Topic 3 redox equilibria oxidation and Reduction Electrochemical Cells and Fuel Cells



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Topic 5.3
REDOX EQUILIBRIA
Oxidation and Reduction

Electrochemical Cells and Fuel Cells

The Electrochemical Series

Spontaneous Reactions


OXIDATION AND REDUCTION
Redox reactions were studied extensively at AS-level. The key points are summarized here:


  • The gain and loss of electrons can be shown by means of half-equations.




  • Oxidation is the loss of electrons. When a species loses electrons it is said to be oxidised.

Eg Fe2+ ­čí¬ Fe3+ + e


  • Reduction is the gain of electrons. When a species gains electrons it is said to be reduced.

Eg MnO4- + 8H+ + 5e ­čí¬ Mn2+ + 4H2O


  • Electrons can in fact never be created or destroyed; they can only be transferred from one species to another. Reactions which involve the transfer of electrons are known as redox reactions.




  • Overall redox equations can be created by combining the half-equations for the oxidation process and reduction processes, after multiplying all the coefficients of the species in one of the half-equations by a factor which ensures that the number of electrons gained is equal to the number of electrons lost.

Eg Fe2+ ­čí¬ Fe3+ + e oxidation

MnO4- + 8H+ + 5e ­čí¬ Mn2+ + 4H2O reduction

Multiplying all coefficients in the oxidation reaction by 5:

5Fe2+ ­čí¬ 5Fe3+ + 5e

means that 5 electrons are gained and five are lost

overall equation:

MnO4- + 8H+ + 5Fe2+ ­čí¬ Mn2+ + 4H2O + 5Fe3+


  • A species which can accept electrons from another species is an oxidising agent. Oxidising agents are reduced during redox reactions.

Eg MnO4- is the oxidizing agent in the above reaction.


  • A species which can donate electrons to another species is a reducing agent. Reducing agents are oxidised during redox reactions.

Eg Fe2+ is the reducing agent in the above reaction.


  • The oxidation number of an atom is the charge that would exist on the atom if the bonding were completely ionic.


In simple ions, the oxidation number of the atom is the charge on the ion:

Na+, K+, H+ all have an oxidation number of +1.

O2-, S2- all have an oxidation number of -2.
In molecules or compounds, the sum of the oxidation numbers on the atoms is zero:

Eg SO3; oxidation number of S = +6, oxidation number of O = -2.

+6 + 3(-2) = 0
In complex ions, the sum of the oxidation numbers on the atoms is equal to the overall charge on the ion.

Eg MnO4-; oxidation number of Mn = +7, oxidation number of O = -2.

+7 + 4(-2) = -1

Eg Cr2O72-; oxidation number of Cr = +6, oxidation number of O = -2.

2(+6) + 7(-2) = -2

Eg VO2+; oxidation number of V = +5, oxidation number of O = -2.

+5 + 2(-2) = +1
In elements in their standard states, the oxidation number of each atom is zero:

In Cl2, S, Na and O2 all atoms have an oxidation number of zero.


Many atoms, including most d-block atoms, exist in different oxidation numbers. In complex ions or molecules, the oxidation number of these atoms can be calculated by assuming that the oxidation number of the other atom in the species is fixed.


  • Oxidation numbers are useful for writing half-equations:


The number of electrons gained or lost can be deduced from the formula:

No of electrons gained/lost =

change in oxidation number x number of atoms changing oxidation number
The oxygen atoms are balanced by placing an appropriate number of water molecules on one side.
The hydrogen atoms are balanced by placing an appropriate number of H ions on one side.



  • Disproportionation is the simultaneous oxidation and reduction of the same species.

There are many d-block species which readily undergo both oxidation and reduction, and which can therefore behave as both oxidising agents and reducing agents. Cu+, Mn3+ and MnO42- are all examples:
Eg Cu+ ­čí¬ Cu2+ + e oxidation

Cu+ + e ­čí¬ Cu reduction


Eg Mn3+ + 2H2O ­čí¬ MnO2 + 4H+ + e oxidation

Mn3+ + e ­čí¬ Mn2+ reduction


Eg MnO42- ­čí¬ MnO4- + e oxidation

MnO42- + 2H+ + 2e ­čí¬ MnO2 + 2H2O reduction



Species such as these are capable of undergoing oxidation and reduction simultaneously. Disproportionation reactions are special examples of redox reactions.

ELECTROCHEMICAL CELLS
1. Electrode potentials
Consider a zinc rod immersed in a solution containing Zn2+ ions (eg ZnSO4):

The Zn atoms on the rod can deposit two electrons on the rod and move into solution as Zn2+ ions:

Zn(s) == Zn2+(aq) + 2e

This process would result in an accumulation of negative charge on the zinc rod.

Alternatively, the Zn2+ ions in solution could accept two electrons from the rod and move onto the rod to become Zn atoms:

Zn2+(aq) + 2e == Zn(s)

This process would result in an accumulation of positive charge on the zinc rod.
In both cases, a potential difference is set up between the rod and the solution. This is known as an electrode potential.
A similar electrode potential is set up if a copper rod is immersed in a solution containing copper ions (eg CuSO4), due to the following processes:

Cu2+(aq) + 2e == Cu(s) - reduction (rod becomes positive)



Cu(s) == Cu2+(aq) + 2e - oxidation (rod becomes negative)
Note that a chemical reaction is not taking place - there is simply a potential difference between the rod and the solution. The potential difference will depend on the nature of the ions in solution, the concentration of the ions in solution, the type of electrode used and the temperature.
2. Creating an emf
If two different electrodes are connected, the potential difference between the two electrodes will cause a current to flow between them. Thus an electromotive force (emf) is established and the system can generate electrical energy.
The circuit must be completed by allowing ions to flow from one solution to the other. This is achieved by means of a salt bridge - often a piece of filter paper saturated with a solution of an inert electrolyte such as KNO3(aq).
The e.m.f can be measured using a voltmeter. Voltmeters have a high resistance so that they do not divert much current from the main circuit.
The combination of two electrodes in this way is known as an electrochemical cell, and can be used to generate electricity. The two components which make up the cell are known as half-cells.
A typical electrochemical cell can be made by combining a zinc electrode in a solution of zinc sulphate with a copper electrode in a solution of copper sulphate.


The positive electrode is the one which most favours reduction. In this case it is the copper electrode which is positive.
The negative electrode is the one which most favours oxidation. In this case it is the zinc electrode which is negative.
Thus electrons flow from the zinc electrode to the copper electrode.

Reduction thus takes place at the copper electrode: Cu2+(aq) + 2e ­čí¬ Cu(s)

Oxidation thus takes place at the zinc electrode: Zn(s) ­čí¬ Zn2+(aq) + 2e
The overall cell reaction is as follows: Zn(s) + Cu2+(aq) ­čí¬ Zn2+(aq) + Cu(s)
The sulphate ions flow through the salt bridge from the Cu2+(aq) solution to the Zn2+(aq) solution, to complete the circuit and compensate for the reduced Cu2+ concentration and increased Zn2+ concentration. The cell reaction including spectator ions can thus be written as follows: CuSO4(aq) + Zn(s) ­čí¬ Cu(s) + ZnSO4(aq).
The external connection must be made of a metallic wire in order to allow electrons to flow. The salt bridge must be made of an aqueous electrolyte to allow ions to flow.
By allowing two chemical reagents to be connected electrically, but not chemically, a reaction can only take place if the electrons flow externally. The chemical energy is thus converted into electrical energy.

3. Designing electrochemical cells
Half-cells do not necessarily have to consist of a metal immersed in a solution of its ions. Any half-reaction can be used to create a half-cell.
If the half-reaction does not contain a metal in its elemental state, an inert electrode must be used. Platinum is generally used in this case, as it is an extremely inert metal.

If a gas is involved, it must be bubbled through the solution in such a way that it is in contact with the electrode.


A few examples are shown below:
a) Fe3+(aq) + e == Fe2+(aq)

A platinum electrode is used, immersed in a solution containing both Fe2+ and Fe3+ ions:



b) Cr2O72-(aq) + 14H+(aq) + 6e == 2Cr3+(aq) + 7H2O(l)

A platinum electrode is used, immersed in a solution containing Cr2O72-, H+ and Cr3+ ions:



c) Cl2(g) + 2e == 2Cl-(aq)



A platinum electrode is used, immersed in a solution containing Cl- ions. Chlorine gas is bubbled through the solution, in contact with the electrode:

d) 2H+(aq) + 2e == H2(g)



A platinum electrode is used, immersed in a solution containing H+ ions. Hydrogen gas is bubbled through the solution, in contact with the electrode:

In addition to making electricity, half-cells provide important information on the relative ability of a half-reaction to undergo oxidation or reduction. The more positive the electrode, the greater the tendency to undergo reduction, and the more negative the electrode, the greater the tendency to undergo oxidation.




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