO3 a suitable solution for a salt bridge?

What is the purpose of a salt bridge?

Answer 3

KNO3 is unreactive with the electrodes AND the ions are free to move

To complete the circuit

Question 4

What features must a solution have to be used as a salt bridge?

Answer 4

• must not react with chemicals in either solution

- ions free to move so conducts electricity

Question 5

What does the voltmeter do?

Answer 5

Prevents electrons flowing -so enables the voltage to be measured.

-otherwise electrons flow from left electrode to right (most reactive metal to the least reactive metal)

Question 6

How can you measure the current in a cell?

Answer 6

Voltmeter is replaced with an ammeter = electrons can flow and a current is produced

Question 7

Why might the current produced by a cell fall to zero after some time?

Answer 7

All the reactants are used up.

Question 8

What will happen to a cell once the reactants are used up?

Answer 8

stops working OR starts to leak

Question 9

Draw a diagram to represent a cell with copper and zinc electrodes

Write half equations for each electrode

Answer 9

Zn(s) –> Zn2+(aq) + 2e-

Cu2+(aq) + 2e- –> Cu(s)

Overall:
Zn + Cu2+ –> Zn2+ + Cu

Question 10

Zn + Cu2+ –> Zn2+ + Cu

This reaction would continue generating a current until what?

Answer 10

• The solid Zinc rod completely reacted.

- All the Cu2+ ions in solution were used up

Question 11

Key facts about electrodes?

Answer 11

~ left electrode is always the -ve electrode as e-s are produced there

~ oxidation always occurs at the -ve electrode

~ right electrode is always the +ve electrode as e-s are used up there

~ reduction always occurs at the +ve electrode

~ e-s flow from the -ve to +ve electrode

Question 12

When do we need to use platinum electrodes?

Answer 12

When there is no solid metal in the reaction, such as when there are metal ions of two different charges in the same solution.

-If this is the case a metal rod made from another unreactive metal is needed to connect the circuit (Platinum).

Question 13

Why is platinum a suitable electrode?

Answer 13

Pt is unreactive AND conducts electricity

Question 14

Why do we use the Standard Hydrogen Electrode?

Answer 14

To compare single electrodes with one another, this half cell is used as the standard electrode.

Question 15

Key Features of the Standard Hydrogen Electrode?

Answer 15

• H2 gas is pumped in at 100kPa
• Electrolyte contains [H+] ions of 1 moldm-3 (usually in HCl)
• Platinum electrode
• Temp of 298K

Question 16

What is the voltage of the Standard Hydrogen Electrode?

Answer 16

ZERO.

Question 17

All the equations shown in the electrochemical series are…

Answer 17

Reduction equations

Question 18

The standard electrode potential of Cu2+ is 0.37V.

Why might the electrode potential of a copper cell not be 0.37V?

Answer 18

If the concentration of the CuSO4 solution is not 1 mol dm-3.

Question 19

Write the conventional cell representation of:

Zn(s) + Cu2+ (aq) —> Cu(s) + Zn2+ (aq)

Answer 19

ROOR

Reduced Oxidised Oxidised Reduced

• Vertical solid lines indicate a phase boundary (i.e. between the solid and aqueous phases).
• A double vertical line in the middle represents the salt bridge.
• Order of electrodes is important too, start from left

Zn(s) │ Zn2+(aq) ‖ Cu2+(aq) │ Cu(s)

Question 20

Conventional cell representation of the Standard Hydrogen Electrode?

Answer 20

Pt(s)│H2(g) │H+(aq) ‖

Question 21

What is the conventional cell representation of a solution containing Fe2+ and Fe3+ ions?

The equilibrium taking place between the Pt and the solution is:

Fe2+ —> Fe3+ + e-

Answer 21

‖ Fe3+(aq), Fe2+(aq) │Pt(s)

Question 22

How can you work out the strongest oxidising agent and weakest reducing agent?

Answer 22

SOWR

Strongest oxidising agent + Weakest reducing agent = Most positive electrode potential

Most negative = weakest oxidising agent + strongest reducing agent

Question 23

How can you work out the E.M.F?

Work it out for:

Li+ = -3.03

Ag+ = +0.80

Answer 23

Most positive - Least positive

Ecell = (+0.80) - (-3.03)

Ecell = + 3.83V

Question 24

What is Cell Discharge?

Answer 24

Cell discharge

When the Ecell value is positive, this means the reaction is feasible and the cell discharges – this means it produces a current.

e.g.

Ag+(aq) + Li(s) <=> Ag(s) + Li+(s)

Question 25

What is Cell Recharge?

Answer 25

Cell recharge

If a reaction is reversible, the cell can be recharged by plugging it into the mains. The reverse reaction will occur when the cell is recharged.

e.g.

Ag(s) + Li+(s) <=> Ag+(aq) + Li(s)

Question 26

Give an environmental advantage and disadvantage of using rechargeable cells.

Answer 26

Advantage: Metals are reused

Disadvantage: Mains electricity is used to recharge, which may come from combusting fossil fuels, which releases CO2(g).

Question 27

What happens to electrode potential values if equilibria shifts?

Answer 27

If more electrons : electrode potential becomes more negative

If less electrons : electrode potential becomes more positive

+ has a knock on effect on the e.m.f. value - may become larger or smaller

Question 28

How can the direction of electron flow be worked out?

Answer 28

Look at diagram.

e. g.
- From right to left
- The Cu2+ is more concentrated on the left, so reduction on Cu2+ is more likely to happen on the left (Cu2+ + 2e- <=> Cu)
- The left electrode is the positive electrode, so the right electrode is the negative electrode