Electrode Potentials

Question 1

What happens when a rod of metal is dipped into a solution of its own? (1)

Answer 1

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

Question 2

What is an electrode/half cell? (1)

Answer 2

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

Question 3

What does the potential difference show and how? (3)

Answer 3

Show where the position of the equilibrium lies
- If there is a LARGE voltage the equilibrium is to the RIGHT
- If there is a SMALL voltage the equilibrium is to the LEFT

Question 4

What is the simplest salt bridge made of? (1)

Answer 4

Filter paper soaked in Potassium Nitrate solution

Question 5

Why are salt bridges necessary? (3)

Answer 5

Complete the circuit
Allows ion movements to balance the charge
Do not react with electrodes

Question 6

Why is KNO3 a suitable solution for a salt bridge? (2)

Answer 6

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

Question 7

Why can the voltmeter give a voltage reading? (1)

Answer 7

Prevents electrons flowing

Question 8

Why might the current produced by a cell fall to zero after some time? (1)

Answer 8

All the reactants are used up

Question 9

What will happen to a cell once the reactants are used up? (1)

Answer 9

Stops working OR starts to leak

Question 10

Why is platinum a suitable electrode? (2)

Answer 10

Pt is unreactive AND conducts electricity

Question 11

When is a platinum electrode used? (1)

Answer 11

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

Question 12

What is the voltage of the Standard Hydrogen Half Cell? (1)

Answer 12

Zero

Question 13

Describe a standard hydrogen electrode (4)

Answer 13

Hydrogen/H2 gas/bubbles
1.0 mol dm–3 H+
At 298K and 100kPa
Pt (electrode)

Question 14

When is the the standard electrode potential value given? (1)

Answer 14

If the cell connected to the standard hydrogen electrode is also in standard condition

Question 15

What is an electrochemical series? (1)

Answer 15

A series of chemical elements arranged in the order of their standard electrode potential

Question 16

The standard electrode potential of Cu2+/Cu is 0.37 V. Why might the electrode potential of the following cell not be 0.37 V? (1)

Answer 16

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

Question 17

Why might other standard electrodes be used occasionally? (2)

Answer 17

Cheaper/easier/quicker to use and provide just as good a reference
Platinum is expensive

Question 18

What does a vertical solid lines indicate in a cell notation diagram? (1)

Answer 18

Phase boundary

Question 19

What does a double vertical line indicate in a cell notation diagram? (1)

Answer 19

Salt bridge

Question 20

Equation for Emf (1)

Answer 20

more positive – least positive

Question 21

What is the function of the platinum electrode? (1)

Answer 21

To allow transfer of electrons/provide a reaction surface

Question 22

Give one reason why the emf of this cell changes when the electrodes are connected and a current flows? (3)

Answer 22

Concentration of the ions change
OR
Are no longer standard
OR
The emf is determined when no current flows

Question 23

State why the electrode potential for the standard hydrogen electrode is equal to 0.00V? (1)

Answer 23

By definition

Question 24

What factors will change E note value? (2)

Answer 24

Concentration of ions
Temperature

Question 25

What happens if you reduce the concentration of ions in the left hand half cell? (4)

Answer 25

Equilibrium moves to the left to oppose the change of removing ions
This releases more electrons
The E note of the left hand half cell becomes more negative
So the emf of the cell increases

Question 26

When is a reaction feasible? (1)

Answer 26

Ecell value is positive (above zero)

Question 27

How are cells recharged? (2)

Answer 27

Reactions are reversible
Reversed by running a higher voltage through the cell then the cell E note value

Question 28

Where are lithium ion cells used? (2)

Answer 28

Mobile phones
Laptops

Question 29

Give an environmental advantage of using rechargeable cells (1)

Answer 29

Metals are reused

Question 30

Give an environmental disadvantage of using rechargeable cells (2)

Answer 30

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

Question 31

What is a fuel cell? (3)

Answer 31

Generates an electrical current from electrochemical reactions.
A fuel cell uses the energy from the reaction of a fuel with oxygen to create a voltage
Does not require electrical recharging

Question 32

Advantages of using fuel cells for energy instead of fossil fuels (2)

Answer 32

Greater efficiency than burning hydrogen in a combustion engine
Less-polluting as water is the only product

Question 33

Disadvantages of using fuel cells for energy instead of fossil fuels (2)

Answer 33

H2 is difficult to store
Fossil fuels are combusted to produce the hydrogen, which releases carbon dioxide

Question 34

Advantages of fuel cells compared to other types of cell (2)

Answer 34

Voltage is constant, as fuel and oxygen is supplied constantly
So concentrations of reactants remain constant.

Question 35

Reactions that take place in an alkali hydrogen fuel cell (2)

Answer 35

At the anode (negative electrode):
H2(g) + 2OH-(aq) –><– 2H2O(l) + 2e-

At the cathode (positive electrode)
O2(g) + 2H2O(l) + 4e- –><– 4OH-(aq)

Question 36

Reactions that take place in an acid hydrogen fuel cell (2)

Answer 36

At the anode (negative electrode):
H2(g) –><– 2H+(aq) + 2e-

At the cathode (positive electrode):
½ O2(g) + 2H+(aq) + 2e- –><–H2O(l)

Question 37

How can the emf of a cell be kept constant (3)

Answer 37

Reagents supplied constantly
So the concentrations of ions are constant
E note value remains constant

Question 38

Give an equation for the reaction that occurs at the positive lithium cobalt oxide electrode
Give an equation for the reaction that occurs at the negative lithium electrode (2)

Answer 38

Positive:
Li+ + CoO2 + e– → Li+[CoO2]–

Negative electrode:
Li → Li+ + e–