Page 163 - Modern Analytical Chemistry

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146 Modern Analytical Chemistry
oxidation
A loss of electrons.
reduction
A gain of electrons.
reducing agent
A species that donates electrons to duction half-reaction. Conversely, the species being reduced is called an oxidizing
another species.
oxidizing agent
A species that accepts electrons from 2+
another species.
Nernst equation
An equation relating electrochemical
potential to the concentrations of *Separating a redox reaction into its half-reactions is useful if you need to balance the reaction. One method for
products and reactants.
Oxalic acid, on the other hand, is oxidized since the oxidation state for carbon in-
creases from +3 in H 2 C 2 O 4 to +4 in CO 2 .
Redox reactions, such as that shown in equation 6.22, can be divided into sepa-
rate half-reactions that individually describe the oxidation and the reduction
processes.
+
H 2 C 2 O 4 (aq)+2H 2 O(l) ® 2CO 2 (g)+2H 3 O (aq)+2e –
–
3+
2+
Fe (aq)+ e ® Fe (aq)
It is important to remember, however, that oxidation and reduction reactions al-
ways occur in pairs.* This relationship is formalized by the convention of calling the
species being oxidized a reducing agent, because it provides the electrons for the re-
3+
agent. Thus, in reaction 6.22, Fe is the oxidizing agent and H 2C 2O 4 is the reducing
agent.
The products of a redox reaction also have redox properties. For example, the
3+
Fe in reaction 6.22 can be oxidized to Fe , while CO 2 can be reduced to H 2 C 2 O 4 .
Borrowing some terminology from acid–base chemistry, we call Fe 2+ the conjugate
reducing agent of the oxidizing agent Fe 3+ and CO 2 the conjugate oxidizing agent of
the reducing agent H 2 C 2 O 4.
Unlike the reactions that we have already considered, the equilibrium position
of a redox reaction is rarely expressed by an equilibrium constant. Since redox reac-
tions involve the transfer of electrons from a reducing agent to an oxidizing agent,
it is convenient to consider the thermodynamics of the reaction in terms of the
electron.
The free energy, ∆G , associated with moving a charge, Q, under a potential, E,
is given by
∆G = EQ
Charge is proportional to the number of electrons that must be moved. For a reac-
tion in which one mole of reactant is oxidized or reduced, the charge, in coulombs, is
Q = nF
where n is the number of moles of electrons per mole of reactant, and F is Faraday’s
–1
constant (96,485 C ×mol ). The change in free energy (in joules per mole; J/mol)
for a redox reaction, therefore, is
∆G =–nFE 6.23
where ∆G has units of joules per mole. The appearance of a minus sign in equation
6.23 is due to a difference in the conventions for assigning the favored direction for
reactions. In thermodynamics, reactions are favored when ∆G is negative, and
redox reactions are favored when E is positive.
The relationship between electrochemical potential and the concentrations
of reactants and products can be determined by substituting equation 6.23 into
equation 6.3
–nFE =–nFE°+ RT ln Q
where E° is the electrochemical potential under standard-state conditions. Dividing
through by –nF leads to the well-known Nernst equation.
balancing redox reactions is reviewed in Appendix 4.