C23 - Redox and Electrode Potentials

Question 1

What are the two common redox titration?

Answer 1

Potassium manganate (VII) KMnO4 under acidic conditions

Sodium thiosulfate (Na2S2O3) for the determination of iodine

Question 2

How is a manganate (VII) titration carried out?

Answer 2

MnO4- ions are reduced so the other chemical used must be a reducing agent that is oxidised.

1. A standard solution of potassium manganate is added to the burette.
2. Using a pipette, add a measured volume of the solution being analysed to the conical flask. An excess of dilute sulfuric acid is also added to provide the H+ ions needed for the reduction of manganate ions. It is self indicating.
3. During the titration, the manganate solution reacts and is decolourised as it’s being added. The end point is the first permanent pink colour, indicating when there’s an excess of manganate ions.
4. Repeat until concordant results are obtained.

Question 3

How are redox equations balanced?

Answer 3

By balancing elements and balancing charge.

Question 4

What can iodine/thiosulfate redox titration be used for?

Answer 4

Determining the ClO- content in household bleach

Determining the Cu 2+ co tent in copper (II) compounds

Determining the Cu content in copper alloys

Question 5

How is an iodine/thiosulfate redox titration carried out?

Answer 5

1) Add a standard solution of Na2S2O3 to the burette.
2) Prepare a solution of the oxidising agent to be analysed. Using a pipette, and the solution being tested to a conical flask.
3) Add an excess of potassium iodide. The oxidising agent reacts with iodide ions to produce iodine, which turns the solution a yellow/brown colour.
4) Titrate the solution with the Na2S2O3. During the titration, the iodine is reduced back to I- ions and the brown colour fades so an end point is hard to see.
5) Starch indicator is added when the end point is being neared and the iodine/thiosulfate solution is a straw colour.

This forms a deep blue/black colour. Once this fades, as thiosulfate continues to be added, the end point would have been reached.

Question 6

What’s a half cell?

Answer 6

Components of a voltaic cell containing the chemical species present in a redox half equation.
E.g. a copper electrode in a copper sulphate solution.

In the cell, the chemicals in the two half cells must be kept apart. If mixed, electrons would flow in an uncontrolled way and heat energy would be released instead of electrical.

Question 7

What do simple half cells consist of?

Answer 7

A metal rod dipped in a solution of its aqueous metal ion.

When two are connected, electrons will flow depending on the relative tendency of each electrode to release electrons.

Question 8

In an electrochemical cell, what is the negative electrode?

Answer 8

The electrode with the more reactive metal. It will be oxidised and lose electrons.

Question 9

In an electrochemical cell, what is the positive electrode?

Answer 9

The electrode with the less reactive metal. It will be reduced and gain electrons.

Question 10

What is standard electrode potential, E ⊖?

Answer 10

The voltage of a half cell compared to that of a standard hydrogen half cell.

The standard conditions are:

• conc of 1moldm-3
• 298K
• 100kPa

Question 11

How are oxidation numbers used to balance (redox) equations?

Answer 11

Ignore spectator ions

Identify the species with changes in oxidation states and balance these first (balance each with the oxidation states of the other).

Balance the rest

Question 12

What are the standard conditions to measure electrode potential values?

Answer 12

Solutions of each ha.f cell have a concentration of exactly 1moldm3

298K

100kPa

Question 13

What does the sign (+/-) of standard electrode potential values show?

Answer 13

How it compares to the E value of hydrogen (0V).

A hydrogen half cell (made of a platinum electrode and hydrogen ions) can be connected to another half cell to identify the standard electrode potential of the other metal.

Question 14

How can a standard electrode potential be measured?

Answer 14

By connecting a half cell to a standard hydrogen electrode (consisting of a platinum electrode, hydrogen gas and H+ ions)

The 2 electrodes are connected by a wire to a high resistance voltmeter to allow controlled flow of electrons.

The solutions are connected by a ‘salt bridge’ containing a concentrated solution of an electrolyte that won’t react (e.g. KNO3).

The value on the voltmeter provides the standard electrode potential value.

Question 15

What does a more negative standard electrode potential (E°) value of a redox system suggest?

Answer 15

The greater the tendency to lose electrons and undergo oxidation.

The less the tendency to gain electrons and undergo reduction.

Question 16

What does a more positive standard electrode potential (E°) value of a redox system suggest?

Answer 16

The greater the tendency to gain electrons and undergo reduction.

The less the tendency to lose electrons and undergo oxidation.

Question 17

What does the standard electrode potential (E°) of metals and non metals suggest?

Answer 17

The more negative the E° value, the greater the reactivity of a metal in losing electrons.

The more positive the E° value, the greater the reactivity of a non-metal in gaining electrons.

Question 18

What’s emf?

Answer 18

The maximum voltage produced by an electrochemical cell.

Question 19

How is the emf of an electrochemical cell identified?

Answer 19

E(cell) = less negative standard electrode potential value - more negative standard electrode potential value

These can be found by connecting two half cells.

E (cell) = E° (positive electrode) - E° (negative electrode)

Question 20

What happens in a zinc-copper cell?

Answer 20

(The 2 standard half cells would be connected to a voltmeter and each other via a salt bridge.)

The copper half cell has the more positive E° value and a greater tendency to undergo reduction and gain e-.

The zinc half cell has the more negative E° value and a greater tendency to undergo oxidation and lose e-.

The zinc electrode is negative the the copper is positive.

Zn + Cu2+ -> Zn2+ + Cu

Question 21

How is emf / E(cell) value calculated?

Answer 21

E° cell = E° (positive electrode) - E° (negative electrode)

Question 22

What is the rule for the feasibility of a redox system?

Answer 22

“Top right - bottom left” (when the standard electrode potential table is arranged from most to least negative).

Question 23

What are the limitations of using standard electrode potential values to predict reactions?

Answer 23

• The don’t demonstrate reaction rate (feasibility is based on G). It may have a high activation energy.
• Concentration. If not 1moldm3, it won’t be standard conditions. For the redox systems, if conc is below 1M, eq will gift to the left to increase electrons. This mass the electrode potential more negative.
• Non standard conditions would affect overall E° value.

Question 24

What are the 3 main types of cell?

Answer 24

Primary
Secondary
Fuel (cell)

Question 25

What are primary cells?

Answer 25

Non-rechargeable, one use only cells.

When in use, electrical energy is produced by oxidation and reduction and electrons.

Question 26

What are secondary cells?

Answer 26

Rechargeable, multi use cells.

Unlike primary cells, their cell reaction producing electrical energy can be reversed.

Question 27

What are fuel cells?

Answer 27

Cells using energy from the reaction of a fuel with oxygen to create a voltage.

Question 28

How does a fuel cell work?

Answer 28

The fuel and oxygen will flow into the fuel cell and the products flow out.
The electrolyte remains in the cell.

They’re continuous and don’t have to be recharged.

The most common are hydrogen and methanol fuel cells.

Question 29

What are the redox and overall equations for the reactions in an alkali hydrogen fuel cell?

Answer 29

Oxidation:
H2 + 2OH- —> 2H2O + 2e-

Reduction:
O2 + 2H2O + 4e- —> 4OH-

Overall:
2H2 + O2 —> 2H2O

Question 30

What are the redox and overall equations for the reactions in an acid hydrogen fuel cell?

Answer 30

Oxidation:
H2 —> 2H+ + 2e-

Reduction:
O2 + 4H+ + 4e- —> 2H2O

Overall:
2H2 + O2 —> 2H2O

Question 31

What does a more positive standard electrode potential value mean?

Answer 31

It’s a stronger oxidising agent.

Equilibrium is more likely to gain electrons, shift to the right and undergo reduction.