Redox & electrode potentials

Question 1

What is an oxidising agent?

Answer 1

A substance that tales electrons from the species being oxidised - it contains the species that is being reduced

Question 2

What is a reducing agent?

Answer 2

A substance that adds electrons to the species being reduced - it contains the species that is oxidised

Question 3

In manganate (VII) titrations, whay is done to Mn04 - ions?

Answer 3

They are reduced

Question 4

How is an MnO4 - titration carried out?

Answer 4

1) A standard solution of potassium manganate (VII), KMnO4, is added to the burette

2) Using a pipette, add a measured volume of the solution being analysed to a conical flask. An excess of dilute sulfuric acid is also added to provide the H+ (aq) ions required for the reduction of MnO4-. An indicator is not required - this titration is self indicating

3) During the titration, the manganate (VII) solution reacts and is decolourised as it is being added. The end point of the titration is judged by the first permanent pink colour, indicating there is an excess of MnO4 - ions present.

4) Repeat the titration until concordant titres are obtained - the titres agree within 0.10cm^3

Question 5

How do you read the meniscus for a KMnO4 titration?

Answer 5

KMnO4 is a deep purple colour so it is difficult to read the bottom of the meniscus through the intense colour. Burette readings are instead read from the top; make sure to read both from the top so the titre is correct

Question 6

What can manganate (VII) titrations be used for the analysis of?

Answer 6

• Iron (II) ions
• Ethandioic acid

Question 7

In iodine-thiosulfate titrations, whats oxidised and whats reduced?

Answer 7

Thiosulfate ions, S2O3 2- are oxidised and iodine, I2, is reduced

Question 8

What are the reactions in an iodine thiosulfate titration?

Answer 8

Question 9

What is the procedure of an iodine-thiosulfate titration?

Answer 9

1) Add a standard solution of Na2S2O3 to the burette
2) Prepare a solution of the oxidising agent to be analysed. Using a pipette, add this solution to a conical flask. Then add an excess of potassium iodide. The oxidising agent reacts with iodine ions to produce iodine, which turns the solution a yellow-brown colour
3) Titrate this solution with the Na2S2O3. During the titration, the iodine is reduced back to I- ions and the brown colour fades quite gradually, making it difficult to decide on an end point. This problem is solved by using a starch indicator. When the end point is being approached, the iodine colour has faded enough to become a pale straw colour. A small amount of starch indicator is added. A deep-blue black colour forms to assist with the identification of the end point. As more sodium thiosulfate is added, the blue-black colour fades. At the end point, all the iodine will have just reacted and the blue-black colour disappears

Question 10

What can iodine/thiosulfate titrations be used to determine?

Answer 10

• The Clo- content in household bleach
• The Cu 2+ content in copper (II) compounds
• The Cu content in copper alloys

Question 11

Copper with an iodine-thiosulfate titration!

Answer 11

Question 12

What does a half-cell contain?

Answer 12

The chemical species present in a redox half-equation

Question 13

How can a voltaic half cell be created?

Answer 13

Connecting 2 different half-cells which then allows electrons to flow. The chemicals in the 2 different half-cells aren’t allowed to mix

Question 14

What is a metal/ metal ion half-cell?

Answer 14

The simplest half-cell that consists of a metal rod dipped into a solution of its metal aqueous ion. At the phase boundary where the metal is in contact with its ions, an equilibrium will be set up

Question 15

How is a reaction written for a metal/ metal ion half-cell?

Answer 15

So that the forward reaction shows reduction and the reverse shows oxidation

Question 16

What is an ion/ ion half-cell?

Answer 16

This contains different ions of the same element in different oxidation states. In this type of half-cell there is no metal to transport electrons either into or out of the half-cell, so an inert metal electrode made out of platinum is used

Question 17

Which is the positive electrode in an operating cell?

Answer 17

The electrode with the less reactive metal - it gains electrons and is reduced

Question 18

Which is the negative electrode in an
operating cell?

Answer 18

The electrode with the more reactive metal - it loses electrons and is oxidised

Question 19

What is the standard electrode potential?

Answer 19

The emf of a half-cell connected to a standard hydrogen half-cell under standard conditions of 298K, solution concentrations of 1 mol dm^-3 and a pressure of 100kPa

Question 20

What does the standard electrode potential look like?

Answer 20

Question 21

By definition, what is the standard electrode potential of a standard hydrogen electrode?

Answer 21

0 V

Question 22

How do you measure a standard electrode potential?

Answer 22

Connect it to a standard hydrogen half-cell

• The 2 electrodes are connected by a wire to allow a controlled flow of electrons
• The 2 solutions are connected with a salt bridge which allows ions to flow. The salt bridge typically contains a concentrated solution of an electrolyte that does not react with either solution. An example, is a strip of filter paper soaked in aqueous potassium nitrate

Question 23

What does an electrode between 2 metal half cells look like?

Answer 23

Question 24

What is the cell reaction?

Answer 24

The sum of the oxidation and and reduction half cells taking place in each half

Question 25

What is emf of an electrochemical cell, Ecell?

Answer 25

The difference between the standard electrode potential of the 2 half cells

Question 26

What is the Ecell equation?

Answer 26

Ecell = EΘ (positive electrode) - EΘ (negative electrode)

Positive electrode - the least negative half cell

Question 27

In a reaction, where will electrons flow from and to?

Answer 27

From the negative cell and to the more positive one

Question 28

What does a lower EΘ value generally show?

Answer 28

The greater the tendency it has to release electrons and shift left for a reaction : Fe 2+ (aq) + 2e- <=> Fe(s)

Question 29

Which electrode does oxidation happen at?

Answer 29

The more negative one as electrons are produced

Question 30

Which electrode does reduction happen at?

Answer 30

The more positive one as electrons are accepted

Question 31

If Cu 2+(aq) + 2e- <=> Cu (s) were to be the more positive half cell, what would the equation be?

Answer 31

Cu 2+(aq) + 2e- -> Cu (s)

Reduction is occurring as electrons are being accepted

Question 32

If Cu 2+ + 2e- <=> Cu (s) were the more negative half cell, what would the equation be?

Answer 32

Cu (s) -> Cu 2+(aq) + 2e-

Oxidation is occurring as electrons are being released

Question 33

What is the rule for an electrode equation to be feasible?

Answer 33

Combining the 2 half equations at each electrode the right way round - the reverse of this would not be feasible

Question 34

What are the limitations of using EΘ values to predict feasibility?

Answer 34

• The electrode potentials may indicate the thermodynamic feasibility of a reaction, but they do not account for the rate of reaction (this could be very slow)
• Standard electrode potentials are measured using standard concentrations of 1 mol dm^-3. Many reactions taking place use concentrated or dilute solutions; therefore the value of the electrode potential would be different from the standard value
• The actual conditions used for the reaction may be different from the standard conditions used to record the EΘ values
• Standard electrode values apply to aqueous equilibria - many reactions take place that are not at equilibria

Question 35

What happens to the electrode potential if the concentration is greater than 1mol dm^-3?

Answer 35

The equilibria will shift to the right, removing electrons from the system and making the electrode value less negative/more positive

Question 36

What happens to the electrode potential if the concentration is less than 1mol dm^-3?

Answer 36

The equilibrium will shift to the left increasing electrons in the system and making the electrode potential more negative/ less positive

Question 37

Whats used in batteries and fuel cells?

Answer 37

Redox reactions where there flow of electrons can be used to generate electrical current

Question 38

What are the 3 main types of cell?

Answer 38

• Primary
• Secondary
• Fuel

Question 39

What is a primary cell?

Answer 39

Non-rechargable cells that are designed only to be used once

Question 40

How do primary cells work?

Answer 40

When in use, electrical energy is produced by oxidation and reduction at the electrodes. However, the reaction cannot be reversed. Eventually, the chemicals will be used up, voltage will fall, the battery will go flat, and the cell will be discarded or recycled

Question 41

What is secondary cell?

Answer 41

Cells that are rechargeable and can be used more than once

Question 42

How does a secondary cell work?

Answer 42

Unlike primary cells, the cell reaction producing electrical energy can be reversed during recharging. The chemicals in the cell are generated and then can be used again

Question 43

What are some common examples of secondary cells include?

Answer 43

• Lead-acid batteries used in car batteries
• Nickel-cadium and nickel-metal hydride ; the cylindrical batteries used in radios, torches etc
• Lithium-ion and lithium-ion polymer cells used our modern appliances - laptops, cameras, mobile phone phones - and also being developed for cars

Question 44

What is a fuel cell?

Answer 44

A cell that uses the energy from the reaction of a fuel with oxygen to create a voltage

Question 45

How does a fuel cell work?

Answer 45

• The fuel and oxygen flow into the fuel cell and the products flow out. The electrolyte remains in the cell
• They operate continuously providing the fuel and oxygen and supplied into the cell
• They do not need to be recharged

Question 46

What are the most common types of fuel cells?

Answer 46

Hydrogen

Question 47

What are the products from a hydrogen cell?

Answer 47

Water only

Question 48

What are the 2 types of electrolytes a hydrogen cell can have?

Answer 48

Acid or alkali