220
JACOBSEN, HOLCMAN, AND SEHESTED
Table II Activation Energies Ea, Arrhenius Preexponential Factor A, Activation Enthalpy ⌬‡HЊ and Activation
Entropy ⌬‡SЊ for Selected Ferryl Ion Reactions at 25ЊC
Ϫ1
Ϫ1
Ϫ1
Ϫ1
Ϫ
Ϫ1
Reaction
Ea (kJ mol
)
A (M
s
)
⌬‡HЊ (kJ mol
)
⌬‡SЊ (J mol 1ЊK
)
FeO2ϩ ϩ HNO2
34.5 Ϯ 0.5
21.3 Ϯ 0.3
22.3 Ϯ 0.5
44.5 Ϯ 1.0
28.1 Ϯ 0.5
(1.2 Ϯ 0.2) ϫ 1010
(5.4 Ϯ 0.5) ϫ 107
(1.3 Ϯ 0.3) ϫ 106
(2.5 Ϯ 0.2) ϫ 1010
(3.4 Ϯ 0.4) ϫ 108
32.0
18.8
19.8
42.0
25.6
Ϫ60.1
Ϫ105.2
Ϫ136.2
Ϫ54.2
ϩ
FeO2ϩ ϩ Mn2
FeO2ϩ ϩ HCOOH
FeO2ϩ ϩ CH2(OH)2
FeO2ϩ ϩ C6H5OH
Ϫ90.0
1
1
CH3CH2OH) ϭ (2.50 Ϯ 0.3) ϫ 103 MϪ sϪ . In
analogy with formic acid this behavior is rationalized
in terms of simultaneous reaction of the ferryl ions
with the primary products formed in the reaction of
FeO2ϩ with ethanol, which on increasing ethanol con-
centrations become less and less prominent.
The reactivity of OH towards the same organic
compounds have been included in Figure 6, to illus-
trate the similarity in the reaction mechanism between
ferryl ions and OH radicals.
Finally, activation parameters for selected reactions
between the ferryl ions and cloud water constituents
were investigated in 1.0 M HClO4-solutions in the
temperature range 5.0–35.0ЊC.
Reaction of FeO2 with formaldehyde was studied
ϩ
in the pH range 0–1.5 with (0.25–5.0) ϫ
10Ϫ3 M CH2(OH)2 and, [Fe2ϩ]0 ϭ 2.5 ϫ 10Ϫ5 M and
[O3]0 ϭ 2.8 ϫ 10Ϫ5 M. The plot of the observed rate
constant vs. concentrations of hydrated formaldehyde
CH2(OH)2 was linear yielding a rate constant
The activation energies and thermodynamic data
are shown in Table II.
k(FeO2ϩ ϩ CH2(OH)2) ϭ 400 Ϯ 50 MϪ sϪ1 indepen-
1
CONCLUSION
dently of the acid concentration, Figure 2.
Reaction of FeO2ϩ with Acetone. The rate constant
Rate constants and activation energies for reactions of
selected inorganic and organic compounds present in
atmospheric water with ferryl ion have been measured.
For rain water with typical acidity of about pH 3
ferryl ion play a role mainly as a temporary OH-radical
sink whereas for more acidic aerosols ferryl ion is
more likely to react as a distinct species. During its
formation, ozone (a two-electron oxidant) is consumed
and FeO2ϩ (a one-electron oxidant) is formed. The re-
sult is a loss of oxidation capacity in the aqueous phase
of the troposphere. In all iron(IV) reactions iron(III) is
formed as one of the products and this might cause
complications in terms of sulphur(IV) oxidation since
autoxidation of HSOϪ3 is initiated by Fe3ϩ [17].
ϩ
of the reaction of FeO2 with acetone was measured
ϩ
in 1.0 M HClO4 solution with [Fe2 ]0 ϭ 2.5 ϫ
10Ϫ5 M, [O3]0 ϭ 3.8 ϫ 10Ϫ5 M, and [CH3COCH3]0 ϭ
(1.0–5.0) ϫ 10Ϫ3 M at 25ЊC. The rate constant
ϩ
1
1
k(FeO2 ϩ CH3COCH3) ϭ 16 Ϯ 2 MϪ sϪ was ob-
tained for the reaction.
ϩ
Reaction of FeO2 with Phenol and Benzoic Acid.
As it was the case with formic acid and ethanol the
plot of the apparent first-order rate constant vs. phenol
concentration was curved indicating rather fast reac-
tions of the ferryl ions with secondary products of
the reaction, Figure 5. The initial slope yields
ϩ
1
1
k(FeO2 ϩ C6H5OH) ϭ (2.8 Ϯ 0.5) ϫ 104 MϪ sϪ
while from the final slope a value of almost a factor
ϩ
of ten lower was obtained k(FeO2 ϩ C6H5OH) ϭ
1
1
(4.0 Ϯ 2.0) ϫ 103 MϪ sϪ indicating that a compli-
cated chain mechanism may be operative when the
surplus of phenol over FeO2ϩ is insufficient.
Financial support from the Commission of the European
Communities within the Environment research program
(contract RINOXA EV5V-ct93-0317) is gratefully acknowl-
edged.
The rate of reaction of FeO2 with benzoic acid is
ϩ
slow and we consider the measured value k(FeO2
ϩ
ϩ
C6H5COOH) ϭ 80 Ϯ 20 MϪ sϪ1 as an upper limit for
1
BIBLIOGRAPHY
this rate constant.
For FeO2 reactions with the organic compounds
ϩ
1. T. E. Graedel, M. L. Mandich, and C. J. Weschler, J.
Geophys Res., 91, (D4), 5205 (1986).
2. T. E. Graedel and K. I. Goldberg, J. Geophys Res., 88,
(C15), 10865, (1983).
3. C. J. Weschler, T. E. Graedel, and M. L. Mandich, J.
Geophys. Res., 19, 505, (1986).
studied, rate constants in the range 3.0 to 4.5 ϫ
105 MϪ sϪ were obtained. These rate constants cor-
relate fairly well with the bond dissociation energy
(BDE) [16] of the reactants, i.e., the plot of log
1
1
k(FeO2 ) vs. BDE is linear, Figure 6.
ϩ