Determination of Aldehydes Using the Reaction with Primary Amines SIDNEY SIGGIA AND EILEEN SEGAL’ Central Research Laboratory, General Aniline 6% Film Corp., Easton, P a .
HE formation of the Schiff base as basis for determining aldeT h y d e s has not been used to any great degree. This reaction undoubtedly has been considered usable by many people, since it is rapid and the excess of amine used should be readily determinable. However, several problems arise xhen one tries this reaction to determine aldehydes, and these problems could account for the little that has been published. These obstacles have been overcome to some degree, and the reaction can be applied to determine some aldehydes with a good degree of accuracy and precision. The method is also fast and is not involved. The method has some other attributes. I t is run in a nonaqueous medium which permits solution of the more insoluble samples; it can be used in the presence of some ketones; and it employs a reaction in an alkaline medium which is desirable in the handling of samples such as acetals, which liberate aldehydes in acid media. I n previous ivork ammonia has been used to determine formaldehyde ( 1 , 3, 4,6). Aniline has also been used to determine formaldehyde nephelometrically ( 7 ) . Phenylenediamine has been used for colorimetric determination of some a,p-unsaturated aldehydes (8). Interferences consist of acids which are approximately as strong as, or stronger than, the salicylic acid used in the titration. -4weak acid will not interfere, because the stronger salicylic acid will titrate the amine away from it. ( I n this case, acids are classified only as “weaker” or “stronger” since, in a nonaqueous medium, the dissociation constants of the various acids are not knonm.) Acid anhydrides and acid halides interfere in that they consume amine. The reaction is as follows:
RlNHz
+ RCHO
+
RCH=SHRi
free amine and ketone. The titration curves obtained in these cases are flattened, indicating that free amine is liberated as the excess is titrated. I n the case of some ketones, the break in the curve is completely obscured. Methyl ketones (acetone and methyl ethyl ketone) and cyclohexanone are in the latter category. With diethyl ketone, however, a distinct break is visible in the curve for the total amount of amine introduced. With ketones present in the aldehydes, the same holds true; methyl ketones and cyclohexanone will obliterate the end point through their weak tying up of amine, but the higher ketones do not affect the final results for the aldehyde, though they do diminish the intensity of the break. If a small amount of aldehyde is present in a large amount of ketone, it would be presumed that the break might be diminished to the extent of being completely obscured.
6
+ HLO
There is a definite equilibrium present in this reaction, so that a nonaqueous system had to be used to get the reaction far enough to completion to be usable. The only water present is that formed in the reaction. The Schiff bases formed deteriorate with strong acids, and one would titrate the original amount of amine put in the system if the mineral acids were used. It is for this reason that salicylic acid is used to titrate the excess amine in this procedure. Salicylic acid has a dissociation constant of 1 X IOp3 (in viater) and is strong enough to effect a good titration of the excess amine without causing the Schiff base to hydrolyze back to the free amine and aldehyde. (This behavior could be looked upon as connected with the equilibrium present in the reaction. The strong acids tie up the free amine very effectively, causing the equilibrium to shift to the left, whereas the salicylic acid does not bind the amine strongly enough to cause the reaction to reverse itself.) The Schiff base itself is a much weaker base than the original amine, so that it causes no trouble in the titration. Only aliphatic amines are usable with this method since titration is done with a relatively weaker acid than is normally used. ilromatic amines are too weak to be effectively titrated with salicylic acid. Lauryl amine is used in the procedure described. This was finally chosen because of its high boiling point and its availability. Butyl amine was the first amine tried and worked satisfactorily, except that due to loss of butyl amine from the reagent through evaporation, repeated standardization was neressary. From the titration curves obtained (Figure 1), it is evident that ketones undergo a reaction with the aliphatic amines to the corresponding Schiff base, but the equilibrium is in favor of the 1 Present address, Polychemicals Department, E. I. d u P o n t de Semours & Co., Inc., Wilmington, Del.
16
18
PP
20
24
26
98
30
MILLILITERS
Figure 1. Titration Curves
_-
-
benzaldehyde - - benzaldehyde plus diethyl ketone, denzaldehyde plus cyclohexanone
. . ..
This method does not work on aliphatic aldehydes, except for the formaldehyde, where it works well. From the titration curves obtained it is concluded that the Schiff bases of the aliphatic aldehydes and the aliphatic amines are not stable enough
Table I. Determination of Aldehl-des
Formaldehyde
Primary Amine Method Time Sample 70 Found 1.5 hr. ), 36.5
Benzaldehyde
Over week end 1 hr.
Salicylaldehyde
y;t&ypn’.
1 hr
1 hr.
J
weighed in. 0.72 Mole weighed in.
36.1 98.7 99.2
0.02 Xole weighed ?ut
99.6 99. 8 99.2 99.4 99.0
7oFound, Other Methods 36.3a 96.3b 99. O b
Aliquot portion used containing 0.02 Mole 0.02 Mole 99.3 98.6~ Cinnamaldehyde 1 hr. weighed out 99.9 0.02 Mole 99.6 99.3b Furfural weighed out 98.0 98.8 Bisulfite method of Siggia and Maxcy (6). Usual experimental error of this method is of the order of i l % of value stated. b 2 4-Dinitrophenylhydrazone method of Iddles, Low, Rosen and H a r t (a). , Results obtained are usually low by 0 t o 2 % , due t o slight s’olubility of precipitate.
..
(I
830
V O L U M E 25, NO. 5, M A Y 1 9 5 3 or, better, that the equilibrium in these cases is too far to the left for the reaction to be used for analytical purposes. No inflection points in the titration curves could be obtained for acetaldehyde, propionaldehyde, or butyraldehyde. Cinnamaldehyde is peculiar in this respect in that it is essentially an aliphatic aldehyde with the phenyl group on the chain, yet it will react completely enough to be used. Crotonaldehyde, which is similar to the cinnamaldehyde except that the phenyl group is replaced with a methyl group, does not show any titration break a t all. Since aldehydes are difficult to obtain in a pure form, difficult to purify, and difficult to keep in a pure form, this study was limited to the aldehydes indicated in Table I and in the text. The aldehydes chosen were those which occur most commonly and which should adequately illustrate the evtent of applicability of the method.
831 solution is added, the cover is replaced, and the mixture is shaken for a few minutes. The mixture is allowed to stand for 1 hour. The solution is then transferred with ethylene glycol-2-propanol to a 250-ml. beaker, and the excess lauryl amine is titrated potentiometrically with 1 N salicylic acid. A blank is also run. CALCULATION
(Blank - titration) X ?jaalicv X M . K . X 100 gramsample X 1000
PROCEDURE
A 0.02 mole sample of the aldehyde is weighed into a 100 0; 150-ml. glass-stoppered flask. Exactly 20 ml. of the lauryl amine
%aldehyde
ACKNOWLEDGMENT
AcknoTTIedgment is made to Dale Eichlin who ran some of the check experiments to establish the validity of the method. LITERATURE CITED
REAGENTS AND EQUIPMEhT
Lauryl Amine Solution. Two moles of lauryl amine weredissolved in 1 liter of ethylene glycol-2-propanol mixture. Lauryl amine was purchased as Armeen 12D from the Chemical Division of Armour & Co. Salicylic Acid. A 1 N solution in ethylene glycol-2-propanol was standardized against alcoholic sodium hydroxide. Equipment. Glass, calomel electrodes were used, along with a model H-2 Beckman pH meter.
=
(1) Foschini, A, and Talenti, hI.,2. anal. Chem., 118, 94-7 (1939). (2) Iddles, H. A , Low, A. W., Rosen, B. D., and Hart, R. T., IWD. ESG.CHEM.,ANAL.ED.,11, 102 (1939). (3) Legler, H., Pharm. Ztg., 76, 1063 (1931). ( 4 ) Meyer, Wr., Ibid., 74, 771-3 (1929). ( 5 ) Siggia, Sidney,and >Iaxcy, William, ~ A L CHmr., . 19,1023 (1947). (6) Smith, C. E., Am. J . Pharm., 1898, 86. (7) Toussaint, G., DBtrie, J., and VBrain, lI.,Compt. rend. SOC. biol. 117, 193-4 (1934). (8) Wearn, R. B., N u r r a y , W. XI., Ramsey, ll., Chandler, N., XSAL. CHEST.,20, 922 (1948). RECEIVED for review October 18, 1952. Accepted J a n u a r y 28, 1953.
Direct Volumetric Assay of Sodium Borohydride and Potassium Borohydride SAUL W. CHAIKIK Stanford Research I n s t i t u t e , Stanford, CaE;,f. of compounds, including the borohydrides, for hydride hydrogen has conventionally been carried out by causing reaction x i t h an alcohol, water, or dilute aqueous acid, and measuring the volume of hydrogen gas evolved a t constant pressure, or the pressure at, constant volume. h recently reported volumetric method (3)of sodium borohydride assay has employed oxidation with an escess of st,andard iodate and, after conversion of the excess to iodine, back titration with thiosulfate. The method described herein involves the direct titration of a sodiuni or potassium borohydride solution n-ith standard sodium hypochlorite, using Bordeaux red as an indicator. The stoichiometry is in accord with the following equation:
PROCEDURE
NALYSIS
A
H+
+ BH4- + 4 OC1- +H3BOj + 4C1- + H?O
The pH of the solution during the titration is critical, having an optimum value in the range 9.6 t,o 10.3. Carbonate buffer is used to maintain the proper pH.
Five milliliters of 2 S sodium hydroxide is introduced into a dry 100-ml. volumetric flask using a funnel to avoid wetting the upper neck with the base. The flask (without stopper) and contents are tared, a quantity of sodium or potassium borohydride about the size of a large pea is added, and another weighing is made. The m i g h t of borohydride should he between 0.1 and 0.3 gram. The volume is made up to 100 ml. with distilled water.
Table I.
Analysis of Sodium Borohydride and Potassium Borohydride
(Comparison of hypochlorite method and hydrogen evolution method) % Borohydride Hydrogen Evolution Sample Hypochlorite KaBHa SaBHa IiBHa KBHa
84 3 , 83.9, 93.9, 92 8 ,
83.8, 83.6 83.7 93.6 92 3 , 92.G
85.0, 84.4, 95.3, 93.3,
84.6 82.9 93.1 94.2
REAGENTS
Sodium Hypochlorite, 0.1 N . A convenient method is to dilute 115 nil. of a commercial 5Tc sodium hypochlorite solution to about 2 lit,ers nit,h nater. Clorox was used, although there seems to be no reason why other brands n-ould not be sat'isfactory. The solution was standardized against standard arsenite usingBordeaux red ( 1 , 2 ) . Carbonate Buffer Mixture. To a mixture of 50 grams of potassium bromide and 100 grams of sodium bicarbonate in about 600 ml. of water n-as added 6 N sodium hydroxide (about 130 ml.) until the nH as 9.6 ( n H meter). The solution Tvas diluted to about 1 liter. Sodium Hydroxide, 2 N . Bordeaux Red Indicator Solution. TKOtenths of a gram of Bordeaux red (Sational Aniline and Chemical Co., Inc.) x i s dissolved in 100 ml. of water. \A
Ten milliliters of the carbonate buffer mixture and 10 to 15 ml. of water are introduced into a titration vessel. A 5-nil. aliquot of the borohydride solution and two drops of Bordeaux red indicator solution are added, and the solution is titrated Fvith standard sodium hypochlorite. The red solution becomes increasingly blue during the course of the titration until just before the end point, when it is almost a pure blue. The end point is marked by a sharp change from blue to yelloJT-green. The per cent borohydride in the sample taken is given by: (111. of SaOCl X Si X mea. wt. of borohvdride X 100 grams of sample in 5-ml. aliquot
The milliequivalent weight of sodium borohydride is 0.004731 and of potassium borohydride is 0.006744.
