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Volume 43, 1910
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Art. VII.—On the Rate of Oxidation of Acetaldehyde to Acetic Acid.

[Read before the Philosophical Institute of Canterbury, 2nd November, 1910.]

Introduction.

In the 1894 number of the “Philosophical Magazine” an account is given of some experiments made by Dr. T. Ewan on the rate of oxidation of acetaldehyde to acetic acid, and the conclusion he comes to from his experiments is that aldehyde is oxidized to acetic acid in the vaporous state at a rate proportional to the concentration or pressure of the aldehyde and to the square root of the oxygen-pressure This, however, did not seem to apply when the pressure of oxygen was above 450mm. of mercury with the temperature at 20° C. In fact, he was unable to obtain any evidence of action with an oxygen-pressure of 599mm. even when the temperature was 21·4° C.

The peculiarity of this sudden cessation of action, and the small number of the experiments of Dr. Ewan which are described, seemed to me to call for more experiments, as he himself says, to clear up this interesting behaviour. No record of the work having been continued either by himself or others could be found in the “Journal of the Chemical Society,” the “Philosophical Magazine,” the “Transactions of the Royal Society,” the “Journal of Physical Chemistry,” or the “American Journal of Science.”

Description of Apparatus.

The apparatus used was in principle the same as that of Dr. Ewan, though with a few modifications, and is shown in the figure on the next page.

In Dr. Ewan's form of apparatus the mixture of aldehyde and oxygen was left in contact with the mercury, thus necessitating the use of bromnaphthalene to protect it, and also resulting in the volume of the reaction-vessel changing as the pressure decreased, the mercury altering its position

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unless constantly rectified. Both of these difficulties were overcome by using a tap K, by means of which the mercury was shut off from contact with the aldehyde and oxygen, except for a few seconds while a reading was being taken. The tap K was a three-way tap the other limb connecting with the air.

The volume of the reaction-vessel thus remained constant, and the reaction could be allowed to proceed an indefinite time without further attention. The readings were always taken with the mercury just up to the tap. The tap used was of the Geissler form, this form being the best able to prevent leakage.

A second modification was made by putting an ordinary three-way tap at D. This was useful for several reasons. At the end of each experiment the whole of the hydrate solution was allowed to run in through the tap B. The three-way tap D could be so turned as to prevent the hydrate running along to the tap K, and so on to the mercury. The aldehyde was contained in a vessel C, with a tap C′. B is a single tap through which the oxygen was introduced. H is a mercury seal covering the junction of the pipette J with the mercury manometer. G is a trap to prevent air reaching the apparatus, if any should happen to leak through the indiarubber tubing. E is a barometer-tube, the lower constricted part of which dips under the mercury in L. The vessel L and the barometer-tube E could be moved up and down. The pressure in the vessel was thus obtained by reading the difference in height between K and E.

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Method of Procedure.

In each experiment the mercury was first brought up to the tap K, the limb opening to the air being free. The tap was then closed, and the barometer and vessel lowered until the mercury in the barometer-tube was below the level of the tap K. The liquid aldehyde was then introduced into the vessel C, and the tap closed from the apparatus. The tube M was then connected with a vacuum pump, and with the tap D turned to connect all the ways with B the whole was evacuated. The tap B was then closed, and K turned to connect the mercury with the apparatus. Aldehyde was then allowed to distil into the reaction-vessel, which it did very readily, the mercury all the time being kept as near to K as possible. The vessel was then evacuated again, and this repeated three or four times, so as to completely replace the air by aldehyde. The final pressure of the aldehyde was taken, and the tap D turned so as to disconnect the right-hand part of the vessel, the reaction-vessel being only connected with the capillary tubing DLK The tap K was also closed. The capillary tubing DBM was then evacuated by means of the pump, and the tap B closed. M was then connected with a tube delivering pure oxygen dried over calcium-chloride. If the total pressure in the reaction-vessel was not required to be more than atmospheric pressure, the tap D was opened so as to connect all three ways, and, as soon as the oxygen had entered, the tap D was again turned, so as only to connect the tube DNK with the reaction-vessel, and the pressure read off. The tubing DB thus did not form part of the reaction-vessel.

If the pressure in the reaction-vessel was required to be greater than atmospheric pressure, the tap D was first turned so as to connect the tube DNK with the tube DB and disconnect the vessel A containing aldehyde. The mercury L was then lowered and the tap K turned. Oxygen was thus drawn over into the pipette J without aldehyde. The tap D was then turned to connect the vessel A with the pipette J, and the mercury head raised, driving the oxygen back into the reaction-vessel A. In this way any pressure obtainable on the gauge could be obtained in the reaction-vessel. The volume, also, not at the temperature of the bath was only that of the capillary tubing DNK.

At the end of each experiment the capillary tubing BD was first evacuated, and the tap B closed. Sodium-hydrate solution of known specific gravity was then allowed to fill the capillary tube. The vessel containing the hydrate was then weighed and again placed under B, the taps B and D being turned so as to let the hydrate into the reaction-vessel. The beaker containing the hydrate was again weighed, and the loss was that due to the hydrate drawn in. In this way the capillary tubing BD was filled with hydrate before both the first and second weighing, and hence introduced no error.

In the case where the final pressure in the reaction-vessel was greater than atmospheric pressure, the pressure was diminished by lowering the mercury head, and drawing some of the gas over into the pipette J, and finally driving it back when the hydrate had been introduced. The hydrate in a few minutes polymerized the whole of the aldehyde remaining, as the pressure soon became constant. Allowance had to be made for the vapour-pressure of the hydrate let in in determining the final pressure of oxygen.

When the whole of the aldehyde had been polymerized, the tap D was turned to disconnect DB, and a portion of the remaining oxygen and

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nitrogen was drawn over into the pipette J by lowering the head and transferred into burette, and analyzed with alkaline pyro-solution in a Hempel pipette, to obtain the percentage of nitrogen present.

From the weight of hydrate drawn in, and the change of pressure after the hydrate had been let in, the quantities of aldehyde and oxygen at the end of the experiment could be calculated.

The volumes of the reaction-vessels used were between 90 and 100 cubic centimetres.

Importance of Cleaning the Reaction-Vessel.

Mention is made by Dr. Ewan of the fact that it seems to be of some importance to keep the apparatus as clean as possible. This was found to be of the utmost importance. The chief difficulty experienced with these experiments was that the aldehyde tended to decrease faster than it should, according to the equation 2C2H4O + O = 2C2H4O2, this being due, no doubt, to polymerization taking place. Aldehyde polymerizes readily in the presence of mineral acids, bases, and many salts, especially sodium-acetate, so that it is of the utmost importance that these should be absent. No reliable correction can be made for polymerization, and any experiment in which it was marked had to be rejected as practically worthless. At the same time, the presence of the polymeric form seemed to have a distinct retarding effect on the action, as will be shown later. For this reason, and to make the conditions strictly comparable, the reaction-vessel and the whole of the capillary tubing and vessel containing the aldehyde were invariably left to stand at least twelve hours, filled with a solution of potassium-permanganate and hydrochloric acid. The manganesedioxide was removed by oxalic acid, and everything washed out with distilled water with the utmost care, and dried in a current of hot air.

Correction for Aldehyde Dissolved in the Acid Formed.

By means of a factor k′, calculated from the excessive decrease of aldehyde as determined at the end of the experiment, on the assumption that this decrease is due to the aldehyde being dissolved in acetic acid according to Henry's law, Dr. Ewan has attempted to make correction for this error. The value of this correction seems to me to be doubtful. In the first place, the value of k′ for any one temperature should be constant, whereas the value of k′ varies from 0·00103 to 0·002873—that is, it is in one case almost three times as great as in another—while the temperature has only varied from 20·2° C. to 20·6° C. Further, this value of k′ at 20·8° C. is 0·001767, or almost twice as great as at 20·6° C. This in itself is sufficient to show that Dr. Ewan's constant is a very uncertain factor. There can be little doubt that where the value of k′ is at all great some of the decrease in the aldehyde has been due to polymerization.

Further, it is extremely unlikely that when such a vapour as that of aldehyde, so near its condensing-point, had once dissolved in the acetic acid under the fairly large pressures at the beginning of the experiment it would vaporize again as the pressure decreased.

In any case, in most of the experiments, the error due to the dissolving of aldehyde by acetic acid is small, because the quantity of acid formed is small, and the surface exposed is also small, as the acetic acid runs down into the drawn-off part of the pipette. In those cases in which

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the aldehyde showed a decided decrease in excess of the oxygen it was assumed that polymerization had taken place, probably on account of traces of impurity in the vessel, and the experiment repeated. The term k′ was therefore neglected, thus greatly simplifying the equation, and introducing no great error. The work of calculating the results when k′ is included is extremely laborious, and little is illustrated when once the principle of the action has been determined. The greatest error due to this will probably be noticed most in the final values of the constant, because there the pressure of the aldehyde is small, and an error of 4 or 5mm. will be most felt.

General Discussion of the Equation Used.

The equation for the direct oxidation of the aldehyde to acetic acid can be written either as 2C2H4O + O2 = 2C2H4O2, in the second case the assumption being that the action goes on between the oxygen atoms and the aldehyde molecules.

If the first equation were the correct representation, the action should proceed at a rate proportional to the square of the concentration of the aldehyde and to the pressure of the oxygen. The second equation suggests that the action proceeds at a rate proportional to the concentration of the aldehyde and to the oxygen atoms.

Everything points to the second equation being the correct one. The values of K′ worked out by this equation give, on the whole, good results, and although there seems to be disquieting differences between some experiments carried out under apparently the same conditions, yet this is probably due to there being some catalytic agent present affecting the action. For each experiment the value of K′ obtained is fairly consistent, and the experiments as a whole show a general consistency.

The equation used was practically the same as that of Dr. Ewan, except that k′, the factor allowing for the aldehyde dissolved in the acetic acid formed, was not taken into account.

Let the partial pressure of aldehyde at the commencement of the experiment be A millimetres, that of oxygen B millimetres, and that of nitrogen N millimetres. P is the total pressure of the gas at any instant. Suppose that after T minutes x millimetres of oxygen have combined with 2x millimetres of aldehyde to form acetic acid. The pressure of oxygen will then be (bx) and that of aldehyde (a — 2x). Also, there will be a certain pressure of acetic-acid vapour, which will be equivalent to the vapour-pressure of acetic acid at that temperature after the acetic acid formed has commenced to condense to liquid.

If, then, the action proceeds with a velocity proportional to the concentration or pressure of aldehyde and of the oxygen atoms, then -dp′/dt = Kp′½p2,
where p′ = partial pressure of oxygen,
p′2 = partial pressure of aldehyde,
dp′/dt = rate of change of pressure of oxygen.

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As Dr. Ewan has shown, the pressure of oxygen atoms should be proportional to the square root of the pressure of oxygen if the law of mass action holds strictly, and a state of equilibrium is assumed, for there is equilibrium between the oxygen atoms formed and the oxygen molecule. The equation therefore becomes K′C(o2) = C(o) × C(o) = C(o)2,
where C(O2) and C(O) represent the concentrations of oxygen molecules and atoms respectively.

From this, then, C(o) = K′C½(o2), or the pressure of the oxygen atoms varies as the square root of the pressure of the oxygen molecules. The actual value of K′ will undoubtedly be very small, as the number of oxygen atoms present can hardly be great even under the most favourable conditions.

In integrating the equation -dp′/dt = Kp′½p2
we put p′ = b - x;p2 = a-2x.


This gives -d(b - x)/dt = K (b - x)½(a-2x),
or dx/dt = K (b - x)½(a-2x),
and integrating it gives
Kt = -½(b - a/2)½ loge (b - x)½-(b - a/2)½/(b - x)½ + (b - a/2)½ + constant,
or, writing the logarithm to the base 10, we get
Kt/2·30 = -½(b - a/2)½ log10 (b - x)½ - (b - a/2)½/(b - x)½ + (b - a/2)½ + constant.


The constant is obtained from the condition that x = o when t = t, and the equation is therefore
K′t = -½(b - a/2)½ log10(b - x)½ - (b - a/2)½/(b - x)½ + (b - a/2)½ + ½(b - a/2)½ log10 b½ - (b - a/2)½/b½ + (b - a/2)½
where K′ = K/2·30.

The values of K′ given in the tables are calculated by means of this equation, and in each case are multiplied by 105.

Experimental Work.

There seems to be nothing more clear from my experiments than that the action is extremely susceptible to the influence of accelerating and retarding influences. Much time was lost in the early part of the work through the sample of aldehyde, prepared by repeatedly redistilling a stock sample, containing paraldehyde. The result of using this sample was that the action would not go on at a rate at all comparable with that which Dr. Ewan observed, though otherwise the conditions were the same. When compared with the results obtained later on with a pure sample the difference is at once apparent.

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The following are some results obtained by using this sample. There can be no doubt that there was the pressure of aldehyde in the pure form there, as the para form, boiling at 124° C., could not give a very great pressure at 21° C., but the presence of the para form is small quantities hindered the action.

The value of k′ is not calculated, as it is obviously worthless.

Time in Minutes. p2. p′. P. Temperature.
0 246 850 1096 20·7° C.
182 246 850 1096
0 250 513 763 20·7° C.
360 741
0 220 498 718 20·6° C.
160 715
0 252 374 621 20·7° C.
988 638
0 303 321 624 25° C.
300 593
0 318 279 627 25° C.
180 615
0 306 391 697 25° C.
720 671

These experiments show that there has been little action going on. At first the aldehyde was not suspected; but when the oxygen had been used undried, and dried over calcium-chloride and sulphuric acid, and when the temperature had been raised to 25° C., and still the action did not proceed at a satisfactory speed, it was thought that the fault was due to the aldhyde.

A second sample of aldehyde was prepared from the original sample by distilling it with dilute sulphuric acid, so as to break down the para form present.

Three distilling-flasks were used. The first contained the mixture of dilute sulphuric acid and aldehyde, the second contained calcium-chloride in fairly large quantities, and the third was empty. All the flasks and connecting tubes were carefully washed and steamed out before being used, so as to remove any trace of acid or base which might polymerize the aldehyde. The flasks were arranged in series.

The second and third flasks were immersed in ice, while the bath containing the first flasks was gradually warmed until the water boiled. The distillation was then stopped. The aldehyde distilled over very readily, being so volatile that some of it passed the second flask, and only condensed on reaching the third flask. Any water-vapour was condensed in the second flask and absorbed by the calcium-chloride. When the water round the first flask had boiled it was detached, and the aldehyde which had condensed in the second flask distilled into the third flask. The temperature was kept down to 25° C., and the latter half of the liquid in the second flask rejected. The purified aldehyde was then transferred into a well-stoppered bottle protected from light, which had been carefully steamed out. This sample held good throughout the series of experiments.

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Experiments Made To Determine Whether Such A Maximum Pressure As That Indicated By Dr. Ewan'S Experiments Did Exist.

In this connection the chief points to be cleared up were: (1) Whether a maximum reaction pressure did exist; (2) whether the pressure of aldehyde had any effect on it; (3) the effect of temperature on this maximum pressure, if it was found to exist.

At the outset it may be said that although the rate of reaction, as indicated by the value of K′, seemed to diminish with higher pressures of oxygen, the pressure of aldehyde remaining constant, yet there was nothing to show that the action came suddenly to an end, as is indicated by Dr. Ewan's experiments. Further, as will be shown later, on raising the percentage pressure of aldehyde the value of K′ again rose, even though the oxygen-pressure remained high. Some explanation from theoretical considerations, as will be shown later, can also be given of the fact that the value of K′ is small when there is a big pressure of oxygen and a small one of aldehyde.

It will be as well at this stage to quote some experiments to illustrate this point.

Except where otherwise stated, the oxygen, prepared from potassium-chlorate and manganese-dioxide, was dried over calcium-chloride. No attempt was made to absorb the acetic acid formed.

Time in Minutes. p2. p′. P. Temperature. K′.
0 349 414 783 20·85° C.
154 256 368 656 21° C. 2·2
203 219 350 601 21° C. 2·5
262 193 338 563 21·1° C. 2·4
318 171 327 530 20·8° C. 2·5
371 155 319 506 20·9° C. 2·4
444 135 309 476 20·9° C. 2·4
509 124 304 460 20·9° C. 2·3

Partial pressure of nitrogen = 20mm.

Partial pressure of acetic acid = 12mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 232 534 793 20·8° C.
198 176 506 721 20·8° C. 1·3
242 169 503 711 20·8° C. 1·2
284 162 499 700 20·8° C. 1·2

Partial pressure of nitrogen = 27mm.

Partial pressure of acetic acid = 12mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 229 598 850 20·8° C.
159 173 570 776 20·8° C. 1·4
826 69 518 622 20·9° C. 1·2

Partial pressure of nitrogen = 23mm.

Partial pressure of acetic acid = 12mm.

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All the other experiments quoted were made in semi-darkness; but the last one quoted was made in absolute darkness. No alteration is, however, noticeable in the reaction.

In the next two experiments the oxygen was dried with concentrated sulphuric acid instead of calcium-chloride. In the first the effect seems to be to show a retardation, but the value of K′ rises again in the second. The extreme sensitiveness of the action to determining agents makes it very difficult to have the conditions exactly the same. In any case there can be no doubt that an action has been going on.

Time in Minutes. p2. p′. P. Temperature. K′.
0 242 592 870 20·9° C.
166 213 577 838 20·9° C. 0·8
219 203 573 824 20·95° C. 0·73
281 194 568 810 20·95° C. 0·71
336 186 564 798 20·9° C. 0·71
459 172 557 777 21° C. 0·70
668 140 542 730 21° C. 0·88
1370 62 505 615 20·9° C. 0·9

Partial pressure of nitrogen = 36mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 240 765 1051 20·9° C.
173 203 747 1008 20·9° C. 1·2
284 157 724 939 20·9° C. 1·2
1030 95 693 846 20·9° C. 0·73

Partial pressure of nitrogen = 46mm.

In the next experiment the oxygen was again dried over calciumchloride :—

Time in Minutes. p2. p′. P. Temperature. K′.
0 250 851 1125 21° C.
165 199 826 1061 21° C. 1·1
193 191 821 1048 21° C. 1·1
221 182 817 1035 21° C. 1·1
249 172 813 1021 21° C. 1·1
449 125 790 952 21° C. 1·1

Partial pressure of nitrogen = 24mm.

The number of these experiments, in every case the action going on regularly puts it beyond doubt that the action is able to go on at pressures of oxygen above 530mm., at least under some conditions. In one

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experiment, to confirm that acetic acid was formed, at the end of the experiment, instead of letting in the hydrate, the reaction-vessel was taken out of the bath, and acetic acid was found in the bottom of the pipette. On breaking the end and letting the liquid out on blue litmus-paper the paper was turned red and the smell of acetic acid was very noticeable.

The results of these experiments are opposed to those of Dr. Ewan, for although the value of K′ shows a decrease with rise of pressure of oxygen there is nothing to show that the action does not go on at least up to very much higher pressures than were found by him It does not, on the face of it, seem probable that an action which went on readily when the pressure of oxygen was 450mm. should cease when the pressure was raised only another 100mm. If there are any oxygen atoms present with the pressure at 450mm., there seems to be no theoretical reason why they should not exist in at least almost as large numbers with the pressure at 550mm., allowing that the external conditions are the same. It seems to me extremely probable that some retarding influence was present of which Dr. Ewan was unconscious. As has been pointed out, in the earlier experiments little or no action went on, on account undoubtedly of the presence of the paraldehyde, especially when the oxygen-pressure was at all high.

Thus, comparing the last experiment quoted with one with similar pressures, only using a different sample of aldehyde, the difference is at once seen :—

Time in Minutes. p2. p′. P. Temperature.
0 246 850 1096
182 246 850 1096 20·7° C.

It is just possible that the sample used by Dr. Ewan was not free from paraldehyde. It must be noted also that the average value of his constant at 21° C., worked out by the same method as mine, is only 1·2, whereas mine is 2·2.

As will be shown in the next section, many other substances affect the speed of the action, showing how easy it would be for the conditions to be different when apparently the same.

Some Catalytic Agents Affecting the Reaction.

1. The effect of impure aldehyde has already been shown by comparing the results obtained with two different samples. Nothing more need be said at this stage.

2. Impurities in the Oxygen.—The oxygen was kept in bell jars over water. To this water considerable quantities of sodium-hydrate were added, to dissolve out any traces of chlorine, or oxides of chlorine, which are evolved when manganese-dioxide and potassium-chlorate are heated. On one occasion pure water was used in the jars, and the oxygen used immediately it had been collected, before the traces of impurities had dissolved. The effect of this trace of chlorine was almost to stop the reaction. The conditions of the experiment were otherwise the same as they had been.

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The following experiment illustrates this :—

Time in Minutes. p2. p′. P. Temperature.
0 255 435 720
120 252 434 717 20·9° C.

Partial pressure of nitrogen = 30mm.

On the next day, using the same sample of oxygen, which had meantime remained undisturbed over the water, the action went on a little more readily, though still much slower than formerly.

Time in Minutes. p2. p′. P. Temperature.
0 266 425 713 20·9° C.
145 680 20·9° C.
199 667 20·9° C.
271 652 20·9° C.
343 643 20·9° C.
1362 198 396 628 20·9° C.

Partial pressure of nitrogen = 22mm.

On shaking the water up to dissolve the chlorine, and allowing it to stand several days, on again using it the action went on with very great rapidity—greater, in fact, than that previously observed.

Time in Minutes. p2. p′. P. Temperature. K′.
0 258 500 782 21° C.
181 142 443 621 21° C. 3·2
230 110 429 575 21° C. 3·4
254 98 422 556 21° C. 3·7
307 78 414 528 21° C. 3·7
409 60 406 502 21° C. 3·3

Partial pressure of nitrogen = 24mm.

The speed of this reaction is somewhat remarkable. Nothing was noticed to account for it. It may be that more water-vapour was present than usual. These results show how susceptible the action is to disturbing effects, the impurities in the oxygen being undoubtedly the cause of the change in the velocity of the reaction in these cases.

3. Vessel.—The speed with which the action proceeds seems also to be influenced by the nature of the reaction-vessel. It has long been known that the form and nature of the vessel has considerable effect on the velocity of gaseous reaction and this affords another marked example of it. Two reaction-vessels were used, so as to enable an experiment to be made each

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day, as the vessel required at least twelve hours for thorough cleaning. Hence, if only one were used, an experiment could not be left to proceed all night and then the same vessel used again the next day; but with two, every day could be made use of. No influence due to the first two vessels was noticed, but towards the end of the series of experiments one of them was broken. The new pipette showed a marked and consistent retarding effect as compared with the old one. Three experiments altogether were made in the new vessel, and corresponding ones in the old one, the effect in each case being the same. The temperature at which each pair was made was different, so that a wide range of conditions was covered.

The first of the following pairs of experiments quoted was made in the new vessel and the second in the older vessel.

Time in Minutes. p2. p′. P. Temperature. K′.
0 255 571 857 40° C.
109 160 527 753 40° C. 3·6
164 149 522 737 40° C. 2·7
223 145 519 730 40° C. 2·2

Partial pressure of nitrogen = 31mm.

Partial pressure of acetic acid = 35mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 254 512 798 40° C.
109 159 467 693 40° C. 3·9
142 140 457 664 40° C. 3·8
218 105 441 613 40° C. 3·7
405 61 419 547 40° C. 3·3

Partial pressure of nitrogen = 32mm.

Partial pressure of acetic acid = 35mm.

The second pair of experiments was made at 48·2° C. The same general effect is noticeable—namely, a fairly rapid falling-off of the speed of the reaction in the new vessel, and a much smaller value of K′ than with the old vessel.

Time in Minutes. p2. p′. P. Temperature. K′.
0 261 507 795 48·2° C.
140 132 447 657 48·2° C. 4·35
170 128·6 445·4 652 48·2° C. 3·76
200 126 444 648 48·2° C. 3·3
232 124 443 645 48·2° C. 2·8

Partial pressure of nitrogen = 27mm.

Partial pressure of acetic acid = 51mm.

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Time in Minutes. p2. p′. P. Temperature. K′.
0 267 496 789 48·2° C.
71 138 437 653 48·2° C. 8·9
85 118 427 622 48·3° C. 9·3
100 108 421 606 48·3° C. 8·8
115 101 418 596 48·3° C. 8·3
175 87 411 575 48·3° C. 6·1
205 80 408 565 48·3° C. 5·7

Partial pressure of nitrogen = 26mm.

Partial pressure of acetic acid = 51mm.

The next pair of experiments was made at 21° C.

Time in Minutes. p2. p′. P. Temperature. K′.
0 254 846 1120 21° C.
200 219·3 828·7 1080 21° C. 0·55
222 212·6 825·4 1070 21° C. 0·63
254 207·3 822·7 1062 21° C. 0·63
285 200·6 819·4 1052 21° C. 0·63
315 195·3 816·7 1044 21° C. 0·67
486 173 802 1006 21° C. 0·68

Partial pressure of nitrogen = 20mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 250 851 1125 21° C.
165 199 826 1061 21° C. 1·1
193 191 821 1048 21° C. 1·1
221 182 817 1035 21° C. 1·1
249 172 813 1021 21° C. 1·1
449 125 790 952 21° C. 1·1

Partial pressure of nitrogen = 24mm.

The fact of these three sets of experiments showing a consistent difference in the rate of the reaction is very strong evidence that the difference was due to the vessel. The conditions otherwise were exactly the same, and the vessels were cleaned in exactly the same manner—namely, by leaving them in a solution of potassium-permanganate and hydrochloric acid for twelve hours at least, and then washing out with oxalic acid and distilled water.

4. Water-vapour.—It is extremely probable that, as water-vapour has the effect of increasing the dissociation of oxygen molecules, its presence would accelerate this reaction. It had been hoped that an experiment might be made in which both the aldehyde and oxygen were dried abso-

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lutely but, while little difficulty was experienced in drying the vessel and oxygen by leaving them in contact with phosphorous pentoxide for a considerable time, the same could not be done with the aldehyde. The most effective drying agents are strongly basic or acidic oxides, or metals, such as sodium and potassium, which show a strong affinity for water.

Now, aldehyde is polymerized rapidly by mineral acids, bases, and many salts. In fact, the utmost precautions must be taken to prevent polymerization, and the presence of the polymeric form, as has been shown, affects the reaction. On this account attempts to obtain a perfectly dried sample of aldehyde have so far failed.

Could such an experiment be made, the result would probably show a great retardation, if not a total cessation of the action.

In view of the number of experiments quoted showing the influencing effects of many substances on the speed of the reaction, little further support is needed to affirm that the action is very sensitive to catalytic agents of all kinds. The experiments of Dr. Ewan can, however, be further quoted to support this.

Four experiments, having their constants worked out, are recorded in the “Philosophical Magazine.” In the first experiment, with the temperature at 20° C., the value of K varies from 0·76 as a minimum to 1·27 as a maximum. This is a very low value, for which he gives no explanation. In the next experiment the value of K varies from 2·81 to 3·57; in the next from 0·96 to 3·18; and in another experiment, in which the initial pressure of oxygen was only 373mm., well below his maximum, at which the reaction goes on constantly, and the initial aldehyde pressure 178·5, the constant begins at a maximum of 2·82, and falls rapidly to 2·59, and then to 2·34—in fact, in this experiment the reaction between the two last readings was extremely slow, as will be shown by quoting the experiment :—

Time in Minutes. p2. p′. P. Temperature. K′.
0 178·5 373·0 559·3
353 122·8 346·8 488·8 20·8° C. 2·82
421 118·2 344·6 481·8 2·59
501 114·2 342·7 476·1 2·34

It will be noticed that in the last eighty minutes the total pressure fell only 5·7mm., whereas in the latter part of a previous experiment, with the oxygen-pressure at 409·1mm. and the aldehyde pressure 119·1mm., the pressure fell 95mm. in 700 minutes, although this latter case includes pressure which should give a much slower rate on the average.

For these variations Dr. Ewan gives no explanation. He mentions that it seems of some importance to keep the apparatus as clean as possible, but he does not say why.

If we accept the idea that the action goes on between the aldehyde molecules and the oxygen atoms, there are two separate actions the speed of which catalytic agents could affect: there is, first, the breaking-up of the oxygen molecule, and, secondly, the combination of the aldehyde with the oxygen atom. It seems to me most probable that in most cases the first of these two actions is the one most affected. Under the most favourable conditions there could hardly be many atoms of oxygen present,

– 47 –

and any agent which tended to lessen the number would have a very marked effect on the speed of the action. Water-vapour is an example of an agent which favours the formation of atoms, and it is possible that paraldehyde-vapour has the opposite tendency.

In any case, the extreme sensitiveness of the action would lend support to the assumption that it is an indirect one, taking place between the oxygen atoms and the aldehyde molecules.

The Effect of Different Percentage Pressures of Aldehyde on the Value of K′

It can be seen from a glance at the experiments quoted that the value of K′ is smaller when the oxygen-pressure is high, the aldehyde remaining constant. It was found, however, that on raising the aldehyde-pressure to something more nearly equal to that of the oxygen-pressure the value of K′ again rose, even when the oxygen-pressure was high. This is further illustrated by experiments made at 25° C. and 30° C. :—

Time in Minutes. p2. p′. P. Temperature. K′.
0 421 602 1060 20·8° C.
73 339 566 954 20·8° C. 2·2
130 278 541 868 20·8° C. 2·4
183 236 523 808 20·58° C. 2·2
231 207 510 766 20·8° C. 2·3
282 192 498 739 20·9° C. 2·3
309 181 493 713 21° C. 2·2
426 138 477 664 20·8° C. 2·0
582 112 465 626 20·8° C. 1·7
1286 60 438 547 20·8° C. 1·2

Partial pressure of nitrogen = 37mm.

Partial pressure of acetic acid = 12mm.

In this experiment, though the oxygen-pressure was 602mm. at the commencement, the action went on rapidly, the value of K′ being practically the same as in the first experiment.

Time in Minutes. p2. p′. P. Temperature. K′.
0 298 475 790 24·8° C.
96 249 451 732 24·8° C. 2·1
154 215 434 681 24·8° C. 2·2
220 182 418 632 24·8° C. 2·2
259 165 410 607 24·8° C. 2·3
305 147 402 581 24·8° C. 2·3
435 137 397 566 24·8° C. 2·4
502 119 388 539 24·8° C. 2·3

Partial pressure of nitrogen = 17mm.

Partial pressure of acetic acid = 15mm.

– 48 –
Time in Minutes. p2. p′. P. Temperature. K′.
0 260 745 1039 25° C.
176 190 711 950 25° C. 1·4
407 158 695 902 25° C. 1·01
478 151 692 892 25° C. 0·91
1325 98 665 812 24·7° C. 0·63

Partial pressure of nitrogen = 34mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 434 691 1165 25° C.
78 333 650 1038 25° C. 2·4
106 320 635 1010 25° C. 2·4
139 294 621 970 25° C. 2·6
220 267 608 930 25° C. 2·6
252 239 593 887 25° C. 2·9
390 145 546 746 25° C. 2·9
1106 36 493 584 24·7° C. 2·06

Partial pressure of nitrogen = 40mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 253 469 734 29·9° C.
45 218 452 702 29·9° C. 3·3
65 203 444 679 29·9° C. 3·6
86 187 436 655 29·9° C. 3·6
130 153 420 605 29·9° C. 3·7
200 110 404 556 30° C. 3·8
444 63 375 470 29·9° C. 3·3

Partial pressure of nitrogen = 12mm.

Partial pressure of acetic acid = 20mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 259 692 987 30·1° C.
66 212 667 934 30·1° C. 1·7
144 173 649 876 30·2° C. 1·8
168 166 646 866 30·2° C. 1·7
239 151 639 844 30·2° C. 1·5
307 136 632 822 30·2° C. 15

Partial pressure of nitrogen = 34mm.

The regularity of this variation calls for explanation. The constant K′ has been worked out on the assumption that the number of oxygen

– 49 –

atoms present is proportional to the square root of the oxygen-pressure, and that the action goes on only between the aldehyde molecules and the oxygen atoms. Now, according to the kinetic theory of gases, these oxygen atoms, if present, would be formed during the collision of two oxygen molecules, and the number formed would be proportional to the number of these collisions. At the same time, the number of molecules re-formed would increase as the number of atoms present increase, until a state of equilibrium would be reached. The value of K′ then, thus worked out, does not take into account the possibility of an oxygen molecule being split up on coming in contact with an aldehyde molecule, and perhaps, in some cases at least, oxidizing the aldehyde molecule at the same time. If this went on to any great extent, then the greater the percentage of aldehyde molecules to oxygen molecules present, the greater would be the excess of oxygen atoms present, caused by contact of aldehyde and oxygen molecules, over the calculated number. The effect, then, if this be a determining factor, of raising the oxygen-pressure and keeping the aldehyde constant would be to lower the value of K′; but on raising the aldehyde-pressure to something nearer to that of the oxygen the value of K′ would again rise. In other words, the value of K′ would be somewhat dependent on the percentage pressures of aldehyde and oxygen. In most of the experiments there seemed to be a more or less distinct lowering in the value of K′ as the proportion of aldehyde became less, on account of the more rapid decrease in the pressure of aldehyde.

At the same time, it must be remembered that at no time during the experiment is a state of equilibrium reached either between the formation of oxygen atoms from molecules and the re-formation of the oxygen molecules, or between the oxygen atoms and aldehyde molecules. Now, if the percentage pressure of aldehyde were great, then an aldehyde molecule would seize upon an oxygen atom in many cases as soon as it was formed —that is, before it had time to come in contact with another oxygen atom —and thus count in the determination of the equilibrium. Thus the fact that oxygen atoms are being removed continually from the sphere of action would prevent a state of equilibrium being reached, and could cause, for any particular pressure of oxygen, the number of oxygen atoms ready to combine with aldehyde molecules to be greater than that calculated from the assumption of a state of equilibrium. Also, the greater the pressure of aldehyde in comparison with that of oxygen, the more would this be felt. This, it seems to me, might be a very considerable factor, and it is hard to see why it should not have at least some effect.

There is another factor which might have some effect, though it would in all probability be slight, if noticeable at all. With higher pressures of oxygen the proportion of oxygen atoms to oxygen molecules would be less than with lower pressures of oxygen. For instance, suppose the pressure of oxygen molecules to be A millimetres in one case and B in another, the value of B being greater than that of A. The number of oxygen atoms in the two cases would be KA½ and KB½ respectively, and the proportion of oxygen atoms to oxygen molecules in the two cases would be K/A½ and K/B½ respectively—that is, in the second case, where B is greater than A, the oxygen atoms are more diluted with the oxygen molecules, which are for the purposes of the reaction inert. This diluting of the reacting substances—namely, the aldehyde molecules and the oxygen atoms—by inert oxygen molecules might have the effect of retarding the action. Unless, however, we assume that time is taken up on collision—and the assumption

– 50 –

of the kinetic theory is that this factor is very small—this dilution should have little effect in reducing the number of collisions between aldehyde molecules and oxygen atoms. It is known that in solution the effect of some inert substance, as far as retarding the reaction is concerned, is very slight, and the same probably holds in the gaseous state. The effect of this factor, if at all noticeable, would be to reduce the value of K′ as the oxygen-pressure increased.

These seem to be the only possible determining factors from the theoretical standpoint, and if, as Dr. Ewan found, the action goes on readily with the oxygen-pressure at 450mm., and ceases practically altogether with the pressure at 550mm., no theoretical explanation seems forthcoming; in fact, it is almost inconceivable that so sudden a break should occur.

The Effect of Temperature on the Reaction.

The effect of rise of temperature on the reaction is, as would be expected, to increase the value of K′. At first increase of temperature by a few degrees seems to have very little effect, but at 50° C. the aldehyde is oxidized very rapidly, practically the whole being converted into acetic acid in the course of a few hours. There can be no doubt that the degree of dissociation of oxygen atoms becomes greater as the temperature rises, and this, together with the increased speed of the molecules, would account for the increased rate of oxidation. It was not considered necessary to make experiments at a temperature higher than 50° C., as little further would be illustrated, the speed at that temperature being sufficiently great to measure accurately. In the first experiment at 21° C. the proportion of aldehyde is rather greater than in the other cases, giving perhaps a higher value of K′ than otherwise.

Time in Minutes. p2. p′. P. Temperature. K′.
0 349 414 783 20·85° C.
154 256 368 656 21° C. 2·2
262 193 338 563 21° C. 2·4
371 155 319 506 20·9° C. 2·4
444 135 309 476 20·9° C. 2·4
509 124 304 460 20·9° C. 2·3

Partial pressure of nitrogen = 20mm.

Partial pressure of acetic acid = 12mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 298 475 790 24·8° C.
154 215 434 681 24·8° C. 2·2
220 182 418 632 24·8° C. 2·2
305 147 402 581 24·8° C. 2·3
435 137 397 566 24·8° C. 2·4
502 119 388 539 24·8° C. 2·3

Partial pressure of nitrogen = 17mm.

Partial pressure of acetic acid = 15mm.

– 51 –
Time in Minutes. p2. p′. P. Temperature. K′.
0 253 469 734 29·9° C.
45 218 452 702 29·9° C. 3·3
65 203 444 679 29·9° C. 3·6
86 187 436 655 29·9° C. 3·6
130 153 420 605 29·9° C. 3·7
200 110 404 556 30° C. 3·8
444 63 375 470 29·9° C. 3·3

Partial pressure of nitrogen = 12mm.

Partial pressure of acetic acid = 20mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 267 496 789 48·2° C.
71 139 437 653 48·2° C. 8·9
85 118 427 622 48·3° C. 9·3
100 108 421 606 48·3° C. 8·8
115 101 418 596 48·3° C. 8·3
175 87 411 575 48·3° C. 6·1
205 80 408 565 48·3° C. 5·7

The fact that the rise in the value of K′ is not regular can be accounted for by the extreme difficulty of having the conditions so exactly the same as not to cause some slight variation in the speed.

Effect of Large Quantities of Nitrogen on the Reaction.

The following experiments may be quoted to show that the presence of nitrogen, even in large quantities, has very little effect on the reaction :—

Time in Minutes. p2. p′. P. Temperature. K′.
0 258 269 656
129 210 245 596 20·9° C. 2·2
318 141 212 494 2·4
411 118 202 461 2·4
1236 63 175 379 1·4

Partial pressure of nitrogen = 129mm.

Partial pressure of acetic acid = 12mm.

Time in Minutes. p2. p′. P. Temperature. K′.
0 304 286 1025 24·8° C.
73 268·6 267·4 984 24·8° C. 2·4
158 237·3 251·7 937 24·8° C. 2·0
188 227·4 246·6 922 24·8° C. 2·1
223 218 242 908 24·8° C. 1·9
283 208 237 893 24·8° C. 1·7

Partial pressure of nitrogen = 435mm.

Partial pressure of acetic acid = 15mm.

– 52 –
General Conclusions.

The general conclusions arrived at are :—

1. That acetaldehyde is oxidized directly to acetic acid, the action taking place between the aldehyde molecules and the free oxygen atoms present.

2. That, with certain modifications, the action can be said to proceed with a velocity proportional to the pressure of aldehyde and to the square root of the pressure of oxygen.

3. The two chief modifications to be remembered are: (a.) That at no time during the reaction has equilibrium been established between the oxygen atoms formed and the oxygen molecules re-formed, the result being that more oxygen atoms are available for combination than would be calculated from a state of equilibrium, and the extent of this excess depends on the percentage pressure of aldehyde. (b.) That it is possible that the collision between an oxygen molecule and an aldehyde molecule may result in the breaking-up of the oxygen molecule.

4. That the existence of a maximum pressure of oxygen above which the action does not proceed is doubtful. Any such maximum is certainly much higher than that indicated by Dr. Ewan. There was no cessation of action with a pressure of 850mm. of oxygen.

5. That the action is greatly influenced by catalytic agents of every kind.

6. The effect of rise of temperature is to increase greatly the speed of the reaction, the value of the constant at 50° C. being four times as great as at 25° C.

Preliminary Note on the White Substance Formed by the Interaction of Mercury, Aldehyde, and Oxygen.

Mention is made by Dr. Ewan of the fact that a mixture of aldehyde and oxygen attacks mercury, forming a white substance, and giving off an inflammable gas which is not absorbed by alkaline pyro-solution. There can be no doubt that the white substance is formed, and experiments were made to establish its nature and manner of formation.

No action was observed between the pure acetaldehyde and mercury.

On allowing oxygen to pass into the reaction – vessel, which had contained the mercury and aldehyde for three hours, the formation of the white substance was soon noted. In fifteen minutes an incrustation had formed on the surface of the mercury. After fifteen hours a considerable quantity of white substance was obtained.

If any inflammable gas unabsorbed by alkaline pyro-solution is formed, it is extremely small in quantity, and not at all as much as would be expected from the quantity of white substance formed.

In appearance and properties the white substance distinctly resembles mercurous acetate, and a considerable amount of evidence has been gained in support of the view that, as formed, it is mercurous acetate, an oxide of mercury being first formed by the action of the oxygen on the mercury, and this then attacked by the acetic acid.

The white substance carries small particles of mercury, which can be separated with very great difficulty.