Molecular Biology of the Cell by Bruce Alberts, Alexander Johnson, Julian Lewis, David Morgan, Martin Raff, Keith Roberts, Peter Walter by by Bruce Alberts, Alexander Johnson, Julian Lewis, David Morg

PANEL 14–1: Redox Potentials 765
HOW REDOX POTENTIALS ARE MEASURED
e –
A reduced and A oxidized
in equimolar amounts
voltmeter
salt bridge
1 M H + and
1 atmosphere H 2 gas
THE STANDARD REDOX POTENTIAL, E′ 0
One beaker (left) contains substance A with an equimolar
mixture of the reduced (A reduced ) and oxidized (A oxidized )
members of its redox pair. The other beaker contains the
hydrogen reference standard (2H + + 2e – H 2 ), whose redox
potential is arbitrarily assigned as zero by international
agreement. (A salt bridge formed from a concentrated KCl
solution allows K + and Cl – to move between the beakers, as
required to neutralize the charges when electrons flow
between the beakers.) The metal wire (dark blue) provides a
resistance-free path for electrons, and a voltmeter then
measures the redox potential of substance A. If electrons flow
from A reduced to H + , as indicated here, the redox pair formed
by substance A is said to have a negative redox potential. If
they instead flow from H 2 to A oxidized , the redox pair is said to
have a positive redox potential.
examples of redox reactions
standard redox
potential E 0
′
The standard redox potential for a redox pair,
defined as E 0 , is measured for a standard state
where all of the reactants are at a concentration of
1 M, including H + . Since biological reactions occur at
pH 7, biologists instead define the standard state as
A reduced = A oxidized and H + = 10 –7 M. This standard
redox potential is designated by the symbol E 0
′ , in
place of E 0 .
c
c
CALCULATION OF ΔG o FROM
REDOX POTENTIALS
To determine the energy change for an electron
transfer, the ΔG o of the reaction (kJ/mole) is calculated
as follows:
ΔG o = –n(0.096) ΔE 0
′ , where n is the number of
electrons transferred across a redox potential
change of ΔE′
0 millivolts (mV), and
ΔE 0
′ = E 0
′(acceptor) – E 0
′(donor)
EXAMPLE:
e –
EFFECT OF CONCENTRATION CHANGES
As explained in Chapter 2 (see p. 60), the actual free-energy
change for a reaction, ΔG, depends on the concentration of the
reactants and generally will be different from the standard freeenergy
change, ΔG o . The standard redox potentials are for a 1:1
mixture of the redox pair. For example, the standard redox
potential of –320 mV is for a 1:1 mixture of NADH and NAD + .
But when there is an excess of NADH over NAD + , electron
transfer from NADH to an electron acceptor becomes more
favorable. This is reflected by a more negative redox potential
and a more negative ΔG for electron transfer.
NADH
oxidized
ubiquinone
NAD +
reduced
ubiquinone
1:1 mixture of
1:1 mixture of oxidized
NADH and NAD +
and reduced ubiquinone
For the transfer of one electron from NADH to
ubiquinone:
ΔE′
0 = +30 – (–320) = +350 mV
ΔG o = –n(0.096) ΔE 0
′ = –1(0.096)(350) = –34 kJ/mole
NADH
NAD +
The same calculation reveals that the transfer of one
electron from ubiquinone to oxygen has an even more
favorable ΔG o of –76 kJ/mole. The ΔG o value for the
transfer of one electron from NADH to oxygen is the
sum of these two values, –110 kJ/mole.
′ ′