Chemistry Reference
In-Depth Information
latter was only obtained with the report of its EPR spectrum in frozen alkaline
solution (Schulte-Frohlinde and Eiben 1962; Eiben and Schulte-Frohlinde 1965)
and the advent of pulse radiolysis (Boag and Hart 1963; Keene 1963, 1964). There
is now vast literature on the physical properties and reaction kinetics of e
aq
−
,
and the latter data are only paralleled in number by those of
•
OH (Buxton et al.
1988).
4.2
Redox Reactions
The hydrated electron is the most powerful reductant (E
7
=
−
2.9 V); H
•
has a
somewhat higher reduction potential (E
7
=
2.4 V; for a compilation of reduction
potentials, see Wardman 1989). Often, both
•
H and e
aq
−
are capable of reducing
transition metal ions to their lower oxidation states [e.g., reactions (4) and (5)].
−
Ag
+
+ H
•
Ag + H
+
→
(4)
Ag
+
+ e
aq
−
→
Ag
(5)
However, there are cases where the reduction potential of H
•
is insufficient to re-
duce the metal ion, and the reduction reaction is only given by e
aq
−
[e.g., reaction
(6) (Baxendale and Dixon 1963); for a review see Buxton and Sellers (1977); for a
compilation of rate constants of ensuing reactions see Buxton et al. (1995)].
Zn
2+
+ e
aq
−
Zn
+
→
(6)
In strongly acid solution, H
•
may even react as an oxidant. For example, H
•
oxi-
dizes Fe
2+
to Fe
3+
[reaction (7)]. A hydride, Fe
3+
H
−
, is thought to be an interme-
diate in this reaction.
Fe
2+
+ H
•
+ H
+
Fe
3+
+ H
2
→
4.3
Dissociative Electron Capture and Related Reactions
The hydrated electron reacts with many compounds which are capable of releas-
ing an anion by dissociative electron capture [e.g., reaction (8)], and, among
others, it was this property which allowed the differentiation between e
aq
−
and
H
•
[reactions (9) and (10)] (Armstrong et al. 1958; Hayon and Allen 1961; Jortner
and Rabani 1962).
e
aq
−
+ ClCH
2
CO
2
H
Cl
−
+
•
CH
2
CO
2
H
→
(8)
H
•
+ ClCH
2
CO
2
H
H
2
+
•
CHClCO
2
H
(9)
→
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