Spectrophotometric Titrations of Carbon Dioxide, Bicarbonate, and Carbonate Solutions A. L. UNDERWOOD and LYMAN H. H O W E 111 Department of Chemistry, Emory University, Atlanta 22, Go.
b Kinetic problems in the alkalimetric titration of carbon dioxide solutions are overcome b y the use of the enzyme carbonic anhydrase. After the titration has been made rapid in this way, instrumental end points based upon linear extrapolation become attractive to improve the precision. The spectrophotometric titrations described are based upon the ultraviolet absorption of the carbonate ion. These titrations eliminate indicator blank corrections and offer better precision than potentiometric or visual titrations. The same end point technique permits titrations of bicarbonate and carbonate as well as carbon dioxide-bicarbonate and hydroxide-carbonate mixtures. HE ALKALIMETRIC titration of carT bon dioxide solutions has been important for many years, but it remains vexatious and imprecise despite repeated attempts to improve it. The difficulty lies partly in the loss of carbon dixoide whenever the solutions are exposed to the atmosphere, but the titration is more intrinsically troublesome because of the unfortunate equilibrium and kinetic conditions attending the formation of carbonic acid from carbon dioxide and water (1, 6, 6). It has been shown recently that the enzyme carbonic anhydrase, which catalyzes the
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Figure 1. Potentiometric titration of 100 ml. of saturated carbon dioxide solution
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ANALYTICAL CHEMISTRY
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reaction, COZ H20 = HzC03, markedly improves the visual carbon dioxide titration by permitting a rapid approach to a permanent color change (8). Thus, the kinetic problem is overcome, and one may titrate rapidly, minimizing the loss of carbon dioxide without overrunning the end point. The precision still suffers, however, from the unfavorable equilibrium conditions, as may be seen from the potentiometric titration curve shom-n in Figure 1. I n such a situation, instrumental end points derived by linear extrapolation (conductometric, spectrophotometric, etc.) are generally more precise than visual or potentiometric ones (4, 9). I n the present case, conductometric titrations are unfavorable; the end point breaks are abrupt, but the ionic conductances are such that the changes in slope are too small for high precision. The carbonate ion exhibits an end absorption in the ultraviolet region of the spectrum which is much stronger than that of bicarbonate or carbon dioxide (see Figure 2). This makes possible spectrophotometric end point detection with improved precision in the carbon dioxide titration. I n the work reported here, measurements were performed manually. I n the presence of carbonic anhydrase, the reaction is so rapid that the absorbance does not drift noticeably after the addition of each increment of titrant; thus enough points to define the titration curve can be obtained almost as quickly as the visual end point. Automatic recording of the titration curve based upon a constant flow buret and a logarithmically-attenuated photometric output signal ( 7 ) is suggested as feasible. The Beckman energy recording adapter for the DU instrument might be useful in this connection. The carbonate absorption enables one to follox spectrophotometrically not only the carbon diovide titration, but alkalimetric bicarbonate titrations and acidimetric titrations of carbonate as n-ell. I n these cases. the reactions are rapid, and carbonic anhydrase is not used.
compartment of the Beckman Model
DU spectrophotometer was replaced
by a compartment designed to accommodate this beaker. The base of the compartment was a n aluminum block with a depression that held the beaker firmly in place. The instrument was elevated on a wooden platform so that a magnetic stirrer (with its cover removed) fitted under the titration compartment. The tip of the IO-nil. microburet penetrated a hole in the top of the compartment and dipped into the solution being titrated. This tip was painted black for a distance extending above and below its entry into the compartment to minimize light-piping, and the hole through which it entered was circumscribed with a dark felt gasket to make a flexible but light-tight fit. Titrations were performed at 235 mp in most cases. The enzyme carbonic anhydrase was obtained from the Worthington Biochemical Corp., Freehold, S. J. The solid enzyme was kept in a freezer as recommended. Portions of 0.1 mg. were used in each carbon dioxide titration, dispensed as aliquots of a freshly-prepared solution containing 1 mg. per ml. “Carbonate-free” 1-11 sodium hydroxide was prepared by eonventional means. A slight concentration of residual carbonate is manifested by the unobjectionable positive slope seen in Figure 3 where the curve should be more nearly horizontal. Sodium hydroxide solutions do not absorb a t the n-avelengths used in these titrations,
APPARATUS A N D REAGENTS
The titration vessel was a 150-ml. beaker of Corning Vycor 7910 glass, which has a relatively high ultraviolet transmittance. The light path through the beaker was about 5.5 em. The cell
Figure 2. Ultraviolet absorption spectra of 0.01M solutions of sodium carbonate, sodium bicarbonate, and carbon dioxide
Figure 3. Spectrophotometric titration of 100 ml. of saturated carbon dioxide solution at 235 mp but carbonates do. Consequently, the presence of carbonate as a contaminant in the titrant may result in a small positive slope in that portion of the curve corresponding to the addition of excess titrant. A saturated solution of carbon dioxide from which a1iquots:were titrated was prepared as described previously (8). All other solutions were prepared from reagent-grade chemicals using boiled water.
ducibility of the spectrophotometric end point was estimated in the following manner: Aliquots of a saturated carbon dioxide solution were diluted to 100 ml. and titrated as reproducibly as possible with regard to the time required and other factors. Absorbance readings were corrected for dilution, and the titration curves were plotted on large graph paper so that the over-all precision was not limited by the graphical process itself. A least squares straight line was calculated for a plot of milliliters of carbon dioxide solution taken us. milliliters of sodium hydroxide solution required to reach the first end point. This line was taken as “correct,” and the deviation of each sample from the least squares value was treated as an indeterminate error. Twenty-two samples, ranging from 10 to 100 ml. of saturated carbon dioxide solution, showed an average deviation of o.570.
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Carbon Dioxide Titrations. Figure 3 shows a typical spectrophotometric titration curve. Comparison with the potentiometric curve of Figure 1 shows clearly the superiority of the spectrophotometric end point, where the intersection of the two straight lines can be located more precisely than can the inflection point of the less rapidly changing potentiometric curve. The second end point, corresponding to complete conversion into carbonate, can scarcely be seen in the potentiometr ric curve, while spectrophotometrically, although it is not sharp, it)can be located by a reasonable extrapolation. Noting that the second ionization constant of carbonic acid is only about 5 X 10-11 (I),one may be satisfied with something less than high precision for the second end point, since it represents a titration that would normally not be undertaken by conventional means. It is difficult to assess the reliability of the spectrophotometric end point in the carbon dioxide titration because, apart from the end point problem itself, one must contend with the loss of carbon dioxide whenever the sample solution is exposed to the air. The rate of loss was lowered, but not to zero, by covering the sample solution in the titration cell with a layer of cyclohexane, chosen because of its ultraviolet transparency. I n one experiment where the effectiveness of cyclohexane was roughly evaluated, the loss of carbon dioxide from a stirred solution in a 15-minute period was approximately halved. The repro-
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Figure 4. Spectrophotometric titration of 100 ml. of 1.8 X 10-2Msodium carbonate solution at 235 m p Deviations found with visual titrations for this system vary with the operator, but are frequently more than 1%, with reduction to 0.7 or O.8yO if the titration is performed with exceptional care using a color match with a reference solution. It is probable that part of the deviation was due to nonreproducible loss of carbon dioxide and that the precision of the spectrophotometric end point was actually greater than indicated by the average deviation of 0.5%. Bicarbonate Titrations. Ten aliquots of a sodium bicarbonate solution were diluted to 100 ml. to give final concentrations of 4 X 10-3 to 3 X 10-2M. These solutions were titrated to the spectrophotometric end point with sodium hydroxide with an average deviation of about 3%. This titration is near the limit of feasibility even for the spectrophotometric end point. Goddu and Hume have pointed
out that the product of the concentration and the ionization constant of a weak acid should be of the order of for a satisfactory spectrophotometric end point; where this product is 10-la or less, no appreciable break appears in the titration curve (3). I n the most dilute bicarbonate solution titrated successfully in the present work, the C X , product was about 2 X Mixtures of carbon dioxide and bicarbonate may be titrated, the first end point representing the former and the second end point the sum of the two. This titration would be useful only in situations where errors of the order of 3 to 5% were tolerable. Carbonate Titrations. Figure 4 shows a spectrophotometric titration curve for the titration of sodium carbonate with hydrochloric acid. Since neither bicarbonate nor carbon dioxide solutions absorb appreciably, one sees spectrophotometrically only the first end point in the carbonate titration. I n Figure 4, the end point is extremely sharp, and the precision is correspondingly high. A series of 12 carbonate solutions ranging from about 3 x 10-3 to 4 x 10-2M shom-ed an average deviation of 0.2%. The buret readings &-ere no more precise than this, and hence the precision inherent in the spectrophotometric end point, including the graphical process, rnay be even better than indicated here. It is obvious that the carbonate titration should become much less feasible if appreciable quantities of bicarbonate are present in the carbonate solution. This effect of excess bicarbonate is shown in Figure 5, where it may be seen that the potentiometric end point (and of course the visual one too) is greatly worsened by the bicarbonate. A comparison of the spectrophotometric end points in Figures 4 and 5 shows that
Figure 5. Potentiometric and spectrophotometric titrations of 3.4 X 1 O+M sodium carbonate solution containing 1M sodium bicarbonate VOL. 34, NO. 6 , MAY 1962
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Figure 6. Spectrophotometric titration of solution containing 1.8 X 1 O-*M sodium hydroxide and 1.7 X 10-2M sodium carbonate
bicarhonate makes tlie end point lcss sharp, but to nothing like the degree to which i t affects the potentiometric one. Ten titrations of carbonate solutions in the sanie concentration range as above but which contained a thirty-fold molar
excess of bicarboxitc 4ion cti an average deviation of 0.47G. Sodium Hydroxide Plus Carbonate. The carbonate ion is too strong a base t o be feasibly differentiated froin hydroxide in a potentiometric or a visual titration. Figure 6 s h o n s a spectrophotometric titration of a sodium hydroxide-sodium carbonate mixture, where the hydroxide end point, as expectcd, is not sharp but might bc useful if high precision n.tw not required. The carbonatc cnd point i q still sharp, of course, but the, pr(&ion with n liich carbonate is drtcrmincd in the mixture depends upon both end points and accordingly is lowered to tlie same lcrcl as the hydroxide precision. The relative amounts of hydroxide and carbonatc will obviously also nffcct thc precision. Ten solutions wrre titrated containing equal concentrations of sodiuni hydroxide and sodium carbonatc ranging froin 3 x lop3to 3 X 10-*11f. Thc average deviation for thc sodiuni hydrouidc cnd point \vas 3Yc, and for the sodiuni carbonatcx crid point 3.5Tc.
1 1 ERATURE CITED
( 1 ) I ~ r r i r l i c ~ n ~ t e i rS., i , Iiolthoff, I . XI , i n I
