Carbon Dioxide Content of Ethanolamine Solutions PAUL T. WAGNER and MARSHALL LEW lndostrial Laboratory, M a r e lsland Naval Shipyard, Vallejo, Calif. QLEOLS
ethanolamine solutions are extensively used in indus-
A txial processes designed to remove carbon dioxide from gas streams. I n investigative and control work it is often desirable to determine the carbon dioxide concentration in both the gas and the absorbing liquid. Available analytical methods leave a great deal to be desired in rapidity or simplicity, and a certain amount of accuracy can he sacrificed if these desirable characteristics can be obtained. The gravimetric Knorr alkalimeter method of determining carbon dioxide in solutions remains the most accurate available pi ocedure. Yariations of this method involve liberation of carbon dioxide TYith strong acid, and estimation of the amount of carbon dioxide by direct measurement of the amount of carbon dioxide liberated (5j, and by measurement of the pressure exerted by the liberated gas ( 1 ) . Other methods include precipitat ion as barium carbonate with either gravimetric or titrimetric estimation. illthough each of these methods is particularly useful from one point of view or another, a simple titrimetric procedure would he desirable. Foreman ( 2 ) titrated amine hydrochlorides directly in 85 to 90% ethyl alcohol using phenolphthalein as the indicator, but (lid not work viith amine carbonates. Lovgren ( 4 ) titrated ammonium acetate and ammonium chloride in 50% ethyl alcohol solutions with sodium hydroxide, using thymolphthalein as the indicator. Fritz ( S ) , in the titration of amine salts in ethylenediamine with strong base, was unable to titrate carbonates because of their insolublity in the solvent. The proposed method permits the reaction of amine carbonate with excess sodium hydroxide in aqueous medium. Back titration of excess base is performed in 75% ethyl alcohol solution, because ethanolamines :we neutral in this solvent.
son mith results obtained with the Knorr procedure, in which 1 to 4 sulfuric acid was used to release the carbon dioxide, and .\searite to absorb it. In the proposed method back-titration was atterhpted with aqueous sulfuric acid, but the end point was indefinite. Low values indicated a possible slight hydrolysis of sodium carbonate. The addition of sodium acetate to provide common sodium ioii and back-titration nith alcoholic acetic acid was found to give reliable results. Direct titration with alkali gave results as much as 30% low, as did titrations in which sodium hvdroxide was added in excess but not heated.
ML 0 5 Y A C E i l C ACID tN rrHANOL
Figure 1. Potentiometric Curves A.
B.
2.76 grams per 100 ml. of CO? in 3.3.V diethanolamine Standardization Visual end point
RE.4GEYTS AND SOLUTIONS
Sodiuni hydroxide, 0.5-\:, carbonate-free, standardized with potassium acid phthalate. Sodium mxtate crvstals. C.P Aiizarin Yellow R:0.2% in 9.5% ethyl alcohol. Acetic acid, 0.LV in 95p0 ethyl alcohol, standardized with sodium hydroxide as in the procedure for determining carbon dioside.
Table I. Ethanolamine
Carbon Dioxide in Ethanolamines Proposed Method COz, G r a m d l 0 0 MI. Bnalyst .4 Analyst B
Knorr Alkalimeter COz, Grams/100 111 Analyst h Analyst B
PROCEDURE
1~:sactl~35 ml. of standardized sodium hydroxide are transfrrred into a 500-mi. wide-mouthed Erlenmeyer flask containing about 5 grams of sodium acet'nte and 5 ml. of aqueous ethanolnniine sample. .4fter henting t,o boiling on a hot plate the s2mple is cooled and 120 ml. of 95% ethyl alcohol added, plus 1 ml. of alizarin yellow indicator. The solution is titrated to a light orange end point (similar to salmon-pink) with alcoholic acetic acid. The titrant is st'andardized about once a week in the same manner, leaving out the ethanolamine. h s the end point is somewhat indefinite, it is advisable to t,itrate to the s:ime color as that obtained in standardization. If the amine is highly carbonated it may be necessary to reduce the sample size t,o less than 5 ml. in order to maintain a definite excess of sodium hydroside. Calcu1:ttion is as follows:
5
K OCO?added
In order to establish the end point, potentiometric titrations were made with a Beckman Model G p H meter, using a silver reference electrode and a glass indicator electrode. The silver electrode was cleaned with dilute nitric acid before use. Reagents were added in the same manner as described above except that 1 gram of lithium chloride was added with the sodium acetate, to stabilize the reference electrode. Figure 1 presents two characteristic potentiometric curves. Curve A shorn titration of escess sodium hydroxide in a sample containing 2.76 grams of carbon dioxide per 100 ml. of 3 . 3 5 diethanolamine. Curve B shows the titration of sodium hydroxide against the acetic acid reagent. The smsller change in millivolts per milliliter of titrant of curve A after the end point is attributed to hydrolysis of sodium carbonate. The solution became light orange in color near the point where a maximum change in potential per milliliter of titrant was observed.
Grams of carbon dioside per 100 nil. of absorbing solution = (3.5 -~ X S of NaOH) - i m l . titration X .V of H 0 . h ) x 2.2 ml. of sample EXPERIMENTAL
RESULTS
Carbonate solutions of mono-, di-, and triethanolamine were obtained by bubbling carbon dioxide into 3 N aqueous amine solutions. The method was applied to diethanolamine solutions having different degrees of carbonation and to mono- and triethanolamine a t only one carbon dioxide concentration for each. The accuracy of the proposed method was evaluated by compari-
Results obtained by two analysts and by two methods are shown in Table I. The first line gives the results obtained with no added carbon dioxide, and probably represents the slight carbonation that had taken place due to absorption of carbon dioxide from the atmosphere. 575
576
ANALYTICAL CHEMISTRY
Comparison of the values obtained by the authors' method with those obtained using the Knorr alkalimeter indicate that the proposed method is accurate within the limits of its reproducibility. DISCUSSION
The reaction of the ethanolamine carbonates with excess sodium hydroxide, when heated, can be represented as follows, using monoethanolamine as an example:
+ 2iYaOH --+ Xa2C03 + 2HOCH2CH2SH30H HOCH2CHJVH3HC03 + 2NaOH --+ NaZCO:, + HOCHZCHZNH~OH + H20 (HOCH2CH2iTHs)J3Os
I n either reaction 2 moles of sodium hydroxide are equivalent t o 1 mole of carbon dioxide, whether the carbon dioxide is present as carbonate or bicarbonate. I n back-titrating the excess sodium hydroxide, both hydrolysis of sodium carbonate and ionization of the amine must be repressed. Since ethanolamine solutions in ethyl alcohol are neutral to alizarin yellow, the basic effects of the amine are eliminated by performing the titration in a predominantly alcoholic solution. The low solubility of sodium carbonate in the alcoholic solution is further decreased by the addition of sodium acetate, thus decreasing the concentration of hydroxyl ion. The addition of sodium acetate appears to
decrease the sharpness a t the end point slightly, but still increases the reliability of the method by suppressing hydrolysis of sodium carbonate. .4lthough the end point is slightly indefinite, errois from this source are minimized by titrating the sample to thr same indicator color as that obtained in the standardization of the acetic acid. For a 5-ml. sample a 0.1-ml. net titration represents about 0.02 gram of carbon dioxide per 100 ml. of solution. Results are computed to the nearest 0.01 gram per 100 ml. Even though the percentage error may be high a t low concentrations, the ahsolute value can be expected to have a probable error of about +0.04 gram of carbon dioxide per 100 nil. of sample. R'here a very high degree of accuracy is not desired, the eimplicity of titrimetric analyses makes this method desirable for t h r estimation of carbon dioxide in ethanolamine solutions LITERATURE CITED
(1) Dixon, R. E., and Williams, R. A., Analyst, 74, 360 (1949). Biochem. J., 14,451 (1920). (2) Foreman, F. W., (3) Fritz, J. S., - 4 x a ~CHEY., . 24,306 (1952). (4) Lovgren, S., 2. anal. Chem., 64, 457 (1924). (5) Reed, R. M., and Wood, W. R., Trans. Am. lnst. Chem. E?tgrs.. 37, 363 (1941). RECEIVED for review M a y 26, 1953. Accepted November 12, lQ53. T h e views expressed in this paper are those of the authors and are not t o be construed as the official views of the S a v y Department.
Test for Acyclic Carboxylic Acid Anhydrides DAVID DAVIDSON Department o f Chemistry, Brooklyn College, Brooklyn 10,
x
u
OFTEN reported characteristic test for or-amino acids consists of their reaction with p-nitrobenzoyl chloride in the presence of bases, which results in the appearance of a blue or purple color (3,4,6, 7 ) . The chemistry of this test has been elucidated by Karrer and Keller ( 5 ) ,who showed that in the case of leucine the test depends upon the successive formation of an 5-p-nitrobenzoy1 derivative ( I ) and its azlactone (11). The latter is a weak acid which possesses indicator properties, its anion (111) being blue in pyridine solution. The present author had also reached this conclusion from theoretical considerations and was engaged in a study of this reaction with David Perlman when the work of Karrer and Keller (6) appeared. At that time it was found that the blue color obtained by Waser's procedure ('?), which ordinarily fades more or less rapidly, could be regenerated not only by adding more pnitrobenzoyl chloride but also by adding acetic anhydride. This suggested that the fading and reappearance of the blue color was due to the opening and closing of the azlactone (I1 I).
RCHCOOH I kHCO
RC-CO
RCH-CO I \
I
\
I
N. Y. necessary any stronger base than the pyridine which might conveniently be employed as the solvent for the reagent. This expectation was realized. PROCEDURE
To 1 ml. of a 3% solution of or-(p-nitrobenzoy1)-amino-@toluic acid (4)in pyridine add 30 ml. or 1 drop of the substanor to be tested. With acyclic carboxylic acid anhydrides a blue color develops within a few seconds. This color slowly fades. Acids as strong as salicylic acid inhibit the color formation, b u t this may be overcome by adding a few drops of triethylamine. Under these conditions carboxylic acid chlorides also respond t o the test if a droplet of water is added. The transitory nature of the color makes this test unsuitable for quantitative work ( 8 ) . Positive results were obtained with acetic, propionic, butyric, isobutyric, capric, lauric, myristic, palmitic, stearic, benzoic, anisic, p-bromobenzoic, phenylacetic, 8-phenylpropionic, cinnamic, a-amylcinnamic, polyadipic, and polyazelaic anydrides. Negative results were obtained with maleic, succinic. glutaric, phthalic, and 3-nitrophthalic anhydrides. The azlactones ( a type of anhydride) thus appear to stand between the acyclic and cyclic anhydrides. The lower reactivity of cyclic anhydrides is also evidenced by their lower heats of hydration ( 1 ) . Positive tests were also given by aspirin and o-benzylouybenzoic acid, an anomaly which is discussed elsewhere (9). LITERATURE CITED
(1) Conn, J.
I
s
0 2
I
I1
I11
It seemed likely that if R in the a-amino acid were an acidifying group such as phenyl, the corresponding p-nitrobenaoyl derivative would be a useful reagent for detecting acid anhydrides since the increased acidity of the azlactone (11) would make un-
B.,Kistiakowsky, G. B., Roberts. R. AT., and Smith, G. A , J . Am. Chtm. &c., 64, 1747 (1942). (2) Davidson, D., and Auerbach, L.. [bid., 75, 5984 (1953). ( 3 ) Edlbacher, S., and Litvan. Fr., 2. physiol. Chem., 265, 241 (1940). (4) Ingersoll, A W., and iidams, R., J Am. Chem. SOC.,44, 2930 (1922). ( 5 ) Karrer, P., and Keller, R., H d u . Chim. -Ida, 26, 50 (1943). (6) Kurts. A. C., Proc. SOC.Ezpt2. Bid. M c d . , 46, 339 (1941). (7) Waser, E., Mitt. Gebiete Lebensm. u . Hug., 20, 260 (1929). (8) Weiss, XI., private communication. RECEIVED for review August 20, 1953.
Acoepted November 3, 1953.
