A Simple, Experimental Illustration of the LeChatelier Principle Isabel M. Plaza del Pino and J o s e M. Sanchez-Ruiz Departamento de Quimica Fisica, Facultad de Ciencias, Universidad de Granada, Granada-18071, Spain Many textbooks of physical chemistry discuss chemical equilibrium in terms of the LeChatelier principle, which is usually expressed as below (1,2). (1) If a system at equilibrium is subjected to a perturbation,
the equilibrium will be shifted to partially undo this perturbation.
We assume that the system is a reaction mixture and that it is perturbed by the addition of one of the reactants. Then a n experimental illustration of the principle should include determining the effect of the perturbation in two cases: when no reaction occurs in the system; when the system is allowed to reach chemical equilibrium. I n other words, the effect of the equilibrium shift (the induced chemical reaction) should be measured separately. This report describes a simple experiment that meets the above requirement. The system chosen is a n aqueous solution of pyridoxal %-phosphate, n-hexylamine, and the Schiffs base they form. I t must be noted that these substances have ionizable groups. (Their fully protonated forms are shown in Fig. 1.)Thus, they exist in solution a s mixtures of molecular species that differ in the number of t" bound protons ( 3 . 4 ) .Accordinelv. we mav e x ~ e chvdroeen ions to be released or taken u p i n t h e process. For instance, a t a DH of about 8. these three substances are present mostly a s monoprotonated species. Thus, formation of one mole ofSchiffs hase is accom~aniedhv the release ofahout one mole of protons.
-
HC=O
0
H
II
o OH I -
-
A
~
o
-
H
2
cI ~
~
~
N*
I
H
In general, the chemical reaction that occurs in the system may be represented a s hexylamine + ppidoxal phosphate
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Journal of Chemical Education
Schiffs base + HzO + nHHi
where n~ stands for the number of moles of protons released for every mole of Schiffs base formed. Due to the pK values for the substances involved in the reaction (3, 4), Schiffs base formation is accompanied by proton release. except a t very basic pH. In otber woydi, nH > 0, and hydrogen ions must be considered as "products" in the reaction. Experimental The experimental solution was prepared by mixing equal volumes of 0.01 M pyridoxal phosphate (Merck) and 0.01 M nhexylamine (Jansen). These molar solutions were prepared by weight. Immediately after mixing, the solution was brought to a pH of about 8 with a small volume of 1M KOH. The exact values of the concentrations and the initial pH are not critical for the result of the experiment. Then one half-hour after mixing (to ensure that chemicalequilibriumwasestablished),0.3 mLof0.2M HClO was added to 40 mL ofthe solution . The time-denendence of the pH valur wns Followed using a con\mtmmI electrudepHmeter system. The tme-dependenceof the absorbancr at 4% n m was follou,ed using a Car!-210 wpectrophotomrrer t l a n pathlength).
.
~~~~~~
Results and Discussion For a n aqueous solution of pyridoxal phosphate, n-hexylamine, and Schiffs base, we assume that chemical equilibrium is established. We now perturb the system by adding ~ a small volume of HC1 solution. The main effect of this perturbation is to increase the hydrogen ion concentration. . . kccordmg to the LeChatelier principle a s formulated above, the chemical equilihrium will shift to reducc [he concentration of hydrogen ions. However, the rates of formation and hydrolysis of the Schiff's base formed by pyridoxal 5'-phosphate and nhexylamine are comparatively low (5).Therefore, the effect of the induced equilibrium shift can be monitored by following the time dependence of the pH value. Figure 2a shows a n example of this kind of experiment. Adding HCI immediately decreases the pH, while the concentrations of aldehyde. amine. and Schiff's base remain constant. Then the p~ increase$ over t i e a s the chemical reaction responds to acid addition as predicted by the LeChatelier principle. This result is rationalized below. If we assume that nH is constant within the pH interval of the experiment, the equilibrium between aldehyde, amine, Schiff's base, and hvdroaen ions can be described bv the equilibrium constant below.
K=
Figure 1. Fully protonated forms of pyridoxal 5'-phosphate 1, r? hexylamine 2, and the Schiff's base 3 formed between these two substances.
2
[Schiffs base] [aldehyde][aminel
where ax = 10-pH. The concentration of water, which is ~racticallvconstant, is included in the value of& As usual \r hen deajing with complex equilihna in solution, the equilibrium constant K is defined in terms of the conventional activity for the hydrogen ion (aH)and by the total concen-
trations of all other substances. I t must be noted that, neelectine deviations from ideal behavior. K is a true const& t h a t does not depend on pH. The relative amounts of aldehyde. amine, and Schiffs base present at equihbrium at a &en pH are determined by an observed equilitriumconstant 13, 61. [Schiffs base1 Kobsd = [aldehyde][aminel The two equilibrium constants are related by eq 3 or 4. lag K = log KObsd + n ~ l o aH g
(3)
logKobsd= logK+ n ~ p H
(4)
Clearly, Kobd does depend on pH. This dependency is given by the following equation
which can be easilv obtained bv differentiation from ea 4. The above derivati& assumes that nHis constant, but eq 5 holds true even when nH depends on pH (6). Since hydrogen ions are "products", nH > 0.The initial pH decrease. which was causedbv the wrturbation, lowers the value of K , (eq ~ 5). The chemical equilibrium shifts to the left, thus producing Schiff's base hydrolysis, proton uptake from the medium, and pH increase. (See Fig. 2a.) The hydrolysis of the Schiffs base can be monitored by following the time dependence of the absorbance of the solution a t 465 mn. (See Fig. 2b.) Pyridoxal phosphate and the Schiffs base both show absorption bands in the W-vis region of the spectrum (31, these being the bands of the Schiffs base displaced to higher wavelengths. Only the
EFFECT OF THE E@UlLlBRlUM SHIFT
Schiffs base absorbs sienificantlv a t 465 mn. Thus. the time-dependent decreas'in absordance a t 465 mn observed aRer HCI addition (Fie. 2b) clearlv shows that the eauilibrium does shift to theleft ( ~ c h i f fbase s hydrolysis). I t must be noted that the results of the experiment shown in Figure 2a (the pH-time profile) would have been the same qualitatively if n" < 0,that is, if hydrogen ions were "readants". Then the initial pH decrease would increase the value of Kobd (eq 51, and the equilibrium would shiR to the right (Schiff's base formation), again producing proton uptake from the medium with a pH increase. I t is clear that the results shown in Figure 2 are consistent with the LeChatelier principle as formulated above (statement 1). The work of several researchers (1,2) has shown, however, that such usual formulations of the principle are vague and could lead to wrong predictions. Amore rigorous treatment of the principle, based on the second law, has been summarized by J. de Heer (1,2).The results of this treatment are expressed below. (2a) The change in an intensive variable caused by changing the correspondingextensivevariable is smaller ifchemical eauilibrium is maintained than if no reaction could occur in the svstem. 12h, The change in a n cxrrnsive vanable caused by changing thr rorrespmdmg intensirr vnriahle is larger ifrhrmirnl equlhbnum is maintain~dthan if no reartion could occur in the system. In the experiment of Figure 2a, HCI addition increases thc number of moles of hydrogen ion in the system N H I . which is the extensive variablecorresponding to the chemical potential of the hydrogen ion (pH. Increasing N H (the oerturbat~onlchanges tht, corres~ondinaintensive vanible and increases 'the chemical potential (pH).According to statement 2a. the change in un is smaller if chemical equilibrium is maintainedxhan if no reaction could occur in the system. In other words, the equilibrium shiR (the induced chemical reaction) must lower p ~ This . is wnsistent with the results shown in Figure 2a. Thus, the pH increase caused hy the equilibrium shift lowers the chemical potential of the hydrogen ion run 1, given that and pH pH = F;
+ RTln a, = & - RT(In10)pH
(6)
where p;I and a~ are the standard chemical potential and activity of the hydrogen ion, respectively, and pH = -logax. We use single-ion chemical potentialsand activities for the hydrogen ion. We believe that this is permissible, a t least for illustration, because operational pH values are defined in terms of a conventional single-ion activity scale (7). We have reported a simple experiment that illustrates the LeChatelier principle that we believe has many merits.
SHIFF'S HYDROLYSIS
I
I
I
I
0
8
16
TIME AFTER HCL ~001~10H/rnin F~gLre2 a. Tne tome aependence of me pH after aao t on of HC to an aqdeods sotdtlon of pyrlooxal pnospnare, n-hexylamne, ana the Scn If s oase. b. The t me dependence of tne absoroance at 465 nm after addition of HCI to the solution
1.The effectof the induced equilibrium shift is measured separately. 2. The experiment uses common laboratory equipment and chemicals that are inexpensive and commercially available. 3. The results of the experiment can be interpreted at various levels of sophistication. They can be explained on more than one basis: the usual formulations ofthe LeChatelier principle (such as statement 1); a pH-dependent, observed equilibrium constant (eq 5);the more rigorous treatment of the principle (statement2a) summarized by de Heer ( I , 2).
Literature Cited 1, de H e m J. J Chem. Educ 1#57,34,376. 2 , de Hem, J. P k n a m ~ m i o g i i lThermodynamim with Applications lo Chemiafv; Rentice: Englewood CEffs, NJ, 1986:Chapter 20. 3. Metzler, C. M.: CahiU, A,; Metzleler, D. E. J. Am. Ckem. Soc. 1880,102.60T5. 4. Cortijo,M.; Llor, J.; Sanchez-Ruiz. J.M. J. B i d . Chem, 1988,263,17960. 5. Sanchez-Ruiz, J.M.;Ra~8y8~z-PuBdo,J.M.:Uor, J.;Colhjo,M.J.Ckm.Sa.Psrkin Ram. 111882,1425. 6 . Alberty, R.A. J Am. Chrm. Sa.1988.91.3899. 7. Batea. R. G. Debrminotion of oH.Thoon a d PmefieP. 2nd ed.: Wilev: New Vork, 1973: Chapter 2
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Volume 68 Number 11 November 1991
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