Redox

Question 1

What is a redox reaction?

Answer 1

A reaction involving both reduction and oxidation.

Question 2

What is oxidation?

Answer 2

The loss of electrons (increase in ox. number)

Question 3

What is reduction?

Answer 3

The gain of electrons (decrease in ox. number)

Question 4

What is an oxidising agent?

Answer 4

The specie that is reduced in a reaction and causes another species to be oxidised.

Question 5

What is a reducing agent?

Answer 5

The specie that is oxidised in a reaction and causes another specie tot be reduced.

Question 6

An redox reaction equation can be broken down into?

Answer 6

• The oxidation half equation

- The reduction half equation

Question 7

How to write out half equations? (No oxygen or hydrogen)

Answer 7

• Chlorine gas oxidises iron(II) ions to iron(III) ions. In the process, the chlorine is reduced to chloride ions.
• You would have to know this, or be told it by an examiner.
  1) WRITE WHAT YOU KNOW FOR EACH OF THE HALF EQUATIONS. In the chlorine case, you know that chlorine (as molecules) turns into chloride ions:
  Cl₂ –> Cl-
  2) BALANCE THE ATOMS YOU HAVE.
  Cl₂ –> 2Cl-
  3) NOW ADD IN H₂O, H+, OH- TO BALANCE THE OXYGEN THEN THE HYDROGEN. There is no oxygen or hydrogen in this equation so you can skip this step. You use H+ when the reaction takes place in acidic conditions, and you use OH- when the reaction takes place in alkaline conditions. If it doesn’t say just use H+/H₂O (H₂O if the equation needs oxygen).
  4) ADD IN ELECTRONS TO BALANCE THE CHARGES:
  Cl₂ + 2e- –> 2Cl-

Question 8

Which side of the equation do electrons go on for a reduction and oxidation half equation?

Answer 8

Reduction = gain = left side (not with the products)
Oxidation = loss = right-side (with the products).

Question 9

Example write out a half equation fo manganate ions (MnO₄⁻) reduced to manganese ions (II)?

Answer 9

Step 1) MnO₄⁻ –> Mn²⁺
Step 2) Already balanced
Step 3) MnO₄⁻ –> Mn²⁺ + 4H₂O
MnO₄⁻ + 8H+ –> Mn²⁺ + 4H₂O
Step 4) MnO₄⁻ + 8H+ + 5e- –> Mn²⁺ + 4H₂O
Final half equation: MnO₄⁻ + 8H+ + 5e- –> Mn²⁺ + 4H₂O

Question 10

This is a redox reaction: Mg(s) + 2HCl(aq) –> MgCl₂(aq) + H₂(g)

Write the two half-equations that produced this redox equation?

Answer 10

Look at oxidation numbers of each of the species in the equation to find which species are reduced and which specie is oxidised. Mg oxidised (ox. no increases from 0 to +2), H reduced (ox no. decreases from +1 to 0). Then follow steps about writing half-equations to get:

Mg –> Mg²⁺ + 2e- (Oxidation half equation)
2H⁺ + 2e- –> H₂ (Reduction half equation)

Cl- is not included in either of these equations because it is not involved in the redox reactions, its oxidations remains the same through the reaction (-1).

Question 11

How do we combine two half equations (reduction and oxidation) to produce an overall ionic equation?

*How do we get it from an overall ionic equation to the redox overall equation? Do we need to know the spectator ions?

Answer 11

Make the number of electrons on the left hand side of the reduction half equation equal to the number of electrons on the right hand side of the oxidation equation (ensure the electrons are balanced). You can do this by multiplying one or both of the equations by a certain factor. Then add the two equations, such that the electrons on each side of the equation cancel out..

Question 12

Combine the half equations:
MnO₄⁻ + 8H+ + 5e- –> Mn²⁺ + 4H₂O and
Fe²⁺ –> Fe³⁺ + e-

Answer 12

5 x (Fe²⁺ –> Fe³⁺ + e-) = 5Fe²⁺ –> 5Fe³⁺ + 5e- +
MnO₄⁻ + 8H+ + 5e- –> Mn²⁺ + 4H₂O
= 5Fe²⁺ + MnO₄⁻ + 8H+ –> 5Fe³⁺ + Mn²⁺ + 4H₂O

Question 13

How do we write a redox reaction from words?
Example: In acid conditions silver metal, Ag, is oxidised to silver(I) ions, Ag+, by NO₃- ions, which are reduced to nitrogen(II) oxide, NO.

Answer 13

1) WRITE OUT WHAT YOU KNOW FOR THE EQUATION, IDENTIFYING REACTANTS AND PRODUCTS:
Ag + NO₃- –> Ag+ + NO
2) IDENTIFY OXIDATION NUMBER CHANGES AND BALANCE ATOMS CHANGING:
For Ag, ox. number goes from 0 to +1. For NO₃-, ox. number goes from +5 to +2. Balance equation (except the oxygens and hydrogens):
Ag + NO₃- –> Ag+ + NO (no need to balance)
3) BALANCE OXIDATION NUMBER CHANGES:
For NO₃-, ox. number goes down by 3, so for Ag we want the ox. number to go up by 3, so we multiply Ag (and Ag+) by 3:
3Ag + NO₃- –> 3Ag+ + NO
4) ADD IN ANY H₂O, H+, OH- TO BALANCE EQUATION: Add 3H₂O to RHS, to balance oxygen, and then 4H+ on LHS, to balance hydrogens (Done under acidic conditions.

Why do we not balance oxygen and hydrogens when we are balancing the other elements?When we are balancing oxidation number changes, do we balance the the Ag and the Ag+?

Question 14

What are redox titrations used for?

Answer 14

They are used for finding the amount of species being oxidised or reduced.

Question 15

How do you do a redox titration?

Answer 15

A known concentration of either a reducing agent or oxidising agent is placed m a burette and titrated against an unknown concentration of the chemical that is being oxidised or reduced respectively.

Question 16

How do we know when the titration is finished?

Answer 16

Will often involve species that can self indicate - they change colour between different oxidation states. E.g. MnO₄⁻ (aq) is purple but is reduced to Mn²⁺(aq), which is colourless. For a reaction between iron(II) and acidified maganate (MnO₄⁻ (aq)/H+(aq), end point indicated when excess MnO₄⁻ ions are present - indicated by a faint pink colour. This is because all the Fe²⁺ has reacted and the MnO₄⁻ can no longer be reduced to Mn²⁺.

Question 17

NOTE: We can calculate the moles of Fe²⁺ that’s reacted with MnO₄⁻, because we know the concentration and volume of MnO₄⁻, the volume of Fe²⁺ and the molar ratio. When using moles of Fe²⁺ and Mr, to calculate mass, you use Mr of Fe.

Answer 17

N/A

For the rest of this, you are going to have to practice Redox titration calculations.

Question 18

What does a half cell made up of?

Answer 18

A half cell is made up of an element in two oxidation states.

Question 19

What is the arrangment of the simplest half cell?

Answer 19

Metal placed in an aqueous solution of its ions.

e.g. copper metal strip in a solution of Cu²⁺ ions.

Question 20

For a copper half cell, what is occuring at the surface between the two copper oxidation states?

Answer 20

A reaction at equilbrium reaction:
Cu²⁺(aq) + 2e- ⇌ Cu(s)

The forward equation is reduction
The backward equation is oxidation.

By convention, the equilibrium is written with the electrons on the LHS. The electrode potential indicates its tendency to lose or gain electrons i.e. it will only favour one direction when there is another half cell, that is connected to it, at which point the metal strip will act as an electrode.

Question 21

Half cells can also be made from non-metals. What is the arrangement of a hydrogen half cell?

Answer 21

Hydrogen gas is bubbled into a solution of aqueous H+ ion (e.g. from a solution of HCl) through an inlet of a glass tube. Metal electrodes will react with the Hydrogen gas/aqueous H+ solution, so an inert platinum electrode is used instead which has the sole purpose to be in contact with both the H₂(g) and H+ ions so that when electron transfer occurs between the aqueous hydrogen ions and hydrogen gas, the electrons produced in the half cell can leave, or the electrons produced in the other half cell (oxidation) can enter through a connecting wire.
This would also be the arrangement for e.g. bromine gas and Br+ ions.

*Don’t really understand.`

Question 22

Arrangement for a metal ion/metal ion half cells. E.g. Fe³⁺ and Fe²⁺ ions

Answer 22

• A solution containing Fe³⁺(aq) and Fe²⁺(aq) with the same concentrations.
• An inert platinum electrode to allow electrons to pass into or out of the half cell via a connecting wire.

Question 23

What is the equation that is occuring at a Fe²⁺/Fe³⁺

half cell?

Answer 23

Fe³⁺(aq) + 2e- ⇌ Fe²⁺(aq)

Question 24

Different half cells have different standard electrode potentials. What is the standard electrode potential of a half cell?

Answer 24

The standard electrode potential of a half cell is the emf of a half cell compared with a standard hydrogen half cell measured at 298K with solution concentrations of 1moldm⁻³ and a gas pressure of 100kPa.

The hydrogen half cell has an emf of 0V so can be used as a reference to measure other half cells against. When a hydrogen half cell is connected to another half cell, the emf produced is the standard electrode potential of that half cell

Question 25

When connecting one half cell to another half cell, what else must your require?

Answer 25

You must have a complete circuit:

• Therefore a salt bridge must be present from solution to solution, so that ions can move through it.
• A wire with a voltmeter between the electrodes -the wire allows electrons to flow through it.
• The voltmeter measures the difference in electrode potentials.

Question 26

What does the emf tell you?

Answer 26

The emf tells you the tendency of different half cells to accept or release electrons.

• When a half cell is connected to a hydrogen half cell, if the emf/standard electrode potential is positive, then forward reaction of half cell equation favoured (lower tendency to release electrons = reduction occurs).
• If emf/standard electrode potential is negative, then reverse reaction of half cell equation favoured (greater tendency to release electrons = oxidation).
• The difference is standard electrode potentials of two half cells. is also known as the emf, and thsi emf give us the cell potential.
• The difference is electrode potential between a hydrogen half cell and another half cell is also known as the emf, and this emf gives the standard electrode potential

Question 27

What is the electrochemical series?

Answer 27

A list in order of their standard electrode potentials (e.g. positive to negative or vice versa).

Question 28

In a simple electrochemical cell, there are two half cells with different electrode potentials, in which one half cell release electrons and the other gains electrons.
If neither of the half cells were hydrogen half cells, how would we calculate which one undergoes reduction and oxidation?

Answer 28

The species with the more negative standard electrode potential, E^Φ, will favour backward reaction, so greater tendency to release electrons, so oxidation occurs at this half cell.
The species with the more positive standard electrode potential, E^Φ, will favour forward reaction, so lower tendency to release electrons, so reduction occurs at this half cell.

Question 29

When two half cells are connected together they will have an overal cell potential
How can we calculate the cell potential when two different half cells are connected together?

*When one of these is a hydrogen half cell, the cell potential is the standard electrode potential of the other half cell.

Answer 29

Cell potential, E𝒸ₑₗₗ^Φ = E^Φ (positive terminal) - E^Φ(negative terminal)
Cell potential = More positive standard electrode potential - Less positive standard electrode potential

Question 30

What does a bigger cell potential/emf mean?

Answer 30

The bigger the value of the emf, the further away from the equilbrium position the reaction moves.

Question 31

NOTE: When combining half cell equations, for the reaction that is undergoing oxidation, you add together its products with the reactants of the reaction that is undergoing reduction.

Answer 31

N/A

Question 32

Why might the reaction between half-cells not take place in reality?

Answer 32

The rate of reaction may be very slow, not high enough activation energy
The actual conditions used for the reaction may be different from the standard conditions used to measure the E^Φ values. This will affect the value of the electrode potential.

Question 33

What is the general rule for when a reaction between two half cells will take place?

Answer 33

The larger the difference in values between E^Φ (standard electrode potential), the more likely a reaction will take place. If teh difference between E^Φ values is less than 0.4V, then a reactin is unlikely to take place.

Question 34

As discussed carrying out the half cell reacions in non-standard conditions can affect the electrode potential. Describe how, the electrode potential of this equation: Cu²⁺(aq) + 2e- ⇌ Cu (s), will change if concentration of Cu²⁺(aq) was increase.

Answer 34

• The equilbrium opposes the change by shifting to the RHS.
• Electrons are less likely to be released (preferred direction of reaction is forward so reduction)
• Electrode potential becomes more positive

Question 35

As discussed carrying out the half cell reacions in non-standard conditions can affect the electrode potential. Describe how, the electrode potential of this equation: Cu²⁺(aq) + 2e- ⇌ Cu (s), will change if temperature of reaction was increased.

Answer 35

• Forming bonds = endothermic, Breaking bonds exothermic.
• Forward reaction is forming bonds so is endothermic.
• When temperature increase, equilbrium opposes the change by favouring the endothermic reaction, so shifts to RHS.
• Electrons are less likely to be released (preferred direction of reaction is forward so reduction)
• Electrode potential becomes more positive

Question 36

There are 2 types of electrochemical cells. What are they?

Answer 36

Non-rechargeable cells -provide electrical energy until the chemicals have reacted to such an extent that the voltage falls. The cell is then flat and is discarded.
Rechargeable cells - the chemicals in the cell react, provide electrical energy. The cell reaction can be reversed during recharging. The chemicals in the cell are regenerated and the cell can be used again.

Question 37

Example of rechargeable batteries?

Answer 37

Nickel and Cadmium (Ni-Cad) batteries, used in rechargeable batteries.

Question 38

What are fuel cells?

Answer 38

These are cells which use external supplies of a fuel and an oxidant, which need to be continuosly supplied. As long as fuel and oxidant is supplied, electrical energy will be produced.

Question 39

What is a hydrogen fuel cell?

Answer 39

This is a fuel cell is where the fuel is hydrogen (or hydrogen rich fuels such as methanol) and the oxidant is oxygen.

Question 40

How does a hydrogen fuel cell work?

Answer 40

1) At anode (negative electrode), the platinum catalyst splits the H2 into protons and electrons
2) The polymer electrolyte membrane only allows the H+ ions to travel across it and forces, the e- to travel through an external connecting wire to the athode.
3) An electric current is created in the circuit as a result.
4) At the cathode (positive electrode), O₂ combines with H+ from the anode and the e- from the external circuit to make H₂O. This is the only waste produce.

Question 41

The hydrogen fuel cell that I have described is what type of hydrogen cell?

Answer 41

It is an acidic hydrogen fuel cell.

Question 42

For an acidic hydrogen cell, what is the equation that is occuring at the anode (negative electrode) and cathode (positive electrode)?

Answer 42

Anode: H₂(g) –> 2H+(aq) + 2e-
Cathode: 1/2O₂(g) + 2H+(aq) + 2e- –> H₂O(l)

Question 43

If we combined the half equations at the cathode and anode of the acidic hydrogen cell, what would the overall equation produce?

Answer 43

1/2O₂(g) + H₂(g) –> H₂O(l)

Question 44

How is an alkaline hydrogen fuel cell produced?

Answer 44

Same as an acidic fuel cell, but the electrolyte is alkaline and the polymer electrolyte membrane only allows OH- ions across.

Question 45

For an alkaline hydrogen cell, what is the equation that is occuring at the anode (negative electrode) and cathode (positive electrode)?

Answer 45

Anode: H₂(g) + 2OH-(aq) –> 2H₂O(l) + 2e-
Cathode: 1/2O₂(g) + H₂O(l) + 2e- –> 2OH-(aq)

Question 46

If we combined the half equations at the cathode and anode of the alkaline hydrogen cell, what would the overall equation produce?

Answer 46

1/2O₂ + H₂ –> H₂O

Question 47

Advantages of Electrochemical cells (not fuel cells)

Answer 47

Cheap to make and have relatively high power densities (a small cell can produce a lot of energy).

Question 48

Disadvantages of Electrochemical cells

Answer 48

The production of the cells involves the use of toxic chemicals, which need to be disposed of correctly of once the cell has reached the end of its life span. These chemicals may need to be neutralised before disposal.
The chemicals used to make the cells are often flammable also.

Question 49

How is the hydrogen fuel stored in a hydrogen fuel cell?

Answer 49

It is stored as COMPRESSED gas. If it cannot be compressed, you may need to liquidify it.

Question 50

In an electrochemical cell, how do you know which electrode is the positive one and the negative one?

Answer 50

The electrode at which reduction occurs(gains electrons by attracting it to them) is the positive electrode/cathode. The electrode at which oxidation occurs (loses electrons) is the negative electrode/anode.