3.1.11 Electrode Potentials and Electrochemical Cells

Question 1

What happens when a rod of a metal is dipped into a solution of its own ions?

Answer 1

An equilibrium is set up between the solid metal and the aqueous metal ions

Question 2

Write a half equation for zinc (s) to zinc (II)

Answer 2

Zn (s) ⇌ Zn²⁺ (aq) + 2e⁻

Question 3

Write a half-equation for copper (II) to copper (III).

Answer 3

Cu²⁺ (aq) ⇌ Cu³⁺ (aq) + e⁻

Question 4

What is the simplest salt bridge made of?

Answer 4

Filter paper soaked in saturated solution of KNO₃ (potassium nitrate)

Question 5

Why are salt bridges necessary?

Answer 5

• Complete the circuit, but avoid further metal/ion potentials as does not perform electrochemistry.
• Allows ion movement to balance the charge. Do not react with electrodes.

Question 6

What symbol is used to represent a salt bridge in standard notation?

Answer 6

||

Question 7

What type of species goes on the outside (furthest from the salt bridge) in standard cell notation?

Answer 7

The most reduced species

Question 8

What does | indicate?

Answer 8

Phase boundary (solid/liquid/gas)

Question 9

How would an Aluminium/Copper cell be represented?

Answer 9

Al(s)|Al³⁺(aq)||Cu²⁺ (aq)|Cu(s)

Question 10

What happens at the left-hand electrode?

Answer 10

Left hand electrode is where oxidation occurs.
Left hand electrode is the half cell with the most negative E° value

Question 11

What happens at the right-hand electrode?

Answer 11

Right hand electrode is where reduction occurs.
Right hand electrode is the half cell with the most positive E° value

Question 12

Which side of the cell has the most negative E° value?
What happens to the metal with the most negative E° value?

Answer 12

Oxidation - left hand electrode

Question 13

Draw the standard hydrogen electrode.

Answer 13

Question 14

What conditions is the standard hydrogen electrode used in?

Answer 14

Temperature = 298 K
Pressure = 100 kPa
[H⁺] = 1.00 mol dm⁻³

Question 15

What is the standard hydrogen electrode used for?

Answer 15

Comparing other cells against. E° of SHE is defined as 0, so all other E° values are compared against it.

Question 16

Why might you use other standard electrodes occasionally?

Answer 16

• They are cheaper/easier/quicker to use and can provide just as good a reference.
• Platinum is expensive

Question 17

If an E° value is more negative, what does it mean in terms of oxidising/reducing power?

Answer 17

Better reducing agent (easier to oxidise)

Question 18

If an E° value is more positive, what does it mean in terms of oxidising/reducing power?

Answer 18

Better oxidising agent (easier to reduce)

Question 19

What factors will change E° values?

Answer 19

• Concentration of ions
• Temperature

Question 20

What happens if you reduce the concentration of the ions in the left hand half cell?

Answer 20

Equilibrium moves to the left to oppose the change of removing ions. This releases more electrons, the E° of the left hand cell becomes more negative, so the e.m.f. Of the cell increases.

Question 21

How do you calculate the emf of a cell from E° values?

Answer 21

Right E° - Left E°

Question 22

When would you use a Platinum electrode?

Answer 22

When both the oxidised and reduced form of the metal are in aqueous solution.

Question 23

Why is Platinum chosen?

Answer 23

• Inert so does not take part in the electrochemistry
• Good conductor to complete circuit

Question 24

How would you predict if a reaction would occur?

Answer 24

• Take the 2 half equations.
• Find the species that is being reduced (this is effectively the right hand electrode)
• Calculate its E° value minus the E° value of the species that is being oxidised (effectively the left hand cell).
• If E° overall > 0, reaction will occur.

Question 25

What was the first commercial cell made from (Daniell cell)?

Answer 25

Zinc/copper (II)

Question 26

What are zinc/carbon cells more commonly known as?

Answer 26

Disposable batteries

Question 27

What are the two reactions that take place in zinc/carbon cells?

Answer 27

Zn oxidised to Zn²⁺ NH₄⁺ reduced to NH₃ at carbon electrode

Question 28

What are the reactions that occur in a lead/acid battery (car batteries)?

Answer 28

Pb + SO₄²⁻ -> PbSO₄ (s) + 2e⁻ PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ -> PbSO₄ + 2H₂O

Question 29

How are cells recharged (if they are rechargeable)?

Answer 29

Reactions are reversible and are reversed by running a higher voltage through the cell than the cell's E°

Question 30

Nickel/cadmium cells are rechargeable AA batteries etc. What reactions occur at the electrodes?

Answer 30

Cd(OH)₂ (s) + 2e⁻ -> Cd (s) + 2OH⁻ NiO(OH) (s) + H₂O + e⁻ -> Ni(OH)₂ (s) + OH⁻

Question 31

Where are lithium-ion cells used?

Answer 31

- Mobile phones - Laptops

Question 32

What reactions occur on discharge in lithium-ion cells?

Answer 32

Li⁺ + CoO₂ + e⁻ -> Li⁺[CoO₂]⁻ Li -> Li⁺ + e⁻

Question 33

What is a fuel cell?

Answer 33

A cell that is used to generate electric current. Does not require electrical recharging.

Question 34

What are the reactions that take place at the two electrons in an alkaline hydrogen fuel cell?

Answer 34

2H₂ + 4OH⁻ -> 4H₂O + 4e⁻ O₂ + 2H₂O + 4e⁻ -> 4OH⁻

Question 35

Draw a diagram of a hydrogen fuel cell.

Answer 35

Question 36

Why is it better to use a fuel cell than to burn H₂ in air, even though the same overall reaction occurs?

Answer 36

- In combustion, sulfur containing compounds and nitrogen containing compounds are produced due to the high temperatures and the S and N in air. These are bad for the environment. - This does not occur in a fuel cell. The only product is water. - More efficient.

Question 37

Disadvantages of fuel cells?

Answer 37

Hydrogen is a flammable gas with a low boiling point making it difficult and dangerous to store and transport and expensive to buy. Fuel cells have a limited lifetime and use toxic chemicals in their manufacture.

Question 38

How do you find the weakest reducing agent from a table of electrode potential data?

Answer 38

Most positive E° value. Then it is the product of the reduction equation i.e. imagine equation going from right to left.

Question 39

What is the reason that some cells cannot be recharged?

Answer 39

Reaction of the cell is not reversible - a product is produced that either dissipated or cannot be converted back into the reactants.

Question 40

Why might the emf of a cell change after a period of time?

Answer 40

Concentration of the ions change - the reagents are used up.

Question 41

How can the emf of a cell be kept constant?

Answer 41

Reagents are supplied constantly, so the concentrations of the ions are constant. E° remains constant.