Fundamentals of Electrochemistry - W.H. Freeman

1504T_ch14_270-297 02/15/06 7:12 Page 291
A 2
is the dominant form of ascorbic acid, and no protons are involved in the net redox reaction.
Therefore, the potential becomes independent of pH.
A biological example of E°¿ is the reduction of Fe(III) in the protein transferrin, which
was introduced in Figure 7-4. This protein has two Fe(III)-binding sites, one in each half of
the molecule designated C and N for the carboxyl and amino terminals of the peptide chain.
Transferrin carries Fe(III) through the blood to cells that require iron. Membranes of these
cells have a receptor that binds Fe(III)-transferrin and takes it into a compartment called an
endosome into which H is pumped to lower the pH to 5.8. Iron is released from transferrin
in the endosome and continues into the cell as Fe(II) attached to an intracellular metaltransport
protein. The entire cycle of transferrin uptake, metal removal, and transferrin
release back to the bloodstream takes 1–2 min. The time required for Fe(III) to dissociate
from transferrin at pH 5.8 is 6 min, which is too long to account for release in the endosome.
The reduction potential of Fe(III)-transferrin at pH 5.8 is E°¿ 0.52 V, which is too
low for physiologic reductants to reach.
The mystery of how Fe(III) is released from transferrin in the endosome was solved by
measuring E°¿ for the Fe(III)-transferrin-receptor complex at pH 5.8. To simplify the chemistry,
transferrin was cleaved and only the C-terminal half of the protein (designated Trf C )
was used. Figure 14-12 shows measurements of log5[Fe(III)Trf C ]/[Fe(II)Trf C ]6 for free protein
and for the protein-receptor complex. In Equation 14-26, E E°¿ when the log term is
zero (that is, when [Fe(III)Trf C ] [Fe(II)Trf C ]). Figure 14-12 shows that E°¿ for
Fe(III)Trf C is near 0.50 V, but E°¿ for the Fe(III)Trf C -receptor complex is 0.29 V. The
reducing agents NADH and NADPH in Table 14-2 are strong enough to reduce Fe(III)Trf C
bound to its receptor at pH 5.8, but not strong enough to reduce free Fe(III)-transferrin.
log {[oxidized form]/[reduced form]}
2.5
2.0
1.5
1.0
0.5
Free
Fe-Trf c
Fe-Trf c -receptor
complex
0.0
E' °
E' °
–0.5
–0.5 –0.4 –0.3 –0.2
E (V vs. S.H.E.)
–0.1
Figure 14-12 Spectroscopic measurement
of log {[Fe(III)Trf C ]/[Fe(II)Trf C ]} versus potential
at pH 5.8. [S. Dhungana, C. H. Taboy, O. Zak, M. Larvie,
A. L. Crumbliss, and P. Aisen,“Redox Properties of Human
Transferrin Bound to Its Receptor,” Biochemistry, 2004,
43, 205.]
Terms to Understand
ampere
anode
cathode
coulomb
current
E°¿
electric potential
electrode
Faraday constant
formal potential
galvanic cell
half-reaction
joule
Latimer diagram
Nernst equation
ohm
Ohm’s law
oxidant
oxidation
oxidizing agent
potentiometer
power
reaction quotient
redox reaction
reducing agent
reductant
reduction
resistance
salt bridge
standard hydrogen electrode
standard reduction potential
volt
watt
Summary
Work done when a charge of q coulombs passes through a potential
difference of E volts is work E q. The maximum work that can
be done on the surroundings by a spontaneous chemical reaction is
related to the free-energy change for the reaction: work G. If
the chemical change produces a potential difference, E, the relation
between free energy and the potential difference is G nFE.
Ohm’s law (I E/R) describes the relation between current, voltage,
and resistance in an electric circuit. It can be combined with
the definitions of work and power (P work per second) to give
P E I I 2 R.
A galvanic cell uses a spontaneous redox reaction to produce
electricity. The electrode at which oxidation occurs is the anode,
and the electrode at which reduction occurs is the cathode. Two
half-cells are usually separated by a salt bridge that allows ions to
migrate from one side to the other to maintain charge neutrality but
prevents reactants in the two half-cells from mixing. The standard
reduction potential of a half-reaction is measured by pairing that
half-reaction with a standard hydrogen electrode. The term “standard”
means that activities of reactants and products are unity. If
several half-reactions are added to give another half-reaction, the
standard potential of the net half-reaction can be found by equating
the free energy of the net half-reaction to the sum of free energies of
the component half-reactions.
The voltage for a complete reaction is the difference between
the potentials of the two half-reactions: E E E , where E is
the potential of the half-cell connected to the positive terminal of the
potentiometer and E is the potential of the half-cell connected to
the negative terminal. The potential of each half-reaction is given by
the Nernst equation: E E° (0.059 16/n) log Q (at 25°C), where
each reaction is written as a reduction and Q is the reaction quotient.
The reaction quotient has the same form as the equilibrium constant,
but it is evaluated with concentrations existing at the time of interest.
Electrons flow through the circuit from the electrode with the more
negative potential to the electrode with the more positive potential.
Complex equilibria can be studied by making them part of an
electrochemical cell. If we measure the voltage and know the concentrations
(activities) of all but one of the reactants and products,
the Nernst equation allows us to compute the concentration of the
unknown species. The electrochemical cell serves as a probe for
that species.
Biochemists use the formal potential of a half-reaction at pH 7
(E°¿) instead of the standard potential (E°), which applies at pH 0.
E°¿ is found by writing the Nernst equation for the half-reaction and
grouping together all terms except the logarithm containing the formal
concentrations of reactant and product. The combination of
terms, evaluated at pH 7, is E°¿.
Summary 291