Peter Jones, M a x L. Haggen and Jethro L. Longridge The University Newcastle upon Tyne, England
The Hydration of Carbon Dioxide A double dock experiment
The uncataly~edrates of hydration of carbon dioxide and dehydration of carbonic acid are sufficiently slow as to be limiting factors in the overall process:
=
+HI
=
+I1+
=
-Ha0
COszPCOZ- HCOs Con (soln) -H -H+ +Ha
= Con (gas)
In their detailed discussion of the biological implications of this aspect of the physical chemistry of carbon dioxide, Edsall and Wyman (I) describe a simple demonstration of the time requirement of the hydration reaction. This experiment, which uses a "clock" technique and is suitably called the "Soda Water Clock," is, in our opinion, well worth inclusion in a repertoire of "clock" demonstrations. A suitable recipe is described below. By using a "Double Clock" extension of this experiment we have developed a version which provides a quantitative kinetic experiment suitable for undergraduate laboratories. I n this experiment the rate constant for the hydration reaction, COz HzO -4 HzCOa, is determined a t about 0°C. The material and apparatus used are simple and inexpensive; the essential requirement for successfully performing the experiment is accurate volumetric work. The theory of the method and some experimental results are given below. The experimental method is similar to that used 50 years ago by Thiel (9). Kern (5) has recently published an excellent review of both early and recent work on this reaction and has pointed out that, although Thiel's method was sound in principle, his interpretation of the reaction was incorrect.
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The Soda Water Clock: A Demonstration
For demonstration purposes carbon dioxide solutions are conveniently prepared using "Sparklets" syphons. A Sparklets Globemaster syphon, when charged with one cartridge of carbon dioxide, yields a solution of 0.1 M COI.
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Procedure. Cool the syphon of soda water and a bottle containing distilled water and bromothymol blue indicator to 0-2'C. Charge a 2-ml hypodermic syringe with 2 ml of 0.4 M sodium hvdroxide solution. Take about 25 ml cooled distilled water and indicator in a. 1W ml conical flask. Add about 25 ml cooled soda water down the side of the flask. Swirl the contents of the flask and inject the sodium hydroxide solution.
The solution, which was originally yellow in color, immediately becomes blue. After about 30 sec the solution becomes yellow again. The explanation is 610
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as follows. The soda water contains mainly unhydrated COz and a little HzCOa. The HnCOapresent reacts very rapidly with OH- to give HC03- and C032-. Before further reaction can occur the COz must be hydrated. Since the hydration reaction is slow, the pH remains high and the solution blue for a considerable period. A follvw-up ezperimenl. Repeat the experiment described above hut add a few drops of fresh blood to the flask at the third stage. Now the blue color produced on injection of the alkali is discharged immediately.
The slowness of the dehydration of carbonic acid provides the biological necessity for the enzyme carbonic anhydrase which is found in the blood. This catalyses the dehydration of H2C03 (and inevitably the hydration of COz)most efficiently. The Double Clock Experiment
A feature which complicates interpretation of the soda water clock experiment is the occurrence, simultaneously with the hydration reaction, of the process: COS
+ OH--
HCOJ-
To avoid serious interference from this competing process it is necessary to work a t pH < 9. This might appear to present serious difficulties to the development of a quantitative "clock" experiment. However, the simplicity of the "clock" technique may he retained in the following way. The experiments are carried out as before; i.e., initially enough alkali is added to increase the pH of the reaction solution to 12-13, but two indicators are now employed. Only one indicator is added initially and this is chosen to give an end point a t pH 9, thus marking the start of-the pH range in which the hydration reaction predominates. The stop watch is started a t this point, (t = t,). A small volume of a second indicator solution is immediately added. This will give a color change a t a time (t = tz) which corresponds to a decrease in the pH of the reaction solution by about one more unit. Phenolphthalein and bromothymol blue are suitable indicators. An alternative is to use phenolphthalein only but to increase its concentration about tenfold after the first color change. Since the observed endpoint corresponds to a fixed color density (i.e., in the same flask with the same volume of solution, a fixed concentration of alkaline form of phenolphthalein,) whereas the pH depends on the color ratio, the pH of the endpoint will depend on the total indicator concentration. I t is, of course, necessary to keep the total indicator concentration
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very small a t all times. The pH's a t which the endpoints are observed are determined in blank experiments. In our laboratory students are provided with a pH (glass electrode) curve for the titration of 50 ml of 0.1 M tris-base (NH2C(CH20H),) with HC1 a t 0°C. By repeating this titration using indicators a t the same concentrations as in the kinetic experiments the required pH's are obtained by interpolation from the calibration curve. The theory of the method is as follows. The reaction is carried out under conditions such that the hydration reaction predominates, so that effectively:
If Co and [OH-],, are the initial carbon dioxide and hydroxide ion concentrations and CI, [OH-]I; C2, [OH-], are the corresponding concentrations a t the first (t = t3 and second (t = t2)color changes, then
The experiment relies on the occurrence of a perceptible carbonate-bicarbonate buffer action a t pH 5 9. The buffer action is allowed for as follows: Conservation of carbon. CI
=
Co - [HCOI-1, - [COaZ-11
(3)
where [HC03-]1 and [COsZ-I1are the bicarbonate and carbonate ion concentrations a t t = tl. Electroneutrality of the solution (ignoring the indicator). Initially
[Na+] is constant, so that when [H+] is sufficiently small we have [HCOI-1,
+ 2[COs2-1, = [OH-lo
The pH of tk sohiion ad t
- [OH-],
(6)
= t, is given by
. IHCOa-Ir [OH-], = K -K, [CO,l-l,
From equations (3), (6), and (7) remembering that in these experiments [OH-lo >> [OH-Il, we obtain
(phenolphthalein) method. For greater accuracy in these experiments the C02 solutions were prepared in the reaction flask by neutralization of sodium bicarbonate with hydrochloric acid. The experimental procedure was as follows: (1) The stock solutions were sodium bicarbonate (about 0.05 M),standard hydrochloric acid (ahout 0.2 M), and sodium hydroxide (a number of solutions in the range 0.05-0.4 M). The final reaction solution compositions were hydrochloric acid to neutralization Sodium bicarbonate 2 ml sodium hydroxide solution (varioua concentmtions) 0.5 ml firat indioator solution (standard phenolphthalein solution diluted tenfold with distilled water). distilled wster to total volume of 50 ml. (2) Titrate the sodium bicarbonatesolutionwiththeHCI either to the methyl orange endpoint or with a pH meter to pH = 4.0. Standardize the sodium hydroxide solutions using the eame standtlrd acid. Success in these experiments depends on careful volumetric work. (3) Put in reaction flask sodium biwbonate solution, distilled water and first indicator. Cool (this is to avoid loss of CO1 on neutralization). Add the calculated volume of HCl slowlv with gentle swirling. Cool in 0-2'C thermostat or ice b a t h h reactFon temperature. (4) Cool NaOH solution and second indicator solution. Charge cooled syringes with these solutions. (5) Swirl the reaction flmk contents (a magnetic stirrer is useful but violent shaking will result in loss of COP)and inject the alkali. (6) Note the first color change and etart the stop watch (1 = t,). (If a split second stop watch is available it is not inconven~entto start timing a t (5)). Swirl the reaction solution and immediately inject the second indicator solution. (7) Stop the watch at the seeond color change ( l = k).
+
In these experiments the first color change occurred a t pH = 9.0, the second a t pH = 8.2. The appropriate equilibrium data are pK, = 14.94; ~ K =I 10.55. In Figure 1 (upper curve) the data of Table 1 are compared with equation (9). The rate constant obtained for the slope of the graph was 1.9 X see-'. This compares with the literature value (1) of 2.0 X lo-= sec-I a t 0' C. Figure 1 (lower curve) shows the data obtained in another set of experiments where the two indioator method (phenolphthalein and bromthymol blue) was used and the carbon dioxide solutions were prepared from a concentrated stock solution obtained using a soda water syphon. The results show much more scatter (due to loss of C02 during transfer from the syphon) and the rate constant is somewhat smaller, k = 1.2 X 10-asset-'at 2.0 0.5'C. Experiments may be extended to demonstrate catalysis of the reaction. The hydration of COe is catalysed by bases in general. Although the "Double Clock" method is not suitable for bases with appreciable
*
Table 1.
The corresponding expression for C2 is obtained similarly. Since [OH-], and [OH-I2 are constants the expression in brackets in equation (6) is a constant (BI). The corresponding factor in the equation for C2 we may write as Bz. Provided that pH < 8 a t t = tz it is sufficiently accurate to take B2 = 1.00. Hence from equation (1)
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+ +
Typical Dato for Hydration of Carbon Dioxide at 2'C
Cn1M X loS)
[OH-1, 1M X 10Si
t, (sec)
Table 1 presents the data obtained in a set of experiments a t 2.0 + 0.5'C using the dngle indicator Volume 41, Number 1 1 , November 1964
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buffering action in the pH range 7-9 (e.g., arsenious acid, pII: = 9.45), the catalytic action of a sufficiently weak acid such as formaldehyde hydrate (4) (PIC. = 13.9) is readily shown. A marked reduction in (h - 1,) is observed with formaldehyde concentrations as low as 0.03 M. The rate equation is now
where [F-] is the anion CH,(OH)O-. The experiments are not readily interpreted quantitatively since [F-] depends upon pH according to
+
IF-] = 1F01/(1 IHtl/K,) where [Fa]is the total formaldehyde concentration.
techniques. "Clock" experiments have long been used in lecture demonstrations for their dramatic and aesthetic appeal. They deserve further recognition in the undergraduate laboratory for their precision and simplicity. The latter aspect is particularly inlportant if it is considered ( 6 ) that the best place in the undergraduate course for the teaching of kinetics is the freshman year. Here the large numbers demand sin~ple,robust apparatus if costs are not to be prohibitive. In our own first year class the special scientific interests of the students cover the range from biological to applied science. The hydration of carbon dioxide and its reverse process are reactions of ubiquitous inherent interest. We have noted the physiological importance of the dehydration of carbonic acid; for students of applied science one may note the utility of the Giammarco-Vetrocoke process (7) in which catalysis of the hydration reaction by As103 is used to facilitate the stripping of COI froln natural gas and from ammonia syuthesis gas. Acknowledgment
We wish to thank Sir Owen Wansborough-Jones aud British Oxygen Chemicals Ltd., for gifts of apparatus and materials, Dr. I
