Chapter 23 - Redox & Electrode Potentials

Question 1

How do you make a simple half cell?

Answer 1

A metal rod dipped into a solution of its aqueous metal ion

Question 2

What is an ion half cell and give an example?

Answer 2

• It contains ions of the same element in different oxidation states
• For example a mixture of aqueous iron(ii) and iron(iii)
• The redox equilibrium would be Fe3+(aq) + e- -> Fe2+(aq)

Question 3

What sort of electrode is used in an ion half-cell and why?

Answer 3

There is no metal to transport electrons either into or out of the half cell so an inert metal electrode e.g platinum is used

Question 4

What is the standard electrode potential?

Answer 4

The e.m.f of a half-cell compared with a standard hydrogen half-cell measured at 298k, solution concentrations of 1moldm-3 and a gas pressure of 100kPa

Question 5

What does the standard electrode potential show?

Answer 5

The tendency to gain electrons compared with the hydrogen half-cell

Question 6

Give an example of a salt bridge and explain why this is used?

Answer 6

A strip of filter paper soaked in aqueous potassium nitrate

Contains a conc. electrolyte that does not react with either solution

Question 7

What does a more negative electrode potential value suggest?

Answer 7

• A greater tendency to lose electrons and undergo oxidation
• Less tendency to gain electrons and undergo reduction

Question 8

What does a more positive electrode potential value suggest?

Answer 8

• A greater tendency to gain electrons and undergo reduction
• Less tendency to lose electrons and undergo oxidation

Question 9

In an ion half cell how can getting 1 moldm-3 solutions be difficult? What is the solution?

Answer 9

• It could be challenging to dissolve enough solute to get this conc.
• If there are equal ion concentrations it can give you the same electrode potential value

Question 10

How do you calculate standard cell potential?

Answer 10

Electrode potential of positive electrode - Electrode potential of negative electrode

Question 11

Given the equations:
A: Cr3+(aq) + 3e- -> Cr(s)
B: Cu2+(aq) + 2e- -> Cu(s)
C: Ag+(aq) +e- -> Ag(s)

And their standard electrode potentials A: -0.77 B: +0.34 C: +0.80

Will C react with A and B?

Answer 11

• Redox system C has a more positive electrode potential value
• C will therefore have a greater tendency to be reduced than A or B
• Oxidising agent on the left of C (Ag+(aq)) should react with reducing agents on the right of A and B (Cr(s) and Cu(s))

Question 12

How is reaction rate a limitation of predictions using electrode potential values?

Answer 12

• Some reactions have very large activation energies
• Electrode potentials may indicate the feasibility of a reaction but they give no indication on the rate of the reaction

Question 13

How is concentration a limitation of predictions using electrode potential values?

Answer 13

• Many reactions take place using solutions which are not 1moldm-3 and so the value of the electrode potential would be different

Question 14

Give an example of how concentration could alter electrode potential

Answer 14

E.g for Zn2+(aq) + 2e- -> Zn(s)

• If the conc. of Zn2+(aq) is greater than 1moldm-3, equilibrium will shift to the right, removing electrons and the electrode potential would be less negative
• If the conc. of Zn2+(aq) is less than 1moldm-3, equilibrium will shift to the left, increasing electrons in system and the electrode potential will be more negative

Question 15

What is a primary cell?

Answer 15

• Non-rechargeable

- Electrical energy is provided by oxidation and reduction at the electrodes

Question 16

What is a secondary cell?

Answer 16

• Rechargable

- The cell reaction producing electrical energy can be reversed during recharging

Question 17

What is a fuel cell?

Answer 17

It uses the energy from the reaction of a fuel with oxygen to create a voltage and the changes that take place at each electrode

Question 18

Advantages of fuel cell?

Answer 18

(1) Fuel cells can operate continuously provided that fuel and oxygen are supplied into the cell
(2) Fuel cells do not have to be recharged

Question 19

Will Cl2 react with Fe2+

Answer 19

The Electrode potential value for Cl2/Cl- is more positive than Fe3+/Fe2+ so Cl2 is more readily reduced than Fe3+, therefore Cl2 will react with Fe2+

You want the left hand side of the more positive and the right hand side of the more negative

Question 20

How do you calculate standard cell potential?

Answer 20

Electrode potential of cell being reduced - electrode potential of cell being oxidised

Question 21

Half equations for the reaction of Fe2+ with Cr2O72-

Answer 21

Cr2O72−(aq) + 14H+(aq) + 6e− → 2Cr3+(aq) + 7H2O(l)

Fe2+(aq) → Fe3+(aq) + e−

Question 22

Ionic equation for reaction of Fe2+ with MnO4-

Answer 22

MnO4−(aq) + 5e− + 8H+(aq) → Mn2+(aq) + 4H2O(l)

Fe2+(aq) → Fe3+(aq) + e−

Question 23

How do you balance an equation using oxidation numbers?

Answer 23

1) Assign oxidation numbers to identify atoms that change their oxidation number
2) Balance the amount of oxidation and reduction using only species that have elements which have changed oxidation number
3) Balance any remaining atoms

Question 24

How do you check whether an equation is balanced?

Answer 24

Ensure both sides of the equation are balanced by charge

Question 25

How do you predict a half equation?

Answer 25

1) Assign oxidation numbers
2) Balance the electrons according to oxidation numbers
3) Balance any remaining atoms
- if oxygen has been lost a likely product is H2O
- balance hydrogen produced with H+ on other side of the equation

Question 26

What is added to the Burette and what is added to the comical flask in a manganate titration?

Answer 26

Burette - standard solution if potassium manganate (VII)

Conical flask - solution being analysed and an excess of sulfuric acid to provide H+ ions for reduction of MnO4-

Question 27

When is the end point of a manganate titration and what is the colour change?

Answer 27

• The first permanent pink colour

- colourless to light pink

Question 28

How do you analyse the percentage purity of an iron(II) compound?

Answer 28

1) calculate the amount of known oxidising agent that reacted
2) Determine the amount of Fe2+ that reacted (in sample)
3) Scale up to find the amount of Fe2+ in the original solution
4) Find the mass of the hydrated iron/Fe
5) Find percentage purity

Question 29

How can you be presented with non-familiar redox titrations?

Answer 29

1) Mangante (VII) titrations can be used to analyse reducing agents that reduce MnO4- to Mn2+ e.g fe2+ and ethanedioic acid
2) KMnO4 can be replaced with another oxidising agent e.g acidified dichromate (VI)

Question 30

What is added to the Burette and what is added to the conical flask in an iodine-thiosulfate titration?

Answer 30

Burette - Na2S2O3

Conical flask - oxidising agent to be analysed and potassium iodide (liberates iodine)

Question 31

What are the colour changes during iodine-thiosulfate titration?

Answer 31

1) When potassium iodide is added to oxidising agent it produces a solution which is a yellow brown colour

2) During titration solution turns from brown to pale straw colour (iodine getting reduced)

3) After starch indicator is added it turns from black to colourless at end point

Question 32

What other oxidising agents can be analysed using iodine?

Answer 32

1) Chlorate(I) ions ClO-

2) Copper(II) ions Cu2+