5.6 REDOX AND ELECTRODE POTENTIALS

Question 1

What is an oxidising agent?

Answer 1

An oxidising agent takes electrons from the species being oxidised and contains the species being reduced (oxidising agent is reduced)

Question 2

What is a reducing agent?

Answer 2

A reducing agent gains electrons from the species being reduced and contains the species being oxidised (reducing agent is oxidised)

Question 3

What is an example of an oxidation half equation?

Answer 3

Oxidation half equation:
Mg = Mg2+ + 2e-

Question 4

What is an example of a reduction half equation?

Answer 4

Reduction half equation:
Cl2 + 2e- = 2Cl-

Question 5

What are the steps in constructing overall half equations?

Answer 5

Constructing overall half equations:
1. write out the two half equations
2. multiply the half equations so that the number of electrons being lost and gained is the same
3. cancel out the electrons
4. if necessary, cancel out other species which appear on both sides
5. add the remaining equations together

Question 6

What are the steps for constructing half equations?

Answer 6

Constructing half equations:
1. balance the species which changes oxidation state
2. balance out any oxygen atoms with H2O
3. balance out any hydrogen atoms with H+
4. add up the total charges on both sides of the equation
5. use negative electrons to balance the side with the greater positive charge so that each side has the same total charge (not always zero)

Question 7

What is electrochemistry?

Answer 7

Electrochemistry is the study of the relationship between electricity and chemical reactions
there are two major electrochemical processes:
- the generation of electric current from spontaneous chemical reactions
- the use of electricity to cause non-spontaneous chemical reactions to occur

Question 8

What is a voltaic cell?

Answer 8

Voltaic cell:
- converts chemical energy into electrical energy
- requires chemical reactions that transfer electrons from one species to another (redox)
- e.g zinc undergoing oxidation (loses 2 electrons), copper undergoes reduction (gains 2 electrons), copper forms a solid green layer on the piece of zinc, copper sulfate solution begins to fade
- often involve metal displacement reactions
- if the redox reaction takes place in the same beaker, then it occurs spontaneously, but the energy involved in the electron transfer can’t be harnessed (heat energy released rather than electrical energy)

CuSO4 + Zn = ZnSO4 + Cu
Zn = Zn2+ + 2e-
Cu2+ + 2e- = Cu
Zn + Cu2+ = Zn2+ + Cu

Question 9

Half cells and voltaic cells.

Answer 9

Half cells and voltaic cells:
- a half cell contains the chemical species present in a redox half equation
- a voltaic cell can be made by connecting two half cells -allowing electrons to flow. the direction of flow depends on the relative tendency of each electrode to release electrons
- the two half cells must be kept apart (although connected) in order to harness electrical energy

Question 10

What is the anode and cathode?

Answer 10

The anode is the electrode where oxidation occurs (negative electrode)
The cathode is the electrode where reduction occurs (positive electrode)

Question 11

What is the purpose of the salt bridge?

Answer 11

Purpose of the salt bridge:
- allows current to flow by maintaining electronegativity
- prevents an excess of charge building up in each half cell, which would cause electrons to stop flowing
- ions flow out of salt bridge
- e.g filter paper soaked in potassium nitrate

Question 12

What are standard electrode potentials?

Answer 12

Standard electrode potentials:
- the emf (electromotive force) of a half cell connected to a standard hydrogen half cell under standard conditions (100kPa, 298K, 1moldm-3)
- indicates the tendency of a species to be reduced and gain electrons

Question 13

How do you measure standard electrode potentials?

Answer 13

Measuring standard electrode potentials:
- to measure a standard electrode potential, the half-cell is connected to a hydrogen half-cell under standard conditions
- whatever the reading is on the voltmeter is taken as the standard electrode potential for that half cell (as the standard hydrogen half-cell potential is given as 0.00V)

Question 14

What is the equation to calculate cell potential (e cell)?

Answer 14

E cell = cathode potential - anode potential
(more pos - more neg)

Question 15

When is a redox reaction feasible?

Answer 15

A redox reaction is feasible if the e cell is calculated to yield a positive value

Question 16

What are the limitations of e cells?

Answer 16

Limitations of e cells:
- doesn’t take into account the reaction rates (kinetics) of the redox reactions
= could be a very slow reaction due to high activation energy
- the concentration of solutions may not be standard
= if the concentration was greater than 1moldm-3 the equilibrium would shift right so the e cell value would become less negative
- other standard conditions may not be used
- some reactions are not aqueous, for which standard electrode potentials are applied

Question 17

What are the types of half cells?

Answer 17

Types of half cell:
- metals in solutions in their own ions. e.g Al3+ / Al
- gases in solutions in their own ions. e.g 1/2Cl / Cl- / Pt
- solutions of ions in two different oxidation states. e.g Fe3+, Fe2+ / Pt
- oxidising agents in solutions of acid. e.g MnO4- , H+, Mn2+ / Pt

Question 18

Modern storage cells.

Answer 18

Modern storage cells:
all modern cells are based on two electrodes having different electrode potentials. three main types
1. primary
- non-rechargable
- to be used once (until chemicals are used up)
- used for low current, long storage devices
- e.g clock or smoke detector batteries
2. secondary
- rechargable (use reversible reactions)
- chemicals regenerated during charging
- e.g Pb-H+ car batteries, Ni-Cd cells in radios, Li-ion polymer cells in mobiles, laptops etc
3. fuel cell
- uses energy from redox reactions of a fuel cell with oxygen
- e.g hydrogen fuel cells

Question 19

Non-rechargable cells.

Answer 19

Non-rechargable cells:
- zinc/copper cells
= the Daniell cell 1830s
= not portable because it contained liquids
= used for sending telegraphs
- zinc/carbon cells
= the leclanche cell
= used for ordinary disposable batteries
= a paste not a liquid

Question 20

What are fuel cells?

Answer 20

Fuel cells:
- involve a redox reaction of a fuel with oxygen
- lots of different fuels can be used but hydrogen is the most common
- fuel and oxygen must flow continuously into the cell for it to work (no need for recharge)
- the products flow out; electrolyte remains inside

Question 21

What are the advantages and disadvantages of fuel cells?

Answer 21

Advantages:
- only product is water (no ghg emission like CO2)
- hydrogen can be obtained from water or other hydrogen-rich fuels such as methanol -it is readily available and renewable

Disadvantages:
- hydrogen is highly flammable in the presence of oxygen
- difficult to store and transport hydrogen as a gas (there is a lot of research into how it could be adsorbed onto surfaces or absorbed into materials to make it safer and reduce volume)
- energy required to extract hydrogen from water as other fuels (this can involve use of fossil fue;s)

Question 22

What are hydrogen fuel cells?

Answer 22

Hydrogen fuel cells:
- fuel = H2(g) at 1atm
- can contain either an acidic or alkaline electrolyte (which gives rise to different redox systems)

Question 23

Standard electrode potentials and redox agents./

Answer 23

A more reactive metal will have a more negative emf = strongest reducing agent
e.g K+ + e- = K -2.94V is most likely to be oxidised
Strongest oxidising agent when more positive emf
e.g 1/2F2 + e- = F- 2.89V