Kenneth C. M i e and R. S. Macomber University of Cincinnati Cincinnati, Ohio 45221 ~
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I
Recrystallization:
Unexpected Behavior
An undergraduate organic demonstration or experiment
The theory and application of recrystallization is so fundamental to the organic chemist that i t is virtually impossible to find a lecture or laboratory text that does not describe the technique in some detail. In laboratory manuals the discussion is usually preceded by the construction of normal phase diagrams of the type shown in Figure 1. Such diagrams are then used to show how, as the temperature of a solution is lowered, the solute becomes less soluble and ultimately begins to crystallize, hopefully at a rate different from that of any co-solutes. Solid Solvent =F== Solute Solvent. (1) The thermodynamic interpretation of such a process can he approached as follows. The system will, at any temperature, prefer the side of eqn. (1) with the lower free energy (AG = -RT 1nK). At constant pressure, AG for eqn. (1) is governed by d(AG) = -ASdT &dni (2) Focusing our attention on the limited region near the 01, eqn. (2) point of complete dissolution (where An, becomes (3) d(AG)/dT = -AS For the normal behavior described above, d(AG)/dT must be negative because the right side of eqn. (1) becomes more favored as the temperature increases. This can only be the case if A S is positive, a result which is intuitively pleasing since the disorder of the solute molecules (a measure of the difficulty of finding one) should certainly increase upon dissolution. Unfortunately, a simplified description such as the one above often leaves the student with the impression that the solubility of any substance in any solvent will increase with temperature. This is clearly not the case, but such inverse behavior is relatively rare among organic compounds in organic solvents and is usually unspectacular. We offer here a simple and inexpensive method for providing a demonstration of an exception to the "rule" that solubility increases with temperature. During the course of other work, we prepared pyridine hydrobromide (I).l Although i t has probably been pre-
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liquid solution
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pared countless times, we have located hut one reference to its properties.2 Compound I is interesting in several respects: it sublimes readily, melts a t a reasonably low temperature, and is volatile enough to be analyzed by mass 1 Spectral (pmr, ir, and ms) and analytical data confirmed its structure. =The "Dictionary of Organic Compounds" (4th ed., Oxford University Press, 1965) lists under pyridine that its betaine hydrobromide exists as reddish-yellow plates, mp 213°C.
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Figure 1. Partial phase diagram for normal solubility behavior.
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Figure 2. Partial phase diagram for inverted solubility behavior.
spectrometry, properties not normally associated with salts. But its most interesting property is that cool solutions of I in chloroform redeposit the crystalline salt upon warming, and redissolve the salt upon cooling! We have now studied this peculiar behavior, and it provides an excellent demonstration of solubility decreasing with increasing temperature. All one need do to recrystallize I is to prepare a cold solution in chloroform, then warm the solution up. The absolute solubility of I in chloroform is 18% (w/w) at O0C, hut this drops to 10% (w/w) at 60°C. The effect is spectacular because floculent crystals nearly fill the vessel. The warming, then cooling may be repeated a t will as often as desired. We have cycled the process as many as 31 times with no sign of deterioration other than a slight yellowing of the solution. The thermodynamic "explanation" of this unusual behavior is that since the left side of eqn. (1) now becomes more favored with increasing temperature, A S for the dissolving process must necessarily he negative. This would indicate that the entire system is becoming more ordered upon dissolution, even though the entropy of I itself is inVolume 50, Number 8, August 1973
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creasing. Clearly the determinant here, which was overlooked in the discussion of normal hehavior, is the importance of solvent-solvent and solvent-solute a s well a s solute-solute interactions. Apparently i n t h e I-chloroform system, the solvated molecules of I possess lower total entropy (disorder) t h a n t h a t of crystalline I plus self-associated chloroform. T h e net result is a phase diagram (Fig. 2) which is inverted by comparison to t h e one in Figure 1.
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Experimental Section
Preparation of Pyridine Hydrobromide (1). A h e c k 50-ml flask equipped with gas inlet and outlet tubes and a stirring bar was charged with 15 ml dry pyridine. Hydrogen bromide gas from a lecture bottle was gently passed into the pyridine; care must be taken to avoid clogging the inlet tube with product. The solid product, mobilized by unreacted pyridine, was isolated by filtration, washed with pentane, and sublimed (120°C. 0.12 mm). The hygroscopic colorless crystals of I melted at 219.5-220"C.2
