Unit 1: Section 8 - Electrode Potentials and 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 potetnials 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

What happens at the left-hand electrode?

Answer 9

• Left hand electrode is where oxidation occurs
• Left hand electrode is the half cell with the most negative E0 value

Question 10

What happens at the right-hand electrode?

Answer 10

• Right hand electrode is where reduction occurs
• Right hand electrode is the half cell with the most positive E0 value

Question 11

Which side of the cell has the most negative E0 value? What happens to the metal with the most negative E0 value?

Answer 11

Oxidation - left hand electrode

Question 12

What conditions is the standard hydrogen electrode used in?

Answer 12

Temperature - 298K
Pressure - 100kPa
[H+] - 1.00 mol dm^-3

Question 13

What is the standard hydrogen electrode used for?

Answer 13

• Comparing other cells agaisnt E0
• E0 of SHE is defined as 0, so all other E0 values are compared against it

Question 14

Why might you use other standard electrodes occassionally?

Answer 14

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

Question 15

If an E0 value is more negative, what does it mean in terms of oxidising/reducing power?

Answer 15

Better reducing agent (easier to oxidise)

Question 16

If an E0 value is more positive, what does it mean in terms of oxidising/reducing power?

Answer 16

Better oxidising agent (easier to reduce)

Question 17

What factors will change E0 values?

Answer 17

Concentration of ions
Temperature

Question 18

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

Answer 18

• Equilibrium moves to the left to oppose the change of removing ions; this releases more electons
• The E0 of the left hand cell becomes more negative, so the e.m.f of the cell increases

Question 19

How do you calculate the emf of a cell from E0 values?

Answer 19

E0cell = E0right - E0left

Question 20

When would you use a Platinum electrode?

Answer 20

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

Question 21

Why is platinum chosen as the electrode?

Answer 21

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

Question 22

How would you predict if a reaction would occur?

Answer 22

• Take the 2 half equations
• FInd the secies that is being reduced (this is effectively the right hand electrode)
• Calculate its E0 value minus the E0 value fo the species that is being oxidised (effectively the left hand cell)
• If E0 overall > 0, reaction will occur

Question 23

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

Answer 23

Zinc/copper (II)

Question 24

What are zinc/carbon cells more commonly known as?

Answer 24

Disposable batteries

Question 25

What are the 2 reactions that take place in zinc/carbon cells?

Answer 25

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

Question 26

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

Answer 26

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

Question 27

How are cells recharged (if they are rechargeable)?

Answer 27

Reactions are reversible and are reversed by running a higher voltage through the cell than the cell’s E0

Question 28

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

Answer 28

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

Question 29

Where are lithium-ion cells used?

Answer 29

• Mobile phones
• Laptops

Question 30

What reactions occur on discharge in lithium-ion cells?

Answer 30

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

Question 31

What is a fuel cell?

Answer 31

A cell that is used to generate electric current; does not require electrical recharging

Question 32

What are the reactions that take place at the 2 electrodes in an an alkaline hydrogen fuel cell?

Answer 32

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

Question 33

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

Answer 33

• In combustion, sulfur containing compounds (SO2, SO3) and nitrogen containing compounds (NO2, NOx) 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 effecient

Question 34

Disadvantages of fuel cells?

Answer 34

• Hydrogen is a flammable gas with a low b.p. - hard and dangerous to store and transport and expensive to buy
• Fuel cells have a limited lifetime and use toxic chemicals in their manufacture

Question 35

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

Answer 35

Most positive E0 value - it is the PRODUCT of the reaction equation

Question 36

What is the reason that some cells cannot be recharged?

Answer 36

Reaction of the cell is not reversible - aproduct is produced that either dissipates or cannot be converted back into the reactants

Question 37

Why might the e.m.f. of a cell change after a period of time?

Answer 37

Concentrations of the ions change - the reagents are used up

Question 38

How can the e.m.f. of a cell be kept constant?

Answer 38

Reagents are supplied constantly, so the concentrations of the ions are constant; E0 remains constant