Time-resolved studies of the kinetics of the reactions of CHO with HI and HBr: Thermochemistry of the CHO radical and the C-H bond strengths in CH2O and CHO

Article abstract of DOI:10.1021/jp970443y

Time-resolved studies of formyl radical, CHO, generated by laser flash photolysis of acetaldehyde, have been carried out to obtain rate constants for its bimolecular reactions with HI and HBr. Each reaction was studied in the gas-phase at five different temperatures in the range 293-540 K. The pressure-independent rate constants gave the following Arrhenius equations: for CHO + HI, log(k₂/cm³ molecule^-1 s^-1) = (-11.51 ± 0.01) - (0.83 ± 0.06 kJ mol^-1)/ RT ln 10, for CHO + HBr, log(k₄/cm³ molecule^-1 s^-1) = (-12.59 ± 0.06) - (0.13 ± 0.44 kJ mol^-1)/ RT ln 10. The Arrhenius parameters are consistent with those for reactions of other free radicals with HI and HBr and support the idea of intermediate complexes in these processes. Combination with the known kinetic data on the reverse reactions leads to the thermodynamic estimate for ΔH_f^○(CHO) = 44.29 ± 0.43 kJ mol^-1 (third law method, 298 K value). Second law values are in reasonable agreement. This value corresponds to the bond dissociation energies D(H-CHO) = 370.86 ± 0.28 kJ mol^-1, and D(H-CO) = 63.18 ± 0.46 kJ mol^-1. These values are in close agreement with several previous determinations although more precise. A small unresolved discrepancy exists with the most recent value for the photodissociation threshold for CH₂O.

Full text of DOI:10.1021/jp970443y

J. Phys. Chem. A 1997, 101, 4185-4190
4185
Time-Resolved Studies of the Kinetics of the Reactions of CHO with HI and HBr:
Thermochemistry of the CHO Radical and the C-H Bond Strengths in CH2O and CHO
Rosa Becerra
Instituto de Quimica-Fisica “Rocasolano”, C.S.I.C. C/Serrano 119, 28006 Madrid, Spain
Ian W. Carpenter and Robin Walsh*
Department of Chemistry, UniVersity of Reading, Whiteknights, PO Box 224, Reading RG6 6AD, UK
ReceiVed: February 5, 1997; In Final Form: March 27, 1997X
Time-resolved studies of formyl radical, CHO, generated by laser flash photolysis of acetaldehyde, have
been carried out to obtain rate constants for its bimolecular reactions with HI and HBr. Each reaction was
studied in the gas-phase at five different temperatures in the range 293-540 K. The pressure-independent
rate constants gave the following Arrhenius equations: for CHO + HI, log(k2/cm3 molecule-1 s-1) ) (-11.51
( 0.01) - (0.83 ( 0.06 kJ mol-1)/ RT ln 10, for CHO + HBr, log(k4/cm3 molecule-1 s-1) ) (-12.59 (
0.06) - (0.13 ( 0.44 kJ mol-1)/ RT ln 10. The Arrhenius parameters are consistent with those for reactions
of other free radicals with HI and HBr and support the idea of intermediate complexes in these processes.
Combination with the known kinetic data on the reverse reactions leads to the thermodynamic estimate for
∆Hf°(CHO) ) 44.29 ( 0.43 kJ mol-1 (third law method, 298 K value). Second law values are in reasonable
agreement. This value corresponds to the bond dissociation energies D(H-CHO) ) 370.86 ( 0.28 kJ mol-1,
and D(H-CO) ) 63.18 ( 0.46 kJ mol-1. These values are in close agreement with several previous
determinations although more precise. A small unresolved discrepancy exists with the most recent value for
the photodissociation threshold for CH2O.
Introduction
threshold for onset of photodissociation of CH2O, using the
PHOFEX technique to monitor H atoms.
The formyl radical, CHO, is of importance as a reactive
intermediate in both atmospheric chemistry1 and combustion.2
Reliable values for ∆Hf°(CHO) and the related bond dissociation
energy D(H-CHO) are necessary in order to assess the kinetics
of its elementary reactions. Recent years have seen a gradual
refinement of the values of heats of formation of many organic
free radicals through the application of a variety of techniques3
such that, in many cases, uncertainties have diminished to as
low as (2 kJ mol-1. As far as CHO is concerned, our interest
in this radical began in 1966 when one of the present authors
was involved in an earlier gas phase study4 of the kinetics of
the reaction:
Despite the apparent precision of the current value, we feel
that the importance of this quantity makes the experimental
check of its value worthwhile. This is only true if the technique
can yield a value of equivalent precision, which is the case with
the work described here. Direct kinetic studies of reaction 2
offer this possibility since a combination of the rate constants
of reactions 1 and 2, and their temperature dependencies, can
be used to obtain ∆Hf° for the reaction and thereby ∆Hf°(CHO).
This would eliminate the assumption made earlier4 and lead to
the desired precision. The key to obtaining a reliable value is
to ensure that the temperature ranges of study of (1) and (2)
are as close as possible and to check that the thermodynamic
data derived using both second and third law treatments are in
close agreement. This approach follows closely that of Gutman
and co-workers,13 who have studied the kinetics of many
radicals, R, with both HI and HBr and used the results to obtain
∆Hf°(R). The kinetics of reaction (2) have not previously been
studied.
I + CH2O f CHO + HI
(1)
which led to a value of 30 ( 8 kJ mol-1 for ∆Hf°(CHO). This
determination was by no means the first value of this quantity,
but it was a landmark in the sense that it served at the time to
eliminate some earlier estimates as much as 50 kJ mol-1 lower
in value and also threw some light on the photochemistry of
CH2O. The weakness in this determination was that it required
an assumption about the activation energy of reaction 2, the
reverse of reaction 1:
Following earlier time-resolved investigations of CHO radical
reactions,14,15 we describe here studies in which CHO is created
via excimer laser flash photolysis of acetaldehyde and monitored
2
2
via its X˜ A′ f A˜ A′′ vibronic transition. As well as reaction
2 with HI, the reaction of CHO with HBr, reaction 4, is also
investigated here in order to gain a second handle on the desired
thermochemistry by combination with the known rate data16,17
for reaction 3, viz.,
CHO + HI f CH2O + I
(2)
In the intervening years there have been several determinations
of ∆Hf°(CHO) employing a variety of experimental tech-
niques.5-12 The values tend to cluster in the range 40-45 kJ
mol-1, mostly with uncertainties of at least (4 kJ mol-1. In
the most recent study Chuang, Foltz, and Moore12 have obtained
a value of 41.8 ( 0.8 kJ mol-1 from a measurement of the
Br + CH2O f CHO + HBr
CHO + HBr f Br + CH2O
(3)
(4)
The kinetics of reaction 4 have also not been studied previously.
Quite apart from the thermochemical application, reactions
of CHO radicals with HI and HBr are of interest because the
X Abstract published in AdVance ACS Abstracts, May 15, 1997.
S1089-5639(97)00443-X CCC: $14.00 © 1997 American Chemical Society

4186 J. Phys. Chem. A, Vol. 101, No. 23, 1997
Becerra et al.
study of their kinetics can throw further light on the possible
involvement of intermediate complexes in radical reactions with
hydrogen halides.18-26
Experimental Section
The apparatus and equipment for these studies have been
described in detail previously.14,15 Only essential and brief
details are therefore included here. CHO was produced by the
308 nm flash photolysis of acetaldehyde using an Oxford Lasers
KX2 exciplex laser operating on XeCl emission. CHO con-
centrations were monitored in real time by means of a Coherent
699-21 single mode dye laser pumped by an Innova 90-5 argon
ion laser and operating with Kiton Red. Experiments were
carried out in a variable temperature quartz reactor with
demountable windows.27 The monitoring laser beam was
multipassed 32 or 36 times through the reaction zone to give
an effective path length of up to 1.5 m.
The monitoring laser was tuned to 614.66 nm, corresponding
to a position slightly to the red of the A˜ 2A′′(0,9,0) r X˜ 2A′-
(0,0,0) bandhead at 614.38 nm.28 Small variations in wave-
number (within a spread of 0.5 cm-1 ) were shown to make no
difference to the results. Light signals were measured by a dual
photodiode/differential amplifier combination, and the signal
decays were stored in a transient recorder (Datalab 910)
interfaced to a BBC microcomputer. This was used to average
decays of up to 40 laser shots (at a repetition rate of 1 or 2 Hz).
The averaged decay traces were processed by fitting the data
to an exponential form using a nonlinear least-squares package.
This analysis provided the values for the first-order rate
constants, kobs, for removal of CHO in the presence of known
partial pressures of substrate gas.
Gas mixtures for photolysis were made up, containing 5 or
10 Torr of acetaldehyde, together with appropriate pressures of
added substrate gas (either 1-5 Torr of HI or 10-50 Torr of
HBr). In a few experiments SF6 was added to test for pres-
sure dependence. Pressures were measured by capacitance
manometers (MKS Baratron). The gases used in this work
were obtained as follows. Acetaldehyde was from Aldrich
(99%). Hydrogen iodide was obtained by dehydration of
aqueous hydriodic acid (Aldrich, 99.99% HI in water) using
phosphoros pentoxide. Hydrogen bromide was from Fluka
(purum grade). Sulfur hexafluoride (no GC detectable impuri-
ties) was from Cambrian Gases. All gases were degassed
thoroughly prior to use. Additionally, the hydrogen halides
were distilled at low pressure periodically to maintain their
purity from contamination by decomposition to halogen. They
were only used after checking they were pure white when
condensed.
Figure 1. Second-order plots: examples of the dependence of the decay
constants, kobs, on pressure of (a) HI and (b) HBr. The inset photograph
shows the decay trace for the [HI] ) 3.0 Torr point at 294 K; the vertical
scale shows absorbance, and the time base corresponds to 6.4 µs
between divisions.
CHO partial pressure and photolysis laser energy (50-100 mJ/
pulse), the important point was established that, in the presence
of reactive substrate, the decays became exponential and were
not sensitive to these parameters. An example of such a decay
trace is shown, inset, in Figure 1a. Signal size was dependent
on CH3CHO pressure, and higher pressures were needed at
higher temperatures, because the CHO detection sensitivity
decreases with increasing temperature. Signal decay constants
in the presence of HI or HBr were insensitive to photolysis
repetition rates and the number of photolysis shots, except for
slight increases with number of shots at high HI and high
temperature.
For both HI and HBr, a series of experiments were carried
out at each of five temperatures in the range 293-540 K. At
each temperature five or six runs, usually of 10 shots each, were
carried out at different substrate pressures. A blank in the
absence of substrate was also performed at the same excimer
laser energy, and the decay was treated as a first-order process
for fitting purposes. These experiments showed good linear
dependencies of kobs with partial pressure of substrate, as
expected for second-order kinetics. This is illustrated in Figure
1 for reaction of CHO with both HI and HBr. The second-
order rate constants, obtained by least-squares fitting to these
plots, are collected in Table 1. The rate constants showed no
variation with added SF6. The error limits shown are single
standard deviations. Arrhenius plots of the rate constants for
Results
Earlier work has demonstrated that the 308 nm photolysis of
CH3CHO provides a reliable source of CHO radicals,14,15 via
the reaction
CH3CHO + hν f CH3 + CHO
(5)
The presence of CH3 radicals does not perturb the kinetics of
CHO decay in the presence of reactive substrates. Preliminary
experiments demonstrated that signal decay in the absence of
added substrate was not first order, as expected. Estimates from
the known extinction coefficient,14 suggested initial CHO
concentrations were of the order of 3 × 1014 molecules cm-3
.
Because of our previous study,14 we did not seek to remeasure
the CHO radical recombination constant. Although decay rates
in the absence of added substrate were sensitive to initial CH3-

Reactions of CHO with HI and HBr
J. Phys. Chem. A, Vol. 101, No. 23, 1997 4187
TABLE 2: Arrhenius Parameters for Reactions 1-4
log(A/cm3
reaction
T range/K molecule-1 s-1
)
Ea/kJ mol-1
ref
I + CH2O
CHO + HI
453-574
-9.86 ( 0.04 72.93 ( 0.38 4
294-540 -11.51 ( 0.09
0.83 ( 0.56 this work
6.24 ( 0.47 16
Br + CH2O 223-480 -10.84 ( 0.09
CHO + HBr 293-535 -12.59 ( 0.06
0.13 ( 0.44 this work
TABLE 3: Thermodynamic Quantities for Species Involved
in the Reactions of Interest
species
∆Hf°/kJ mol-1
S°/J K-1 mol-1
CH2O
CHO
Ib
-108.57 ( 0.46a
218.95b
224.65b
180.79
206.59
175.02
198.70
see text
106.76 ( 0.04
26.36 ( 0.21
111.86 ( 0.06
-36.44 ( 0.17
HIb
Brb
HBrb
a From ref 29. b From ref 10.
Figure 2. Arrhenius plots of the rate constants of reaction for CHO
with HI and HBr.
the only problem is that of correcting them to 298 K. For this
purpose midpoint temperatures of the overlapping ranges were
chosen: viz 500 K for reaction pair (1), (2) and 385 K for
reaction pair (3), (4). These give
TABLE 1: Experimental Rate Constants for the Reactions
of Formyl
CHO + HI
CHO + HBr
k/10-12 cm3
k/10-13 cm3
∆H°1,2(500 K) ) (72.93 ( 0.38) - (0.83 ( 0.56) )
72.10 ( 0.68 kJ mol-1 (i)
T/K
molecule-1 s-1
T/K
molecule-1 s-1
294
361
410
456
540
2.21 ( 0.07
2.34 ( 0.12
2.43 ( 0.06
3.09 ( 0.07
3.86 ( 0.17
293
365
414
471
535
2.38 ( 0.09
2.55 ( 0.08
2.31 ( 0.08
2.65 ( 0.11
2.38 ( 0.11
The tabulated enthalpy functions10 H°(T) - H°(298 K) for I,
CH2O, HI, and CHO are used to obtain ∆(H°(500 K) -
H°(298 K)) ) 1.17 kJ mol-1 and thereby ∆H°1,2(298 K) ) 70.93
( 0.68 kJ mol-1. Combination of this value with the known
enthalpies of formation10,29 of I, HI, and CH2O, shown for
convenience in Table 3, leads to ∆Hf°(CHO) ) 42.76 ( 0.85
both reactions are shown in Figure 2. For reaction 4, CHO +
HBr, the plot is linear with an almost zero gradient. For reaction
2, CHO + HI, the plot shows some curvature. We chose to
put the Arrhenius line through the lower temperature points and
make the assumption that the rate constants at 456 and 540 K
were slightly too high. This can be justified as a perturbation
arising from unavoidable formation of small quantities of I2 at
these temperatures. (Use of high pressures of HI led to
curvature in the second-order plot at 540 K.) The Arrhenius
parameters obtained from fitting the data of Figure 2 are
kJ mol-1
.
∆H°3,4(385 K) ) (6.24 ( 0.47) - (0.13 ( 0.44) )
6.11 ( 0.64 kJ mol-1 (ii)
The tabulated enthalpy functions10 H°(T) - H°(298 K) for Br,
CH2O, HBr, and CHO are used to obtain ∆(H°(385 K) -
H°(298 K)) ) 0.58 kJ mol-1 and thereby ∆H°3,4(298 K) ) 5.53
( 0.64 kJ mol-1. Combination of this value with the known
enthalpies of formation10,29 of Br, HBr, and CH2O, shown for
convenience in Table 3, leads to ∆Hf°(CHO) ) 45.27 ( 0.81
log(k2/cm3 molecule-1 s-1) )
(-11.51 ( 0.01) - (0.83 ( 0.06 kJ mol-1)/RT ln 10
kJ mol-1
.
log(k4/cm3 molecule-1 s-1) )
The reaction entropy changes obtained using ∆S° ) R ln-
(Af/Ar) with correction from reaction temperature to 298 K via
use of tabulated Cp values10 are
(-12.59 ( 0.06) - (0.13 ( 0.44 kJ mol-1)/RT ln 10
For reasons given later (see Discussion) we believe that the error
limits for A2 and Ea,2 should be 10 times larger, and therefore
increased uncertainties are used in the calculations below.
∆S°1,2(298 K) ) 28.58 ( 1.89 J K-1 mol-1
(cf. 31.50 from Table 3)
∆S°3,4(298 K) ) 31.79 ( 2.07 J K-1 mol-1
Thermochemical Calculations
The availability of temperature-dependent rate constants for
both reaction pairs (1), (2) and (3), (4) in substantially
overlapping temperature ranges offers the opportunity for
calculation of ∆Hf°(CHO) using both second and third law
procedures for each reaction pair. To facilitate the presentation,
Arrhenius parameters for all four reactions are given in Table
2. For reaction 3 we have chosen to use the data of Nava et
al.,16 rather than that of Poulet et al.,17 because it covers a wider
temperature range.
A. Second Law Determination of the Enthalpy of Forma-
tion of Formyl. In this procedure enthalpy changes for these
reactions are obtained directly from the differences of Arrhenius
activation energies for the forward and reverse reactions, and
(cf. 29.38 from Table 3)
The small differences between the entropies obtained from the
experiments and those from statistical calculations10 lead one
to believe that the values of ∆Hf°(CHO) obtained must be close
to the true value which will lie in the range 42-45 kJ mol-1
.
B. Third Law Determination of the Enthalpy of Forma-
tion of Formyl. In this procedure the Gibbs free energy changes
for these reactions are obtained from the equilibrium constants,
themselves calculated from the ratio of forward and reverse rate
constants (i.e., k1/k2 or k3/k4) at the midpoint temperatures, and
these are then corrected to 298 K using the free energy functions
(G°(T) - H°(298 K))/T.

4188 J. Phys. Chem. A, Vol. 101, No. 23, 1997
Thus, for (1) and (2) at 500 K
Becerra et al.
TABLE 4: Comparisons of Arrhenius Parameters for
Radical Reactions with HI and HBr
K ) k1/k2 ) (1.32 ( 0.11) × 10-6
∆G° ) 56.28 ( 0.35 kJ mol-1
A/10-12 cm3
radical, R
molecule-1 s-1
Ea/kJ mol-1
ref
(a) R + HI
CH3
4.5 ( 0.8
4.5 ( 0.9
3.9 ( 0.8
3.1 ( 0.6
3.1 ( 0.6
-1.2 ( 0.6
-3.2 ( 0.6
-5.1 ( 0.7
-6.3 ( 0.8
+0.8 ( 0.6
18
18
18
18
Then from the tabulated functions10 (G°(T) - H°(298 K))/T for
I, CH2O, HI, and CHO, we obtain
C2H5
i-C3H7
t-C4H9
CHO
∆(G°(500 K) - H°(298 K))/T ) -32.25 J K-1 mol-1
this work
(b) R + HBr
and therefore
CH3
C2H5
i-C3H7
t-C4H9
CHO
1.57 ( 0.26
1.70 ( 0.55
1.58 ( 0.38
1.37 ( 0.47
0.26 ( 0.04
0.64 ( 0.36
-1.6 ( 0.6
-4.2 ( 1.2
-6.4 ( 0.9
-7.8 ( 1.4
+0.1 ( 0.4
-4.5 ( 1.5
19
19
19
19
this work
20
∆(G°(500 K) - H°(298 K)) ) -16.13 kJ mol-1
This gives
CH3CO
∆H°(298 K) ) 56.28 + 16.13 ) 72.41 ( 0.35 kJ mol-1
times, and so the results should not be affected by diffusion or
reaction vessel surface effects. In another type of check, tests
showed that there was insignificant sample depletion by
repetitive laser photolysis (up to 40 shots). The only problem
encountered was an apparent enhancement in decay constants
for HI left in the vessel at high temperatures and pressures. In
the extreme, it was clear that rates were too fast, since at high
[HI] (>4 Torr at 540 K), the decay constant was clearly higher
than the second-order plot through lower [HI] values suggested.
The most plausible explanation for this was the formation of
small quantities of I2 in the cell at the highest two temperatures.
The rate constant for the reaction of CHO + I2 is ca. 20 times
greater than that for CHO + HI.4 Thus, as little as 1% of I2
will cause a 20% increase in decay constant. We thus suspect
that, despite our efforts to purify HI rigorously, small quantities
of I2 are the cause of the apparently high rate constants for CHO
+ HI at 456 and 540 K. This was the reason for not including
these constants in the Arrhenius fit. The consistency of the
rate constants at the other three temperatures is remarkable.
However, in view of this problem the uncertainties in the
Arrhenius parameters must be considered larger than those
quoted. Hence, we have increased these error limits by an order
of magnitude; i.e., A is uncertain to 10(0.09 cm3 molecule-1 s-1
Combination of this value with the known enthalpies of
formation10,29 of I, HI, and CH2O (Table 3) leads to ∆Hf°(CHO)
) 44.24 ( 0.62 kJ mol-1
.
Also, for (3) and (4) at 385 K
K ) k3/k4 ) 8.34 ( 0.85
∆G° ) -6.79 ( 0.32 kJ mol-1
Then from the tabulated functions10 (G°(T) - H°(298 K))/T for
Br, CH2O, HBr, and CHO, we obtain
∆(G°(385 K) - H°(298 K))/T ) -29.62 J K-1 mol-1
and therefore
∆(G°(385 K) - H°(298 K)) ) -11.40 kJ mol-1
This gives
∆H°(298 K) ) -6.79 + 11.50 ) 4.61 ( 0.32 kJ mol-1
Combination of this value with the known enthalpies of
formation10,29 of Br, HBr, and CH2O (Table 3) leads to ∆Hf°-
and Ea is uncertain to (0.56 kJ mol-1
.
(CHO) ) 44.34 ( 0.59 kJ mol-1
.
It is interesting to compare the Arrhenius parameters obtained
in this work with those of reactions of other radicals with HI
and HBr.18,19 These are shown in Table 4. For reactions of
alkyl radicals with HI, A factors are typically ca. (3-5) × 10-12
cm3 molecule-1 s-1, which is very similar to the value obtained
for CHO + HI. The small positive activation energy for CHO
+ HI apparently contrasts with the negative activation energies
seen for alkyl radicals with HI. However, the values are
generally small and are consistent with (although not a proof
of) a reaction proceeding via an intermediate complex. For
reactions of alkyl radicals with HBr, A factors are typically ca.
(1.3-1.7) × 10-12 cm3 molecule-1 s-1, which is ca. 6 times
the value obtained for CHO + HBr. Although this might seem
like an unusually large difference for such apparently similar
reactions, the A factor for the reverse reaction, Br + CH2O
(Table 2 and ref 16), is lower than those for Br atom reactions
with alkanes22 by a similar order of magnitude. The A factor
for CH3CO + HBr20 is also lower than those for alkyl radicals
+ HBr but by a smaller factor. This suggests that for CHO +
HBr, and likewise the Br + CH2O reaction, the activated
complex has a tighter structure than those for the alkyl radical
+ HBr reactions. A possible explanation for this could lie in
the more polar character of the reactants in this case. The carbon
atom of CHO is more positively charged than the odd-electron-
bearing carbon atoms of alkyl radicals while HBr has a bigger
The third law method thus produces almost perfect agreement
between the values obtained from the two reaction systems. It
is intrinsically more precise than the second law method and
leads to an averaged estimate:
∆Hf°(CHO) ) 44.29 ( 0.43 kJ mol-1
From this may be derived the following bond dissociation
energies:
D(H-CHO) ) 370.86 ( 0.28 kJ mol-1
D(H-CO) ) 63.18 ( 0.46 kJ mol-1
Discussion
A. Kinetics. There are no previous studies of the kinetics
of the reactions of CHO with HI or HBr. The laser flash
photolysis technique provides direct and absolute measurements
of the reaction rate constants. A number of aspects of the study
provide assurance that there are no serious errors. The
photolysis wavelength of 308 nm is not absorbed by either HI
or HBr, and so adventitious and undesirable photolytic side
reactions are avoided. CHO radical concentrations are probed
centrally in the reactor, and decay times are of the order of 10
µs. At the pressures of study this is fast compared with diffusion

Reactions of CHO with HI and HBr
J. Phys. Chem. A, Vol. 101, No. 23, 1997 4189
TABLE 5: Summary of Experimental Values of
work rectifies these two defects as well as backing up the result
via the additional study of the reaction pair (3), (4). The
photoionization studies of Warneck5,6 set out to measure
threshold energies for formation of H+ and OH+ from CH2O
and HCOOH, respectively, where CHO is the neutral fragment.
Without any corrections for nonthermal effects these lead to
∆Hf°(CHO)/kJ mol-1 at 298 K
experiment
value
30 ( 8
ref
kinetics of iodination
of CH2O (1966)
4
photoionization threshold
of CH2O, HCOOH (1974)
photodissociation threshold
in CH2O (1978)
direct photoionisation
of CHO (1980)
photodissociation threshold
in CH2O (1983)
CHO radical recombination
(1985)
42 ( 5
39 ( 4
28 ( 5
44 ( 3
43.5 ( 8
5, 6
7
individual (298 K) values of 41 ( 7 and 43 ( 5 kJ mol-1
.
Photodecomposition studies essentially attempt to obtain the
threshold energy of the process
8
CH2O f CHO + H
9
10
which is a measure of D0(CHO). This can be either corrected
to 298 K or used to obtain ∆Hf°(CHO) at 0 K, which can itself
be corrected to 298 K. The difficulty in carrying out the
experimental studies themselves lies in identifying the true
threshold for dissociation without assistance from thermal or
internal vibrational energy in CH2O. Thus, typically values have
tended to be slightly low in the past. Reilly et al.7 cite the value
of D0(H-CHO) ) 360 ( 4 kJ mol-1, corresponding to ∆Hf°-
(CHO) ) 39 ( 4 kJ mol-1 (at 298 K), while Moortgat et al.9
obtained D0(H-CHO) ) 364 ( 3 kJ mol-1, corresponding to
∆Hf°(CHO) ) 44 ( 3 kJ mol-1 (at 298 K). More recently,
Chuang, Foltz, and Moore12 have identified the threshold by
simultaneously monitoring absorption (photoacoustic spectrum)
and H atom dissociation excitation (PHOFEX spectrum) in both
CH2O and CD2O. The threshold is identified by breakoff in
the PHOFEX spectrum. This result leads to D0(H-CHO) )
362.21 ( 0.67 kJ mol-1, corresponding to ∆Hf°(CHO) ) 41.80
( 0.81 kJ mol-1 (at 298 K), by far the most precise value yet
obtained.
reevaluation (1987)
photodissociation threshold
in CH2O (1987)
kinetics of CHO + HI, HBr;
third law value
43.5 ( 4
11
12
41.8 ( 0.8
44.29 ( 0.43
this work
dipole than HI. Therefore, the long-range interactions of these
species should lead to initial formation of a complex of structure:
H
H
Br • • • C
O
If the reorganization of this species is more difficult than that
of the equivalent H-I‚‚‚CHO complex due to stronger binding
forces, then a tighter transition state could well arise. The
activation energy is almost zero for CHO + HBr (just as for
CHO + HI) compared with the negative values for alkyl +
HBr, but this is still consistent with the involvement of an
intermediate complex. For CH3CO + HBr the activation energy
is negative (but small).20 This is not inconsistent with this idea
if methyl groups can assist in lowering the secondary barriers
of rearrangement of complexes of this type. This requires the
second step of the process to be rate determining but is consistent
with the trends of increasingly negative values for the activation
energies of the alkyl radicals + HBr (and HI) with methyl
substitution in the alkyl radical.18,19 Interestingly, it is also
consistent with methyl group substituent trends in the seemingly
quite different reactions of silylene insertion into the Si-H
bonds of methylsilanes (MenSiH4-n, n ) 1, 2, or 3). The
evidence for involvement of intermediate complexes in this class
of reactions is very strong.30,31
In the photoelectron spectroscopic study, Dyke et al.8 meas-
ured the vertical ionization energy of CHO directly and used it
to estimate ∆Hf°(CHO) ) 28.0 ( 5.4 kJ mol-1. It seems
probable that the low value results from an incorrect assignment
of the adiabatic ionization potential. The fifth method36 involves
observation of chemiluminescence from glyoxal, (CHO)2,
formed via CHO radical recombination, giving D298(OHC-
CHO) ) 299.2 ( 1.3 kJ mol-1. This was used by Fletcher
and Pilcher29 to obtain D298(H-CHO) ) 370.3 ( 3.8 kJ mol-1
,
corresponding to ∆Hf°(CHO) ) 43.7 ( 3.8 kJ mol-1. The
JANAF tables10 citing this last data give more conservative error
limits of (8 kJ mol-1
.
Some of these data were briefly reviewed earlier by Timonen
et al.,11 who suggested a consensus value of ∆Hf°(CHO) ) 43.1
( 4 kJ mol-1 (at 0 K), corresponding to 43.5 ( 4 kJ mol-1 at
298 K. It thus appears that all values since 1974 (except those
of ref 8) for ∆Hf°(CHO) at 298 K lie in the range 40-45 kJ
mol-1. Error margins of all studies prior to that of Chuang,
Folz, and Moore12 cover this range. The discrepancy now is
between the more precise values of the latter and the value
derived in this work, the error margins of which do not overlap.
We have looked carefully at the possible errors arising in this
study. We focus first on the third law derived values, since
these are the most precise. Errors in the rate constants (Table
1) are single standard deviations (68% confidence). An increase
to two standard deviations (95% confidence) would lead to a
final uncertainty of ca. (0.6 instead of (0.4 kJ mol-1, not a
large change. The thermodynamic functions of the species
involved in reactions 1-4 should be reliable to a high degree
of precision, since the structural parameters (geometries,
vibrational wavenumbers) required for their calculation are well-
established.10 We estimate that uncertainties deriving from this
source are unlikely to exceed (0.1 kJ mol-1. The second law
derived values of ∆Hf°(CHO), while inevitably less reliable,
do not undermine confidence in the results obtained here. The
The results of the present work are therefore fully in tune
with the long-held, well-known involvement of polar effects in
alkyl radical H atom transfer reactions32-34 and the more recently
proposed arguments in favor of intermediate complexes.18-26
B. Thermochemistry of CHO and the Related Bond
Dissociation Energies, D(H-CHO) and D(H-CO). Previous
experimental values for ∆Hf°(CHO) are listed in Table 5. Only
data since 1966 have been considered as values prior to that
were reviewed by Walsh and Benson.4 The data shown have
been obtained by five different types of measurement: iodine-
formaldehyde kinetic studies, photoionization studies, photo-
decomposition studies, photoelectron spectroscopic studies, and
formyl radical recombination studies. The kinetic study of
reaction 1 carried out previously by Walsh and Benson4 had
two deficiencies as far as the determination of ∆Hf°(CHO) is
concerned. First, it required an estimate of Ea,2 (since this was
not experimentally available), and second it was based on an
older, less reliable value35 for ∆Hf°(CH2O) of -116 ( 6 kJ
mol-1 compared with the more recent and precise flame
calorimetric figure of -108.57 ( 0.46 kJ mol-1 obtained by
Fletcher and Pilcher.29 For some inexplicable reason, the
JANAF tables10 continue to prefer the older value. The present

4190 J. Phys. Chem. A, Vol. 101, No. 23, 1997
Becerra et al.
values from the two reaction pairs (1), (2) and (3), (4) differ by
2.51 kJ mol-1. While the error margins do not quite overlap at
one standard deviation uncertain, they do at two. The values
span the third law derived values. Small differences between
second and third law derived values commonly arise. The
second law values are usually less precise (and therefore less
reliable). This is because a second law value relies on
measurements of the values of the gradients of two Arrhenius
plots. Small distortions (unknown systematic errors) in the rate
constants at the temperature extremes of the kinetic study
(always the most demanding) can easily give rise to errors of a
few kJ mol-1 in activation energy values, even though rate
constant values themselves may be reliable to (5% in them-
selves. The judgment of the quality of a second-order value
for ∆H° of a reaction rests on how well the ∆S° value derived
from the A factor ratio fits the known or estimated ∆S° from
thermodynamic considerations. For both reactions of this study
the agreement is quite good, if not perfect. It is certainly
comparable with the best obtained for the equivalent H
abstraction reactions of I or Br atoms with alkanes.18,19
Nevertheless the third law values are much more precise,
especially in this case where ancillary thermodynamic quantities
are so reliable. The puzzle that remains, therefore, is the origin
of the discrepancy between this work and that of Chuang et
al.12 We can throw no more light on it except to indicate the
implications. If our results are correct, it means that D0(H-
CHO) ) 364.70 kJ mol-1, corresponding to a photodissociation
threshold for CH2O of λ ) 328.0 nm compared with 330.3 nm
observed.12 If Chuang, Foltz, and Moore12 are correct, it means
that either our measured rate constants are too large by ca. a
factor of 2 (1.8 for k2 at 500 K and 2.2 for k4 at 385 K) or k1
and k3 from the published studies4,16,17 are too small by
equivalent factors. Since there are two studies of k3 that are in
close agreement (rate constant differences are e14% over the
whole range 295-480 K), it seems unlikely that this value can
be seriously in error. The consistency of the kinetics makes it
appear unlikely that k1 is also in error. There is thus a strong
case for the reliability of the data based on kinetics.
(Spain) as well as British Gas (UK) to R.B. We thank Monty
Frey for helpful discussions and Ben Mason for preliminary
experimental measurements.
References and Notes
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1185.
(10) Chase, M. W., Jr.; Davies, C. A.; Downey, J. R., Jr.; Frurip, D. J.;
MacDonald, R. A.; Syverud, A. N. JANAF Thermochemical Tables, 3rd
ed.; J. Phys. Chem. Ref. Data 1985, 14 (Suppl. 1).
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Chem. 1987, 91, 5325.
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(13) See ref 3 and references therein.
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Although it is not our prime consideration to review theoreti-
cal work in this area, it is worth noting that the considerable
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most recent of these37 gives D0(H-CO) ) 54.98 kJ mol-1
,
corresponding to ∆Hf°(CHO) ) 47.61 kJ mol-1. The agreement
with experiment is reasonable, if we accept the the expressed
uncertainties of (4 kJ mol-1 of earlier theoretical calcula-
tions.38,39
It is also interesting to note the close similarity in values of
D(H-CHO) and D(CH3CO-H) for which the most recent
values are 373.8 ( 1.5 kJ mol-1 (experiment23) and 374.9 (
2.8 kJ mol-1 (theory40). It appears that Me-for-H replacement
in formaldehyde marginally increases the remaining aldehydic
C-H bond strength.
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(40) Bauschlicher, C. W., Jr. J. Phys. Chem. 1994, 98, 2564.
Acknowledgment. Support for this work was provided by
the UK EPSCR in the form of an earmarked studentship for
I.W.C. and by grant project PB94-0218-C02-O1 of the DGICYT