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 5

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

Answer 5

||

Question 6

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

Answer 6

The most reduced species

Question 7

What does|indicate?

Answer 7

Phase boundary (solid/liquid/gas)

Question 8

What happens at the left-hand electrode?

Answer 8

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

Question 9

What happens at the right-hand electrode?

Answer 9

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

Question 10

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

Answer 10

Oxidation - left hand electrode

Question 11

What conditions is the standard hydrogen electrode used in?

Answer 11

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

Question 12

What is the standard hydrogen electrode used for?

Answer 12

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

Question 13

Why might you use other standard electrodes occassionally?

Answer 13

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

Question 14

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

Answer 14

Better reducing agent (easier to oxidise)

Question 15

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

Answer 15

Better oxidising agent (easier to reduce)

Question 16

What factors will change E0 values?

Answer 16

Concentration of ions
Temperature

Question 17

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

Answer 17

• 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 18

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

Answer 18

E0cell = E0right - E0left

Question 19

When would you use a Platinum electrode?

Answer 19

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

Question 20

Why is platinum chosen as the electrode?

Answer 20

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

Question 21

How would you predict if a reaction would occur?

Answer 21

• 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 22

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

Answer 22

Zinc/copper (II)

Question 23

What are zinc/carbon cells more commonly known as?

Answer 23

Disposable batteries

Question 24

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

Answer 24

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

Question 25

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

Answer 25

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

Question 26

How are cells recharged (if they are rechargeable)?

Answer 26

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

Question 27

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

Answer 27

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

Question 28

Where are lithium-ion cells used?

Answer 28

* Mobile phones * Laptops

Question 29

What reactions occur on discharge in lithium-ion cells?

Answer 29

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

Question 30

What is a fuel cell?

Answer 30

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

Question 31

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

Answer 31

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

Question 32

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

Answer 32

* 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 33

Disadvantages of fuel cells?

Answer 33

* 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 34

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

Answer 34

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

Question 35

What is the reason that some cells cannot be recharged?

Answer 35

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

Question 36

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

Answer 36

Concentrations of the ions change - the reagents are used up

Question 37

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

Answer 37

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