Indian Journal of ChemistryVol. 17A, March 1979, pp. 264-266

Kinetics of Oxidation of Benzaldehyde & SubstitutedBenzaldehydes by N -Bromosuccinimide

(Mrs) SUSHAMA KANDLIKAR, B. SETHURAM & T. NAVANEETH RAO*

Department ofrChemist.ry, Osmania" University, Hyderabad 500007

Received 29 April 1978; revised 25 September 1978; accepted 27 October 1978

Title investigation has been followed titrimetrically in the presence of Hg(II) acetate,sulphuric acid and 50% acetic acid (v/v). The overall order of the reaction is found to be two,unity each with respect to oxidant and substrate. The rate of reaction is inhibited by theaddition of H+ions. The rate decreases with increasing percentage of acetic acid, suggesting adipole-dipole type of reaction. A mechanism in which NBS abstracts a hydride ion from the-CHO group of the aldehyde in the rate-determining step, has been proposed. This is supportedby the fact that electron withdrawing groups decrease the rate and etectron-reteastng groupsincrease the rate. Hammett's plot of log k vs IJ is linear with a slope p = -1·14 at 50°. The pvalues increase with temperature suggesting that it is an entropy controlled reaction.

SO far no kinetic work appears to have beenreported on the oxidation of aldehydes byN-brotnosuccinimide, though product. analysis

has been made-. In the present work the kineticsof oxidation of benzaldehyde and substituted benz-aldehydes by NBS in the presence of Hg(II) acetate,sulphuric acid and 50% acetic acid has been followedtitrimetrically, with a view to elucidating themechanism of the reaction.

Materials and MethodsAll the chemicals used were of highest purity and

.wherever necessary they were further purified by. standard methods. Sulphuric acid was used formaintaining the pH and mercuric acetate was usedas a source of Hg(II) ions. The reaction wascarried out in vessels wrapped in black clothto avoid any photochemical reaction. Solutions ofaldehyde and NBS containing 0'02M Hg(OAc)2 werethermostated 'for 30 min separately before mixingto initiate the reaction. The rate of reaction wasfollowed by estimating the concentration of NBSat regular intervals by iodometry. The oxidationproduct was identified as benzoic acid.

Results and DiscussionUnder the conditions of [NBS]~[aldehyde] the

disappearance of NBS followed first order kinetics as~~~d~~~u~~~~~w~~~effect of varying the initial [NBS]. at a constant [alde-hyde] also showed that the order with respect to NBSto be one (Table 1). The data in Table 1 also revealsa first order dependence of the rate on [substrate].Hence, the rate law of the reaction could be written as

_ d[~~SJ = k"[NBS][aldehyde] ... (1)

It was observed that change of concentration ofHg(II) acetate over a four-fold range (0·025 to Q·1M)had negligible effect on the rate. The function of

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added mercuric acetate is only to fix up Br" formedin the course of reaction as HgBr2 or HgBri-.The rate was found to be inhibited by the additionof H+ ions (Table 2). This may be explained byassuming an equilibrium between unprotonated andprotonated NBS and that the unprotonated speciesis the reactive one.

TABLE 1- EFFECT OF VARYING [NBS] AND[BENZALDEHYDE] ON THE OXIDATION OF BENZALDEHYDE

{[H2S04] = 5·00 x 10-3M; [Hg(OAc)2] = 2·00 x 10-2M; solvent= 50% HOAc (v/v); temp. = 50°}

[NBS]M

[Benzaldehyde]M

0'0500·0500·0500'0500'0500·0500·0500'0500·0750'1000·1250'1500·2000'2500·300

k' X 104

sec-1k" X 10'

litre mol-! sec-!

0·0020'0030·0040'0050·0060'0080'0100'0050'0050'0050'0050'0050'0050'0050·005

5'655·605'505'705'605'555'605'708'47

11'214'217·222-628'033·3

1·131'12HO1'141'121'111'12H41'131'12H4H51·131-121-11

TABLE 2 - EFFECT OF VARYING [H+]* ON THE REACTION

{[NBS] = 5'00x 10-3M; [benzaldehyde] = 5'00x 10-2M; temp.= 50°; solvent = 50% HOAc (v/v); [Hg(OAc)2] = 2'00

X 10-2M}

103 [H+]*, l'v1.103 v, min-1

0·062'1

2'5048'3

7'5031'4

12'511'5

*Amount of added H+ at constant [HSO;;l and [HOAc].

-,\

KANDLlKAR et al.: OXIDATION OF BENZALDF:HYDE BY N13S

The influence of dielectric constant (D) of themedium on the rate of reaction was studied inbinary solvent mixtures of acetic acid and water.With increasing acetic acid content (decreasing D)of the medium the rate of the reaction decreasedto' a larger extent (Table 3). The plot of log kvs D-1/2D+1 was linear with a positive slope,indicating dipole-dipole type of reaction>, The ab-sorption maximum for NBS was found to be, thesame in the presence and absence of benzaldehydeindicating no complex formation between these twounder the experimental conditions.

In view of the overall second order kinetics andon the basis of solvent and [H+] effect studieson the reaction, a hydride ion transfer mechanism(Scheme 1) has been' suggested.

The mechanism shown in Scheme 1 gets supportfrom the substituent effect studies also. It wasobserved that (Table 4) the reaction is facilitatedby electron donating groups in the benzaldehydering. The rates of oxidation of meta- and para-substituted benzaldehydes are in the order ofp-CH3>m-CH3> H >m-Cl >m-Br>m-N02>p-N02•

(Table 4). This shows that the presence of electronwithdrawing groups in the aldehyde ring makes itless reactive by creating a more positive chargeon the carbonyl carbon atom of aldehyde groupwhich in turn makes it difficult for the hydrogento be removed as hydride ion. A plot of log kvs Hammett's cr values" for the meta- and para-substituted benzaldehydes was found to be linearwith the slope (reaction constant p) of -1·14 at50°. The negative p value indicates that mostprobably a positively charged activated complex isdeveloped during the course of the reaction. Theincrease in p with increase in temperature (-1·14,-0·71 and -0·42 at 50°, 62·5° and 68·0° respec-tively) suggests that the reaction is an entropycontrolled one.

In the case of ortho-substituted benzaldehydes,it was observed that the nitro group retards therate of reaction to a greater extent than methylgroup. However, the decrease in rate with chloroor bromo group in the ortho position is not as muchas when it is in meta position. This could bedue to the +M effect in addition to -1 effect ofthese groups in the ortho position. However, the over-all effect of these substituents (N02, Cl and Br) in theortho position is to decrease the rate in comparison toortho-methyl which is usually taken as a standard.

To study the applicability of Taft's equation- forortho-substituents, the values of log k" /k~ were

TABLE 3 - EFFECT OF SOLVENTCOMPOSITIONON THEREACTION RATE

{[NBS] = 1·00 x 10-3M ; [benzaldehyde] = 1·50 x 10-2M;temp.= 50°; [H2S04] = S'OOx 10-3M; [Hg(OAc)2] = 2·00

x 10-2M}

103 X k' (min-I) in solvent containing HOAc (%. v/v)

20·020·1

90·01·30

40·014·3

60·06'20

80'01'70

TABLE 4 - ACTIVATIONPARAMETERSFOR THEOXIDATION OF BENZALDEHYDEBY NBS

{[NBS] = 5·00 x 10-3M; [benzaldehyde] = 5'00 x 10-2M;temp.e- 50°; [H2S04] = s'00x10-3M; [Hg(OAcl.l = 2·00

X 10-2M; solvent = 50% I-WAc}

Substi- k"x 103 ~Ej: ~Ht l1Gt l1Sttuent litre mol+! sec.-I e.u.

H 11'5 23-4 22-8 21'8 3'20m-CH3 13-2 18·3 17'7 21'2 -9'30m-Cl 4·50 13'7 13'1 22'4 -28'8m-Br 3·30 14·1 13·5 22'5 -27'8m-NO. 1'90 23·8 23'2 23'0 0'620o-CR. 5·40 22'8 22'2 22·3 -0·310o-Cl 3'70 13'7 13·1 22·5 -29'1o-Br 3'20 16'0 15'4 22'5 -22·0o-NO! 1'40 5'50 4·90 23·1 -56,3p-CHa 17·5 20·8 20'2 23·4 -10·0P-NO. 1'50 22·3 21·7 24·0 -8,95

l1Et. ~Ht and l1Gt values in kca.l mo.r",

plotted against cr* values at 50° according toequation

log k"/k~ = P*cr*+oEs ..• (2)where k" and k~ represent the rate constants for theoxidation of any ortho-substituted benzaldehyde andortho-tolualdehyde respectively, cr* and E; are thepolar and steric parameters, P* and 3 are thereaction constants for the same two effects res-pectively. The plot of log k"/k~ vs cr* did notgive a good correlation, though the plot was fairlylinear. Hence, using the method of solving thesimultaneous equation", it was found that the kineticresults' for ortho-substituents could be fitted wellin the equation (2) with p* = -0·23 and 3 = 0·16.The plot of (log ktrlk~-3Es) vs cr* gave a goodlinear plot with a correlation coefficient r = 0·989,showing the presence of both polar and steric effectsin the reaction. Also it was found that the

~.

",,'" 10",. + B, -.~ ~,==='[cO"~~:~:t~~J1Activated c o rnpl e x Sla,.

"B, + [j'" + ,,",CO,"o

(

Scheme 1

265

INDIAN J. CHEM., VOL. 17A, MARCH 1979I

calculated values of log k" /k~ according to equation(2) correlated better with the experimental values.This correlation indeed shows the presence of bothsteric and polar effects in the reaction.

The negative ~* value shows that the reaction isfacilitated by high electron density at the reactionsite. Also the p* value is different and lower thanP value (-1·14 at 50°) obtained in the case ofmeta-substituted benzaldehydes. This again confirmsthat in audition to polar effects, the rates of reactionof aitha-substituted benzaldehydes are affected bysteric factors also.

The L~.Et values calculated from the Arrheniusplot for all aldehydes and other thermodynamicparameters are presented in Table 4. From theseresults it may be considered that both the para-meters (LlHt and LlSt) are important in controllingthe rate of reaction, though the entropy factorseems to be the predominating. This point was also

266

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verified by calculating the isokinetic temperatures(~) from the slope of the plot of LlHt vs LlSt.The ~ value was found to be 300 K which isbelow the temperature range used in the presentwork (320 to 338 K). Constancy of LlGt valuescalculated for all aldehydes indicates that probablythe same mechanism prevails in all the cases.The negative values for entropy of activation reflectsthat the activated complex is more rigid.

References1. FILLER, R., Chem, Rev., 63 (1963), 21.2. HAMMETT,L. P., Physical organic chemistry (MCGraw-Hili,

New York), 194P, 193.3. AMIS, E. S., Solvent effects on reaction rates and mechanism

(Academic Press, London), 1966.4. TAFT, Jr., R. W., NEWMAN, M. S. & VERHOCK, F. H.,

J. Am. chem, Soc., 72 (1950), 4511.5. ANDERSON,R. L. & BANCROFT,T. A., Statistical theory in

research (McGraw-Hill, New York), 1952, Ch. 14. .6. LEFFLER, J. E., J. org, Chem., (1955), 1202.