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The high-temperature oxidation of aldehydes

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Page 1: The high-temperature oxidation of aldehydes

THE HIGH-TEMPERATURE OXIDATION OF ALDEHYDES

R. R. BALDWIN, D. H. LANGFORD,* M. J. MATCHAN,~ R. W. WALKER, AND D. A. YORKE:~

Department of Chemistry, The University, Hull, Yorkshire, England

Studies of the oxidation of aldehydes have given the velocity constants k~ = 0.013 and 0.076 liter mole -I sec -~ for ttCHO and C~HsCHO, respectively, and k4 = 1.36 X 106, 1.82 X 10 s, 2.41 )< 10 ~ liter mole -1 sec -1, for HCHO, CaHsCHO, and n-C3HTCHO, respectively, at 440~ on the assumption that k~ = 1.8 X 10 ~. These are preliminary values which may be modified slightly by a more-detailed treatment.

AH -b02 = A- k i l O 2

HO2 -b AH = A -b H~O,

2 HO~ = H~O~ q- 02.

(1)

(4)

(5)

Values of 0.028, 0.088, and 0.155 for C:I-I6, C3H8, and i-C4H10, respectively, have been ob- tained at 440~ for the velocity constants for reaction of HO~ radicals with the hydrocarbon relative to their reaction with HCHO.

Velocity constants for the reactions of C~H~ radicals with O2, H: and D~, for the reaction of CHs radicals with CH3CHO, and for the reaction of n-C3H7 radicals with O~, have also been obtained. The formation of CH3CHO in the oxidation of C~H~CHO, and of C~HsCHO in the oxidation of n-C~HTCHO, is shown to be due to radical attack at the secondary CH, position of the aldehyde, rather than to reaction of the alkyl radicals with 0,. The negative temperature coefficient of about 20 kcal mole -1 for the oxidation of CH5 radicals is attributed to the occurrence of the reversible reaction CHa + O~ -~ M = CH~02 4- M, and information is also provided on the possible reactions of the CH302 radical.

Although most of the features of hydrocarbon oxidation are qualitatively understood, I very few of the velocity constants of the elementary reac- tions involved are known. The lack of quantita- tive data arises first from the complexities in- herent in hydrocarbon oxidation, and second, from the difficulty in finding reliable sources of the radicals. Two particular features of hydro- carbon oxidation have aggravated the problem; first, the fact that the nature and relative con- centrations of the radicals vary as the reaction proceeds, and second, the fact that the reaction intermediates are more reactive than the parent hydrocarbon, so that the study of hydrocarbon oxidation soon becomes a study of the reactions of its intermediates.

With the development of pressure transducers for the measurement of small and rapid pressure

* Senior Lecturer at the Hull College of Tech- nology.

t Now at Glaxo Laboratories Limited, Ulverston. ~: Now at B.P. (Chemicals) Limited, Epsom.

changes, it has been possible s recently to study the oxidation of simple aldehydes between 400 ~ and 500~ Such studies have the advantage that the main products, normally CO and an olefin, are much more stable than the parent aldehyde, so that the mechanism of oxidation is relatively simple. Such oxidations thus can be used as a source of H02 radicals, and of specific alkyl radicals. The simplicity and potential of the method can be illustrated by reference to the oxidation of C2HsCH0 in aged boric-acid-coated vessels, for which the following basic mechanism can be written2;

C2HsCHO -{- 02 = C2H5C0 -~ HO2 (1)

C2H~CO ~- M = C2H5 ~- CO ~- M (2)

C~H5 + 02 = C~H4 + H0~ (3)

H02 q- C2HsCHO = H202 + C~HsCO (4)

2 HOe -- H202 ~- 02. (5)

251

Page 2: The high-temperature oxidation of aldehydes

252 LOW-TEMPERATURE OXIDATION OF HYDROCARBONS

Stationary-state t reatment gives the rate equa- tion

- - d E A H ] / d t = d (AP) /d t

= k4(kJks)ll2EAH]312EO2"]lm, (i)

where A p is the pressure change. Figure 1 (curve A) shows the variat ion of

reaction rate with temperature, and indicates the existence of a negative temperature coeffi- cient between 350 ~ and 400~ Detailed studies were thus made at 440~ in the hope tha t com- plications resulting from peracetic acid formation would be absent, and that the simple mechanism above would operate. A standard mixture con- taining 4, 30, and 26 mm Hg of C2HsCHO (AH), Om and Nm respectively, was selected. With this low pressure of C3HsCHO, reasonable rates (half- life time ca. 60 sec) were obtained, which could be followed with the aid of a pressure transducer. The effect of independent variation of C3H~CHO and 03 concentrations was examined by inter- changing each with Nm while the effect of inert- gas addition and vessel diameter was also examined.

Analysis of reaction products over a range of conditions showed that A p gave a precise measure of reaction, up to 50~o reaction. The main reaction products were C2H4, CO, and H203, consistent with the above mechanism. The other primary products were C02 and CH3CtIO in about 10% yields, and traces of C~He and C2H40. No analysis was made for tI20.

The AP- t ime curves were markedly auto- catalytic, so that measurement of the initial rate was difficult, but orders of 1.5, 0.4, and 0.1 in AH, 02, and N2, respectively, were obtained, in close agreement with Eq. (i); the slight diver- gences may well be due to failure to measure the initial rate because of the finite mixing period of 2 sec before the reaction vessel tap was closed. For the maximum rate, the order in AH at low AH concentrations was 1.5, increasing toward 2.0-2.5 at the highest AH concentration of 10 mm Hg. The order in 03 was 0.1-0.2, while the order in N2, obtained from the plot of log rate-log total pressure on addition of up to 600 mm Hg of N2, was 0.2 at low N2, increasing toward 0.5 at high N2 concentrations. Both initial and maximum rates were independent of vessel diameter over the range 15-55 mm.

The autocatalytic nature of the reaction can be at tr ibuted to the homogeneous decomposition of H302 according to Reaction (6), the OH radicals then undergoing Reaction (7) ;

H202 + M' = 2 OH -}- M' (6)

OH + C2H~CHO = H20 + C~HsCO. (7)

I t is now necessary to write differential equations for both AH loss and for H~02 formation, and these equations, given below, can be solved only by numerical integration using computer methods,

- - d [ A H ] / d t = klEAH][O2] + (k4/ks~m)GEAH ]

+2k6EH20:][-M'-], (ii)

1.2 -I

0.8 [ ~ k o .c_

E

0 I I I 1 1.2 1.3 1'4 1.51

YT x 10 31.6

FIG. 1. Variation of reaction rate with temperature for acetaldehyde and propionaldehyde oxidation" Curve A: C~H~CHO = 4 m m H g ; 0 2 = 30 ramHg;N~ = 2 6 m m I t g . CurveB:CH3CHO = 2 mm Hg; 02 = 30 mm Hg; N2 = 28 mm Hg.

Page 3: The high-temperature oxidation of aldehydes

HIGH-TEMPERATURE OXIDATION OF ALDEHYDES 253

d~I202-]/dt = kIEAHT[023 -]- (h/ks~12)G[AH],

(iii)

where

G= ks1/~[H02], (iv)

and is given by

G 2= k,[AH][O2] -5 k6[H2023[M']. (v)

Since k8 is accurately known, 3 the only un- known parameters are kl and (ka/k51/2), and step- wise integrations have been carried out for a range of values of kl and (k4/k51/2) of interest; suitable step lengths were chosen to give adequate accuracy. The approximate magnitude of kl can be found from the shape of the AP- t ime curve, as shown in Fig. 2. A value of 10 -4 (mm Hg sec) -1 gives maximum rate initially, and the rate falls off sharply as the reaction proceeds (curve A). With k, = 10 -s, the S-shaped nature of the curve is excessive (curve C), and a value of about 10 -6 gives a close fit between observed and calculated AP- t ime curves. For values of k, in this range, the primary initiation process (1) is negligible compared to the dissociation of H202 at maximum rate, so that the parameter (]Q/k5112) c a n be ob- tained from the maximum rate for the standard mixture. The value of kl can then be obtained in

two ways:

(a) as the value [1.4 X 10 -s (ram Hg sec) -1] that gives the experimentally determined ratio of initial ra te /maximum rate:

(b) as the value (2.0 X 10 -6) which, as shown in Fig. 2 (curve B), gives the best fit to the observed AP- t ime curve.

The discrepancy between these two values prob- ably arises from an accumulation of small effects, such as the difficulties in the measurement of initial rate, and the differences between A[AH] and AP resulting from further reactions of the products (particularly oxidation of CH3CHO), and a mean value of 1.7=[=0.3X10 --6 ( m m H g sec) -1, or 0.076 liter mole -1 sec -1 at 440~ will be adopted. Confirmation of this value is obtained by plotting the estimates of initial rate against [AHJ312E02]II2 for a wide range of mixture com- position. Equation (i) gives the gradient as ka(kl /kS) 1/2, and from the value obtained for (k4/]% 1/2) from maximum rate measurements, kl = 2.25 X 10 -6 ( m m H g s e c ) -1. This value is probably a bit too high because of the experi- mental difficulty that the rate cannot be measured until 2 sec after the gases have been mixed.

These values, and the others quoted in Table I, have been obtained by making the reasonable assumption that the efficiency of the aldehyde in Reaction (6) is four times tha t of N~. With this value, the computer program predicts a rather

I I I [ l {

A B

2

0- :Z

E E

1

i I 1 i 0 2 0 4 0 6 0 8 0 100

Time in sec

Fx~. 2. Ap vs time relationship in the oxidation of propionaldehyde at 440~ C2HsCHO = 4 mm Hg; 02 = 30 mm Hg; N2 = 26 mm Hg. - - - , computed curves; X, experimental points. Curve A: k, = 10 -4 (mm Hg sec)-l; (k4/k5 w2) = 0.2 (mm Hg see) -1/2. Curve B: k~ = 2.0 X 10 -~ (mm Hg sec)-l; (k4/k511~) = 0.206 (ram Hg sec) -1/~. Curve C: kl = 10 -8 (ram Hg see)-1; (k4/k5 u2) = 0.2 (mm Hg sec) -1/2.

Page 4: The high-temperature oxidation of aldehydes

254 LOW-TEMPERATURE OXIDATION OF HYDROCARBONS

smaller effect of inert gas than is found experi- mentally for C~H~CHO. A likely explanation of this discrepancy is that C2H~CHO is particularly effective in Reaction (6), possibly as a result of complex formation. ~-7 Computer t reatment sug- gests that the most satisfactory interpretation of the effect of addition of N2 is obtained if C2H~CHO is some 30 times more efficient than N2 in Reaction (6). With this value, the dissocia- tion of H202 is increased significantly and the value of (k4/k5112) will be reduced by a factor of about 2, kl being proportionally increased. In the case of C2HhCHO, allowance has been made for the occurrence of Reactions (9 ) and (10), given later in this paper. Since these represent propaga- tion reactions involving radicals that do not undergo termination reactions, they accelerate the reaction, and thus reduce the value of (k~/k51/2) by about 15%. Reaction (13), given

later, has a similar effect but, because the yields of C02 are smaller, it has not been allowed for in this paper.

With C2HhCHO, studies have been made only at 440~ so tha t the activation energies have not been directly determined. From the value of kl at 440~ however, use of pre-exponential factors of 109, 10 l~ and 10 n gives values of E1 equal to 33, 36, and 40 kcal mole -1, respectively, which may be compared with the estimated s endothermicity of 40-41 kcal mole -1. Even with the smallest value of E1---- 33 kcal mole -1, kl would decrease to ca. 6 X 10 -1~ liter mole -1 sec -~ at 123~ so that the values of 3.75 ~ l0 -a and 1.0 ~ 10 -~ liter mole -~ sec -1 at 123~ reported by Farmer and McDowell, 9 and by Combe, Niclause, and Letort, ~~ respectively, for the initiation process in CH~CHO oxidation, are clearly too high for a homogeneous initiation by (1) and must represent either a heterogeneous initiation, or some secondary initiation process.

With n-CaHTCHO, the basic mechanism is similar to that for C~H~CHO. Because of the fast rate with this aldehyde, significant reaction has occurred by the time the "init ial" rate is meas- ured, and the possibility of initiation by secondary products cannot be excluded, so tha t the value obtained for k~ is not considered reliable.

In the case of HCHO, the mechanism is even simpler, comprising Reactions (1F), (3F), (4F), (5), (6), and (7F).

Since the reaction involves no pressure change in the early stages, it has been followed by analysis for the CO produced. In the later stages of the reaction, pressure changes occur due to decom- position of H202, and to oxidation of the CO.

The figures in Table I, obtained using the simple procedure outlined in the previous para- graphs, must be regarded as prelm~inary values. They have been obtained using a value for k6 ob- tained a from studies at atmospheric pressure on the assumption that the decomposition of H~02 is in its second-order range, so that k6 is directly proportional to pressure. There are indications that the plot of the apparent value of/c~ against pressure does not pass through the origin, and until this behavior is understood, the values of k4/k61/2 represent an upper limit, which may be high by 50-100%, while the values of kl may increase by a similar proportion. However, this uncertainty about k~ does not affect the relative values of k~/k51/~, which should be accurate to 20%.

With HCHO, studies over the range 400 ~ 500~ have enabled the activation energy E4 to be estimated. The simple procedure gives (E4F -- -~Eh) = 10 =t= 2 kcal mole -1, A4F/A~ 1/2 = 2.4 X 104 (liter mole -1 sec-1) 1/2. Accepting the only currently available value l~ for k5 of 1.8 X 109 liter mole -~ sec -1, and assuming that it is independent of temperature, A4F = ].0 X 109, E4F = 10 -4- 2 kcal mole -1. The only other reaction of HO~ radicals for which a reasonable estimate of activation energy is available ~ is Reaction (4H), for which E4H = 25 -4-- 5 kcal mole -1.

HO2 + H2 = H202 + H. (4H)

S i n c e s AH4F ~--- - -2 kcal mole -1, AH4~ = 14 kcal mole -1, the value of a in the Evans-Polanyi 12 relation E = a AH + c, is close to unity. Reason- able estimates of the activation energy for reac- tions of H02 with hydrocarbons and related com- pounds thus can be made from a knowledge of their thermochemistry.

TABLE I

Values at 440~

HCHO + 02 = HCO + H02 (1F)

HCO + 02 = H02 + CO (3F)

HO2 + HCHO = H202 + HCO (4F)

OH + HCHO = H20 + HCO. (7F)

kl, (k4/kY~), Aldehyde liter mole -1 sec -1 (liter mole -1 sec-l) ~/~

HCHO 0.013 32 C~HbCHO 0.076 43 n-CsHTCHO - - 57

Page 5: The high-temperature oxidation of aldehydes

HIGH-TEMPERATURE OXIDATION OF ALDEHYDES 255

By adding a hydrocarbon to reacting HCHO -+- O~ mixtures, and by measuring the relative amounts of products formed from the hydro- carbon and from HCHO, the relative velocity constants for attack of HO2 on the hydrocarbon and on HCHO have been obtained. At 440~ for C2H~, C3H8, and i-C4H~0, the ratios are 0.028, 0.088, and 0.155, respectively.

C:H~ ~ 02 -- CHaCHO -[- OH (8)

In the oxidation of C2H~CHO at 440~ signi- ficant yields (ca. 10%) of CH3CI-IO are obtained. I t is tempting to suggest that this results from Reaction (8), which is of a type that has fre- quently been suggested in hydrocarbon oxidation. However, studies 13 of the addition of C2H6 to slowly reacting mixtures of H~ ~ O2 at 500~ show that at least 96% of the C2H~ oxidized appears initially as C2I'Ia, and that no more than 0.4% appears as CHaCHO. To explain these observations in terms of a competition between Reactions (3) and (8) would require Reaction (3) to have an activation energy of at least 60 kcal mole -~, which is clearly impossible. I t is thus concluded that CH3CHO is formed as a result of attack at the CH2 position, the CHaCHCHO radical then reacting with O~ to give CHaCHO:

OH ~ C2H~CHO = CHaCHCHO -+- H20 (9)

CH3CHCHO + 02 = CH3CHO ~ CO ~ OH.

(lo)

The present results do not permit a decision as to whether the CH3CHCHO radical is formed pre- dominantly by HO2 attack, or by attack through the less-selective OH radical.

A similar mechanism will account for the significant yields of C2I.I~CHO (ca. 8%) formed in the oxidation of n-C3HTCHO under similar conditions. This again contrasts with the very low yields of C2HsCHO found ~4 when small amounts of C3Hs are added to slowly reacting mixtures of H2-r 02 at 480~ indicating that C2HsCHO is only a very minor product from the reaction of C3H~ radicals with 02:

0 //

CHaCH2CHCHO + O~--~ CH3CH2CHC I \ O H

O

CHaCH2CHO + CO -~ OH.

L

i i

i AL-

x/x ~

/ /

/ /

? r

i

j /

/

0 0.1 O . 2 I

FiG. 3. Plot of [-C2H4J/rC~H(] against ['O2]/rAH] for oxidation of propionaldehyde at 440~

A similar mechanism accounts for the formation of aldehydes in the oxidation of straight-chain alcohols.~-~8

Although studies of aldehyde oxidation are thus important in providing information about reactions of HO:, they also enable reactions of alkyl radicals to be examined. With C2HsCHO, small amounts of C2H6 are found in the reaction products, presumably from the reaction

C2H5~ C2H~CHO = C2H6~ C2H5CO. (11)

Competition between Reactions (3) and (11) gives the relation

dEC2H4]/dEC2He] = k3EO2]/ku[AHJ.

As Fig. 3 shows, the plot of d[C2H4]/d[-C~H6"] (obtained from the initial yields of C2H4 and C2I.I6) against EO2]/[AH] gives a good straight line, the gradient giving ka/kll = 41 -r 5. Using the Brinton and Volman expression TM for kll, k3---8.2 X 107 liter mole -1 sec -1 at 440~ in reasonable agreement with the estimate made by Sampson 2~ of 1.0 "~ l0 s at 623~ Moreover, since the ratio d[-C2H4-]/dgC2H6] is not affected by inert-gas addition, the reaction forming C2H4 cannot have any third-order component, in con- trast to the reaction of CH3 radicals with Oz. [The possibility of an equilibrium C2H5 -r 02 + M ~ C2H502 ~ M, followed by C2H502---* C2H4 ~ HO2, cannot be excluded.]

By measuring the increase in the yield of C2H6 or C2H~D in the presence of H2 or D2, it is possible to obtain the velocity constant ratios

Page 6: The high-temperature oxidation of aldehydes

256 LOW-TEMPERATURE OXIDATION OF HYDROCARBONS

kn/k~ = 8.1 4- 1, kn/kl~D = 21 4- 2 at 440~

C2Hs + H2 -- C2H~ + H (12)

C~Hs + D2 = C2HaD + D. (12D)

From the value of kn, k12 = 2.15 X 10 ~ and klw = 8 .3X 104 liter mole -I see - I at 440~ Combination of these values with estimates at lower temperatures 21-2s has enabled the accurate Arrhenius parameters Alz = 6.5 X 109, A12/)= 4.1 X 109, E~ = 14.1 kcal mole -~, and E~o = 14.7 kcal mole -~, to be obtained.

With C2H~CHO, significant yields of C02 (about 10%) are formed. This is undoubtedly a primary product, although the autocatalytic nature of the [CO2] vs percent reaction curve indicates that CO2 is also formed in the later stages of the reaction by the oxidation of CO. Confirmation that this "latter reaction is unim- portant in the early stages comes from the ob- servation that the initial addition of 30 mm Hg of CO only doubles the initial rate of formation of COs. The C02 almost certainly results from the reaction

C2HaCO + 02 = C2H~C08---* COs, (13)

which competes with the reaction

CsH~CO + M = C~I~ + CO + M, (2)

This pyrolysis can be followed by measuring the relative yields of C2I-~ and Carte, since d[C2H42/ all-Call(]--]c14p[M]/k3~[02-]. Because of the subsequent oxidation of both olefins, the actual yields of C~H4 and CsH6 are measured at various stages in the reaction, and the ratio [C2H4]/ [CaH(] is extrapolated to zero reaction so as to determine its initial value. A plot of d[C~H,]/ dEC3H6J against M at constant F02-] will be independent of M if (14p) is in its first-order range, and will be linear through the origin if it is in its second-order range. Figure 4 shows that a curved relationship is obtained, indicating that the reaction is in its transition range, as expected for this type of radical pyrolysis. No other deter- minations of kt4~ have been made over such a wide pressure range. On the simple Hinshelwood- Lindemann treatment, the apparent first-order velocity constant is given by

= k ~ k b E M ] / ( k ~ E M ] + kb)

A T M ~ A * + M (a)

A*--- products. (b)

Thus, a plot of [M]/(d[-C~H,J/d[C~H6]) against M should be a straight line, for variation of M at constant [02]. The results give a surprisingly good straight line over the pressure range, and the low-pressure and high-pressure values of the ratio k~,~l,/k3 ~ can be obtained.

in which case, d[ CO2"]/ d[" CO-] = k13['O2"]/ k2[-M J, where kls may be a composite velocity constant if the reaction is reversible. Over a fifteenfold 0.24 range of 02 concentration, the plot of FC02]/ ~CO] against ['O2] is closely linear. The variation of [-C02]/[C0"] with variation of M, obtained by addition of N2, indicates that Reaction (2) is in its second-order range, in agreement with ~ 0-20 recent results obtained by Kerr and Lloyd 24 at 7- lower temperatures. Use of other inert gases is planned to examine the relative efficiency of different molecules in Reaction (2). ~: 0.16

Od By assuming that the efficiency of 02 relative to N2 in Reaction (2) is the collision frequency ratio of 0.78, and by assigning the optimum value of I0 to the efficiency of CsI-IaCHO relative to N2, the value obtained for kt3/k2 = 0.21 for 0.12 M = N 2 .

In the case of n-C3HTCHO, a further reaction of alkyl radicals becomes important, namely the pyrolysis reaction (14p), which competes with the oxidation reaction (3p):

CsH7 + M = CHa + 62H4 -{- M (14p)

C~H7 + 02 -- Call6 + H02. (3p)

0"0~

x

/ I .[ I

0 2o0 400 600 N 2 (mm Hg)

FIG. 4. Plot of [C~H4]/[CsH6] for N2 addition in oxidation of n-butyraldehyde at 440~ n- CsHTCHO = 4 mm Hg; 02 = 30 mm Hg.

Page 7: The high-temperature oxidation of aldehydes

HIGH-TEMPERATURE OXIDATION OF ALDEHYDES 257

In calculating M, the efficiency of CaH~CHO relative to N2 has been taken as the value (12) giving the best correlation coefficient for the straight line, while the efficiency of O2 has been equated to the relative collision efficiency of 0.78. From the expression

k14~ = 3.47 X 10 ~a exp (--31,400/RT) sec -~,

given by Laidler and Lin, 25 kap at 450~ = 5.8 X 107 liter mole -~ see -~. Assuming a small activation energy of 4 kcal mole -1 for Reaction (3p), this gives a value of 6.4 X 107 at 480~ which may be compared with the value of 3.8 X 10 ~ obtained from studies 28 of the addition of C3H8 to slowly reacting mixtures of H2 ~ 02.

The oxidation behavior 2~ of CHaCHO in the temperature range 360~176 shows a consider- able contrast to that of the other aldehydes. Over this wide range, the reaction rate only varies by a factor of 4 because of the existence of a negative temperature coefficient in the range 440~176 (Fig. 1, curve B). In contrast to the oxidation of HCHO and C~HsCHO, the oxidation is not autocatalytic above 400~ presumably because H202 is not an important product, being only about 10% of the aldehyde oxidized. This lower yield is not unexpected, since both HCO and C2H5 have been shown to react predominantly with O2 to give H02 radicals, which then give H202. The oxidation process for CHa radicals is less clear, but certainly there is no obvious reac- tion yielding HO~ radica]s directly.

At 540~ a mixture containing 2, 30, and 28 mm Hg of CHaCHO, 02, and N2, respectively, gives high yields of CH4 (ca. 50%) and C2H8 (ca. 7%) compared to those of the oxygenated products CHaOH and HCHO, which are both present in ca. 15%-20% concentrations. At 440~ the yields of CH, (ca. 10%) and C2H6 (ca. 0 .5~) are significantly reduced. Studies of

the relative yields of CH4 and C2H6 at 440 ~ and 540~ have given values for the ratio (kl~,/k15112) of 57.5 and 110 liter ~[2 mole -~/2 sec -~1~, respec- tively. If the transition-state theory prediction that k~5 is proportional to T -2 is used with a value of k15 = 2 . 9 X 10 ~~ at 165~ suggested by Shepp's recalculations 28 of the results of Kistia- kowsky and Roberts, ~ then kn~ = 6.0 X 106 liter mole -1 sec -1 at 440~ and 9.7 X 106 at 540~ Combination of these values with results at lower temperatures obtained by Dodd, ~~ and by Volman and Brinton, ~ gives the Arrhenius parameters All~ -- 1.6 ~ 0.6 X 109 liter mole -~ sec -~, and En~ = 8.2 4- 0.5 kcal mole-I:

If the oxidation process for Ctta radicals is represented by the over-all equation

CHa ~ 02---~ oxidation products (ox)

then from the ratio [-CH4J/([-I-ICH0J~ [-CH3OH]) -- 1.25 at 540~ and 0.10 at 4400C, Eua -- Eo~ -- 29 kcal mole -1.

With the above value of Ella, Eox -- --21 kcal mole -1. This large negative value immediately excludes the direct reaction C H a ~ 0 3 = HCHO ~ OH as the main oxidation process, though a small contribution from this reaction cannot be excluded. At lower temperatures, the formation of the peroxy radical CH302 by the termolecular reaction (16) has been generally accepted:

CHa ~ 02 ~ M = CHaO2 ~ M. (16)

I t is only possible to explain the large negative temperature coefficient for kox by assuming that Reaction (16) is reversible, and that this reverse reaction competes with a generalized reaction (x):

CHaO2 ~- X = oxidation products. (x)

The over-all velocity constant ko~ is then given by

kox = k~6[M]k~[X]/(LI,[M] + k~[X]). (vi)

A negative temperature coefficient can only be obtained if Reaction (-- 16)' predominates in the denominator, so that, on the assumption that X is a molecular species whose concentration is temperature independent, Eox approaches the value E l s ~ E x - E-16. Values of 25-30 kcal mole -~ have been estimated a~ for E-16, while El6 is likely to have a small negative value, aa perhaps up to 5 kcal mole -~. A value for Ex of about 10 kcal mole -1 is thus required to account for the ob- served activation energy associated with kox. If, however, X is a radical, whose concentration can vary with temperature, it is the product kx[-X] that must have a temperature coefficient equiv- alent to 10 kcal mole -1.

If X = CHaCH0, Reaction (17) has the right activation energy, but this reaction, followed by (18), creates excessive chain branching, which is not consistent with the kinetic features of the reaction:

CH3 ~ CHaCH0 = CH4 ~ CHACO ( l la) CH302 ~ CHaCH0 = CHaO2H ~ CHACO (17)

CHa ~ CHa = C2H6. (15) CHa00H = CHaO ~ OH. (18)

Page 8: The high-temperature oxidation of aldehydes

258 LOW-TEMPERATURE OXIDATION OF HYDROCARBONS

If X -- M, the activation energy required for Els of 10 kcal mole -1 would be rather less than estimates made by Heicklen, 32 but (18) cannot be the sole reaction because it fails to account for the significant yields of CH30H that are found. The simplest way of producing CH30H without introducing chain branching is either through the occurrence of Reaction (20), or by the occur- rence of Reaction (21), since the combination of Reactions (18) and (21) does not increase the number of chain centers:

C H 3 0 2 ~ M = H C H O ~ O H ~ - M (19)

CH302 W CH3 = 2 CH30 (20)

CH302 -~- HO2 --- CH~O2H ~- 02. (21)

A preliminary kinetic analysis suggests that at least two of Reactions (17), (19), (20), and (21) are required, and indicates the complexity of the processes involved in the oxidation of CH3 radicals.

ACKNOWLEDGMENTS

Thanks are due to the Gas Research Council for a grant to M. J. M., to the Kingston-upon-Hull Education Committee for provision of facilities for D.H.L., and to the United States Air Force for support under grants AF EOAR 67-01 and EOOAR 68-0013 through their European Office of Aerospace Research.

REFERENCES

1. MINKOFF, G. J. AND TIPPER, C. F. H. : Chemis- try of Combustion Reactions, Butterworths, 1962.

2. BALDWIN, R. R., WALKER, R. W., AND LANG- FORD, D. H. : Trans. Faraday Soc. 65, 792, 806 (1969).

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COMMENTS Z. A. Szabo. The University, Budapest. You

have assumed the formation of C2H4 in your scheme. I wonder what is the subsequent fate of ethylene. Does it interfere with the main kinetics?

Authors' Reply. One advantage of studies of the oxidation of aldehydes is that, in contrast to the position in hydrocarbon oxidation, the major products are less reactive than the parent ma-

Page 9: The high-temperature oxidation of aldehydes

HIGH-TEMPERATURE OXIDATION OF ALDEHYDES 259

terial. We have confirmed this by adding ethylene in amounts greater than those formed during the reaction, and shown that there is little effect on the over-all rate. This addition does increase somewhat the yield of ethylene oxide (C~H40), thus accounting for the autocatalytic nature of the [-C2H403-time curve. Once all the propion- aldehyde has been consumed, the C2H4 will be oxidized, predominantly to HCHO and CO, by attack of OH radicals formed from the residual H20~, but we have not considered this system a convenient way of studying these reactions and have preferred to add C2H4 to slowly reacting mixtures of H2 + 0~.

J. H. Knox, University of Edinburgh. From the k~,/k4 values obtained in this work, one may calculate ratios for the rates of H02 attack at primary, secondary, and tertiary C-H bonds as 1:6:22 for 440~ These ratios are in remarkable agreement with those obtained by Irving (un- published data) from competitive oxidations of ethane, propane, and isobutane at 300~176 namely 1:5:20, suggesting that there is large participation of H02 (or R02) radicals in these low-temperature oxidations, the ratio for OH being about 1:4:11 at 300~ (see R. R. Baker, R. R. Baldwin and R. W. Walker, This Sym- posium).