Reactions of the ferryl ion with some compounds found in cloud water

Article abstract of DOI:10.1002/(SICI)1097-4601(1998)30:3<215::AID-KIN7>3.0.CO;2-V

The lifetime of the ferryl ion FeO, (minutes at pH below 1) is reduced by a factor of 50 at pH above 3. This is rationalized in terms of an acid-base equilibrium between two different hydrolytic forms of this species (Fe(OH)₂⁺⁺ and F

Full text of DOI:10.1002/(SICI)1097-4601(1998)30:3<215::AID-KIN7>3.0.CO;2-V

Reactions of the Ferryl Ion
with Some Compounds
Found in Cloud Water
FRANK JACOBSEN, JERZY HOLCMAN, KNUD SEHESTED
Chemical Reactivity, bld. 313, Risø National Laboratory, DK-4000 Roskilde, Denmark
Received 24 February 1997; accepted 14 July 1997
ABSTRACT: The lifetime of the ferryl ion FeO2ϩ, (minutes at pH below 1) is reduced by a fac-
tor of 50 at pH above 3. This is rationalized in terms of an acid-base equilibrium between
two different hydrolytic forms of this species (Fe(OH)ϩ2 ϩ and Fe(OH)3ϩ) with a pKa ϭ 2.0. The
ϩ
rate constants for reactions between FeO2 and selected compounds found in cloud water
were measured in acid solutions by stopped-flow technique. For inorganic reactants:
1
1
1
1
1
1
kHNO2 ϭ 1.1 ϫ 104 MϪ sϪ
,
kNO2 Յ 105 MϪ sϪ
,
kCl ϭ 1.0 ϫ 102 MϪ sϪ
,
kHSO3 ϭ 2.5 ϫ
Ϫ
1
Ϫ
Ϫ
105 MϪ sϪ1, kSO2 ϭ 4.5 ϫ 105 MϪ sϪ1, and kMn(II) ϭ 1.0 ϫ 104 MϪ sϪ were obtained and for
1
1
1
the organic reactants: kHCOOH ϭ 160 MϪ sϪ1, kHCOO ϭ 3 ϫ 105 MϪ1, kCH3COOH ϭ 3.1 MϪ sϪ
,
1
1
1
Ϫ
1
1
1
1
1
1
kCH2(OH)2 ϭ 400 MϪ sϪ
,
kCH3COCH3 ϭ 16 MϪ sϪ
,
kCH3CH20H ϭ 2.5 ϫ 103 MϪ sϪ
,
kC6H5OH ϭ
4.0 ϫ 103 MϪ sϪ1, and kC6H5COOH ϭ 80 MϪ sϪ were obtained. A good correlation between
log(k) and the standard one electron reduction potential (EЊ) indicates that the reactions of
inorganic compounds proceed by electron transfer. The reaction mechanisms in case of or-
ganic compounds are very similar to the reactions of the OH radical, i.e., H-abstraction, and
1
1
1
a fairly good correlation between log(k) and the bond dissociation energy BDE was obtained.
ϩ
Activation parameters were measured for the reaction of FeO2
with HNO2
ϩ
(Ea ϭ 34.5 kJ/mol); Mn2 (Ea ϭ 21.3 kJ/mol); HCOOH (Ea ϭ 22.3 kJ/mol); CH2(OH)2 (Ea ϭ
44.5 kJ/mol); and C6H5OH (Ea ϭ 28.1 kJ/mol). ᭧ 1998 John Wiley & Sons, Inc. Int J Chem Kinet 30:
215–221, 1998.
INTRODUCTION
Fe(II) was identified as a night-time source of OH -
radicals via a Fenton like reaction [1], and at day-time
the direct photoreduction of Fe(III)-complexes also
generates OH and Fe(II) [4]. In an oxidizing atmo-
sphere the concentration of Fe(II) relative to Fe(III) is
surprisingly high, both due to the photoreduction and
the fact that, for typical iron concentrations of 1.8 ϫ
10Ϫ8 M in the aqueous phase of the atmosphere, HO2
reacts almost entirely with ferrous and ferric ions [5]
maintaining a steady-state concentration of iron(II). In
fact, field measurements of soluble iron(II) in fog wa-
ter have demonstrated that a large fraction of total iron
To evaluate the effect of relevant atmospheric reac-
tions, model simulations have proven to be an impor-
tant tool. However, for the model calculations to pro-
vide a high degree of accuracy reliable data has to be
available. So far most of the basic chemical reactions
in the troposphere are known for both liquid and gas-
eous state, but the role and presence of transition metal
ions (TMI) still seems to require special attention.
When TMI was included in model calculations [1–3]
ϩ
exists as Fe2 in solution [6]. Until quite recently it
was believed that the oxidation of ferrous by ozone
produced OH radicals and iron(III), however, Løgager
et al. [7] provided final evidence that iron(IV) (the
ferryl ion) and oxygen was formed.
Correspondence to: K. Sehested
Contract grant Sponsor: Commission of the European Communities
within the Environment research program
Contract grant number: (contract RINOXA EV 5V-ct93-0317)
᭧ 1998 John Wiley & Sons, Inc.
CCC 0538-8066/98/030215-07

216
JACOBSEN, HOLCMAN, AND SEHESTED
The objective of this work was to investigate
the reactivity of the ferryl ions (Fe(IV)) towards
important inorganic and organic cloud water con-
stituents, i.e., HNO2/NOϪ2 , ClϪHSO3Ϫ/SO2, Mn2ϩ,
and HCOOH/HCOOϪ, CH3COOH, CH3CH2OH,
CH3OCH3, CH2(OH)2, phenol, and benzoic acid.
itor its decay. Below pH 1.5 a decrease of absorption
was observed and at pH Ն 2 an increase of absorption
was observed due to the formation of different hydro-
ϩ
lytic forms of ferric ions, Fe3ϩ and FeOH2 .
The observed overall first-order rate constant of the
ϩ
FeO2 decay increases about fifty times when going
from pH 0 to 3. The decay of the ferryl ion is repre-
sented by a sequence of reactions, rate constants are
adopted from [7]:
EXPERIMENTAL METHODS
Materials
ϩ
ϩ
FeO2 ϩ H2O !: Fe3 ϩ OH ϩ OHϪ
1
k1a ϭ (1.3 Ϯ 0.2) ϫ 10Ϫ2 sϪ (1a)
All solutions were prepared using triply-distilled water
and the chemicals, all from Merck, were of analytical
grade, and used as supplied.
ϩ
ϩ
FeO2 ϩ OH ϩ Hϩ !: Fe3 ϩ H2O2
1
1
k1b ϭ (1.0 Ϯ 0.5) ϫ 107 MϪ sϪ (1b)
ϩ
ϩ
FeO2 ϩ H2O2 !: Fe3 ϩ HO2 ϩ OHϪ
O3 Preparation and Measurements
1
1
k1c ϭ (1.0 Ϯ 0.1) ϫ 104 MϪ sϪ (1c)
Stock solution of ozone was prepared by bubbling the
effluent from an Erwin Sander (V) ozonator through
an aqueous HClO4 solution. The ozone concentration
was determined spectrophotometrically at its maxi-
ϩ
ϩ
FeO2 ϩ HO2 !: Fe3 ϩ O2 ϩ OHϪ
1
1
k1d ϭ (2.0 Ϯ 1.0) ϫ 106 MϪ sϪ (1d)
mum absorption wavelength, ␭ ϭ 260 nm, using
max
At low pH all reactions are significantly taking part in
the decay, while only two reactions (1a–1b) are in-
volved at high pH because at micromolar concentra-
tions of FeO2ϩ reaction (1c) is too slow to significantly
influence the observed rate. At all pH’s the rate deter-
mined step is reaction (1a), and the observed rate con-
stant is plotted as a function of pH in Figure 1. We
interpret the s-shaped curve as a result of a protolytic
equilibrium between two different hydrolytic forms of
an extinction coefficient ⑀260 ϭ 3300 MϪ cmϪ [8].
1
1
Hydrogen peroxide was measured by the ferrous
ϩ
method [9], using the Fe3 304 nm absorption
1
1
(⑀304 ϭ 2200 MϪ cmϪ ).
Stopped-flow Apparatus
For the kinetic studies a Hi-Tech SF 51 Stopped-flow
spectrophotometer (SFS) with UV and VIS lamps was
used. The data was recorded on a LeCroy 9400 A dual
175 MHz digital oscilloscope and analyzed on an on-
line IBM PC equipped with software for data analysis.
The optical quartz cell, of 10 mm optical path length,
and the drive-syringes were kept at constant temper-
atures between 5.0–35.0ЊC within Ϯ0.2ЊC. The
stopped-flow apparatus was equipped with two point
mixing facility allowing 10–20 ms delay time after
the first mixing.
RESULTS AND DISCUSSION
؉
Influence of Acidity on the Decay of FeO2
The hydrolysis of the ferryl ions was studied
in 1.0–10Ϫ3 M HClO4 with initial concentrations
ϩ
[Fe2 ]0 ϭ 5.0 ϫ 10Ϫ5 M and [O3]0 ϭ 6.0 ϫ 10Ϫ5 M.
The ferryl ion was formed by oxidation of Fe2ϩ by O3
in acid solution [7,10]. At pH 0–1 the half-life of
ϩ
FeO2 is in the order of minutes, but decreases with
Figure 1 pH dependence of the rate constant k1a for
ϩ
ϩ
Ϫ5
FeO2 Initial concentrations; [Fe2 ]0 ϭ 5.0 ϫ 10 M and
increasing pH.
The 320 nm absorption of FeO2ϩ was used to mon-
[O3]0 ϭ 6.0 ϫ 10 M in perchloric acid solutions at 25ЊC.
Ϫ5

REACTIONS OF FERRYL ION
217
ϩ
iron (IV) (possibly Fe(OH)2ϩϩ L Fe(OH)3 ) with
pKa around 2.0.
concentrations would prevent the formation of ferryl
Ϫ
ions as ozone reacts fast with NOϪ2 , k(O3 ϩ NO2 ) ϭ
1
3.7 ϫ 105 MϪ sϪ1 [12].
The reaction of FeO2ϩ with HNO2, whether it goes
via H-abstraction or electron transfer, results in for-
ϩ
Reactions of FeO2 with Inorganic
Compounds
ϩ
mation of Fe3 and NO2 (reaction 2). From our ex-
ϩ
Ϫ
Reaction of FeO2 with HNO2 /NO2 . Reaction of
periments we are unable to deduce whether NO2 reacts
with the ferryl ion but this may be likely if the rate of
FeO2 with nitrous acid was studied in 1.0 and
ϩ
such a reaction exceeds 106 MϪ sϪ . If this is the case
the observed rate constant k2 should be divided by 2.
The final products of the reaction between HNO2 and
FeO2ϩ are Fe3ϩ and NO3Ϫ ions.
1
1
0.1 M HClO4 with initial concentrations of [Fe2 ]0 ϭ
ϩ
2.5 ϫ 10Ϫ5 M and [O3]0 ϭ (2.5–4.0) ϫ 10Ϫ5 M and
[HNO2] ϭ (1.2–5.0) ϫ 10Ϫ4 M.
ϩ
ϩ
Reaction of FeO2 with ClϪ. Reaction of FeO2
FeO2 ϩ HNO2 !: Fe3 ϩ NO2 ϩ OHϪ (2)
ϩ
ϩ
with chloride was studied in 0.1 M HClO4 with
ϩ
(3–10) ϫ 10Ϫ3 M ClϪ. The rate constant k(FeO2
ϩ
The decay of the ferryl ion was followed at 320 nm.
From the plot of the apparent first-order rate con-
stant vs. HNO2 concentration, k2(FeO2 ϩ HNO2) ϭ
(1.1 Ϯ 0.1) ϫ 104 MϪ sϪ was obtained, Figure 2.
The slope of the straight line in the first-order plot is
identical for both HClO4 concentrations allowing to
estimate the upper limit of the rate constant for the
reaction between ferryl ion and the nitrite ion to
k(FeO2 ϩ NOϪ2 ) Յ 105 MϪ sϪ1 using pKa ϭ 3.3 for
nitrous acid [11]. It was not feasible to measure at
pH Ͼ 1, as larger HNO2/NOϪ2 concentrations are nec-
essary to compete with a faster decay of the ferryl ion
(the conversion rate of ferryl ions into OH radicals
(reaction (1a)) constitutes the lowest limit for the ob-
ClϪ) ϭ (100 Ϯ 10) MϪ sϪ1 was obtained from a plot
of the observed first-order rate of decay at 320 nm vs.
ClϪ concentration, Figure 2. It is unclear whether the
1
ϩ
1
1
reaction FeO2 ϩ ClϪ is a real electron transfer reac-
ϩ
tion because of the extиremely hicgohuopnlee eElЊe(cиtCroln/Crled)uϭc-
Ϫ
tion potential of the Cl/ClϪ
2.41 V [13].
ϩ
1
If this is the case the one electron reduction poten-
ϩ
tial of the FeO2 /Fe3ϩ couple should be in the vicinity
of 2.0 V. It is however possible that the ClϪ complexes
ϩ
FeO2 prior to the reaction so that Cl2Ϫ is formed di-
rectly whereby a substantially lower reduction poten-
tial EЊ(ClϪ2 /2ClϪ) ϭ 2.1 V [13] has to be overcome.
ϩ
Ϫ
Ϫ
Reaction of FeO2 with HSO3 and SO2. The
served rate constant). However, larger HNO2/NO2
reaction of FeO2 with bisulphite was followed
ϩ
at 320 nm, by observing the build-up of Fe3
ϩ
(as Fe(OH)2 and/or Fe(OH)2ϩ) in 1.0 ϫ 10Ϫ3 M
ϩ
HClO4 solutions at 25ЊC. At this pH more than
90% of the bisulphite is in the HSOϪ3 form based
on pKa ϭ 1.81 [11]. With the initial concentra-
ϩ
tions
applied
[Fe2 ]0 ϭ [O3]0 ϭ 5.0 ϫ 10Ϫ5 M,
FeO2 (2.5 ϫ 10Ϫ5 M) reacts with (1.25–5.0) ϫ
10Ϫ5 M HSO3 . Under these nonpseudo-first-order
ϩ
Ϫ
conditions only the initial rate method could be used
to determine the rate constant. The apparent rate con-
stant was measured within the first 100 ms of the re-
action and kapp was a linear function of the bisulphite
Ϫ
concentration, [HSO3 ], and a second-order rate con-
ϩ
stant for the reaction was evaluated: k(FeO2
ϩ
HSO3 ) ϭ (2.5 Ϯ 1.0) ϫ 105 MϪ sϪ . At pH 1 more
Ϫ
1
1
than 90% of the bisulphite is in the SO2 form and the
ϩ
rate constant k(FeO2 ϩ SO2) was determined to
ϩ
1
1
k(FeO2 ϩ SO2) ϭ (4.5 ϩ 1.0) ϫ 105 MϪ sϪ
by
observing the 320 nm decay of the ferryl ion.
ϩ
ϩ
Reaction of FeO2 with Mn2 . Reaction between
ϩ
Figure 2 Dependence of the observed rate of FeO2 decay
FeO2ϩ and Mn2ϩ was studied in 1.0 M HClO4 at 25ЊC.
on the concentration of: chloride, [Cl 2] ϫ (5 ϫ 104) M,
Ϫ
Equal concentrations of Fe2 and O3 were used
ϩ
(ٗ) pH 1; nitrous acid, [HNO2] ϫ 104 M, (O) pH 0, (᭝) pH
1.5; and formaldehyde, [CH2(OH)2] ϫ 103 M, (᭞) pH 0,
(᭛) pH 1,5.
([Fe2 ]0 ϭ [O3]0 ϭ 5.0 ϫ 10Ϫ5 M) and the man-
ϩ
ganese concentration was kept in the range

218
JACOBSEN, HOLCMAN, AND SEHESTED
(0.5–2.0) ϫ 10Ϫ4 M. Manganic and ferric ions, iden-
tified by resulting absorption spectra [14], were
formed as products according to:
reaction (1a) is spontaneous (⌬rG Յ 0), a lower limit
for the standard one electron reduction potential for
ϩ
ϩ
the (FeO2 /Fe3 ) couple can be estimated to
ϩ
ϩ
EЊ(FeO2 /Fe3 ) Ն 0.87 V. Greenwood and Earnshaw
[15] suggested a standard one electron reduction po-
tential of a Fe(IV) compound to be in the order of
2 V, sufficient for ClϪ to be oxidized by electron trans-
fer as indicated by the linear relationship between
log(k) and standard one electron reduction potential,
Figure 3.
ϩ
ϩ
FeO2 ϩ Mn2 ϩ 2Hϩ !:
ϩ
ϩ
Fe3 ϩ Mn3 ϩ H2O (3)
The plot of the apparent first-order decay rate vs.
ϩ
[Mn2 ] was linear yielding
a
rate constant
k(FeO2ϩ ϩ Mn2 ) ϭ (1.0 Ϯ 0.1) ϫ 104 MϪ sϪ . This
value is about an order of magnitude lower than
the value measured for the ferryl ion reactions with
Fe2ϩ [7,10].
ϩ
1
1
؉
Reaction of FeO2 with Organic
Compounds
In general the rates of FeO2 reactions with inor-
ϩ
ϩ
ganic substrates show the characteristics of electron
transfer reactions by the linear relationship between
log(k) as a function of the standard one electron re-
duction potential EЊ [13], Figure 3.
Reaction of FeO2 with HCOOH/HCOOϪ. Reac-
ϩ
tion of FeO2 with formic acid was studied in 1.0 M
ϩ
HClO4 with [Fe2 ]0 ϭ (2.1–2.5) ϫ 10Ϫ5 M, [O3]0 ϭ
(2.4–2.5) ϫ 10Ϫ5 M and (0.125–15.0) ϫ 10Ϫ3 M
HCOOH. The plot of the observed first-order rate con-
stant as a function of formic acid concentration, Figure
4, is nonlinear.
From this linear relationship one would predict a
ϩ
ϩ
value of log k(FeO2 ϩ Fe2 ) to be about 5.4 for the
reaction between ferryl and ferrous ions. The reported
ϩ
ϩ
values in the literature are log k(FeO2 ϩ Fe2 ) ϭ
The bimolecular rate constant obtained from the
5.16 [7] and log k(FeO2 ϩ Fe2 ) ϭ 5.26 [10]. In the
initial slope is k(FeO2 ϩ HCOOH) ϭ 350 MϪ sϪ
ϩ
ϩ
ϩ
1
1
case of hydrogen peroxide the predicted value of log
about twice of the value obtained from the final slope.
This can be rationalized in terms of a simultaneous
reaction of the ferryl ions with HO2 radicals formed
as the product of the one-electron oxidation reaction
of formic acid and a subsequent reaction with oxygen
ϩ
k(FeO2 ϩ H2O2) is 4.0 and the values of
ϩ
log k(FeO2 ϩ H2O2) in the literature are; 4.0 [7] and
3.98 [10]. The literature values of rate constants of
FeO2ϩ reactions with Fe2ϩ and H2O2 has been included
in Figure 3 for comparison. On the assumption that
ϩ
ϩ
Figure 3 log k(FeO2 ) as a function of the standard one
Figure 4 Dependence of the observed rate of FeO2 decay
electron reduction potential (E0) for the electron transfer re-
action of FeO2 with inorganic compounds. (᭹) literature
values [7,10] (the point for NO2 also includes HNO2).
on (᭹) the formic acid concentration and (O) the ethanol
concentration at 1.0 M HClO4 and 25ЊC; [Fe2 ]0 ϭ
ϩ
ϩ
Ϫ
Ϫ5
[O3]0 ϭ 2.5 ϫ 10 M.

REACTIONS OF FERRYL ION
219
ϩ
FeO2 ϩ HCOOH ϩ Hϩ !:
ϩ
Fe3 ϩ COOH ϩ H2O
(4)
(5)
и
и
COOH ϩ O2 !: CO2 ϩ HO2
FeO2 ϩ HO2 ϩ Hϩ !: Fe3 ϩ O2 ϩ H2O (6)
ϩ
ϩ
HO2 ϩ HO2 !: H2O2 ϩ O2
(7)
At higher concentrations of formic acid, reaction (4)
competes efficiently with reaction (6) whereby the fi-
nal slope represents the rate constant of the elementary
ϩ
1
1
reaction (4), k(FeO2 ϩ HCOOH) ϭ 160 MϪ sϪ .
Thus, at high HCOOH concentrations the forma-
tion of H2O2 by reaction (7) is enhanced. This for-
mation of hydrogen peroxide was experimentally ver-
ified by the ferrous–ferric oxidation [9], although the
determination was only semi-quantitative because of
the presence of HCOOH. From the dependence of the
ϩ
apparent second-order rate constant kobs(FeO2
ϩ
ϩ
Figure 5 Dependence of the observed rate of FeO2 decay
HCOOH/HCOOϪ) on pH in the range 0.0–3.0, Table
I, a rate constant for the reaction of the ferryl ion
on phenol concentration in 1.0 M perchloric acid solution at
ϩ
Ϫ5
25ЊC with [Fe2 ]0 ϭ 2.5 ϫ 10 M and [O3]0 ϭ 2.8 ϫ
10 M.
with the formate ion k(FeO2 ϩ HCOOϪ) ϭ (3.0 Ϯ
Ϫ5
ϩ
1.0) ϫ 105 MϪ sϪ was deduced using pKa ϭ 3.75
1
1
for the formic acid dissociation.
ϩ
Reaction of FeO2 with CH3COOH. The rate
ϩ
constant of the reaction of FeO2 with acetic acid
nonlinear function of the ethanol concentration, Figure
ϩ
4. From the initial slope a rate constant k(FeO2
ϩ
was measured in 1.0 M HClO4 solution with
1
1
CH3CH2OH) ϭ (5.0 Ϯ 0.5) ϫ 103 MϪ sϪ was ob-
[Fe2 ]0 ϭ 2.5 ϫ 10Ϫ5 M, [O3]0 ϭ 3.8 ϫ 105 M and
ϩ
[CH3COOH]0 ϭ (5.0 Ϫ 10.0) ϫ 10Ϫ2 M at 25ЊC.
tained, while the final slope, as in the case of for-
ϩ
mic acid, yields half of this value k(FeO2
ϩ
The rate constant k(FeO2 ϩ CH3COOH) ϭ 3.1 Ϯ
ϩ
0.5 MϪ sϪ1 was obtained. This value is considered to
1
be merely an upper limit since for such a low rate
constant reasonable doubts exist as to whether the
measured value is pertinent to the reaction in question
or is due to a minute amount of a highly reactive im-
purity.
ϩ
Reaction of FeO2 with Ethanol and Formal-
ϩ
dehyde. The reaction of FeO2 with ethanol
was studied in 1.0 M HClO4 solutions at
25ЊC with [CH3CH2OH]0 ϭ (1.0–5.0) ϫ 10Ϫ4 M,
ϩ
[Fe2 ]0 ϭ 2.5 ϫ 10Ϫ5 M,
and
[O3]0 ϭ 2.5 ϫ
10Ϫ5 M. The observed first-order rate constant was a
Table I Apparent Second-Order Rate Constants for
Ferryl Ion Reactions with Formic Acid at Different pH
Ϫ2
Ϫ1
Ϫ1
pH
kobs ϫ 10
M
s
0.0
1.0
1.5
1.8
3.0
1.6
5.8
16.0
23.0
ϩ
Figure 6 log k(FeO2 ) and log k(OH) for the H-abstraction
reactions of FeO2 and OH radicals as a function of the bond
dissociation energy, BDE, for the organic compounds stud-
ied.
ϩ
750.0

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
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For FeO2 reactions with the organic compounds
ϩ
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(BDE) [16] of the reactants, i.e., the plot of log
1
1
k(FeO2 ) vs. BDE is linear, Figure 6.
ϩ

REACTIONS OF FERRYL ION
221
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