An azeotropic mixture is one whose boiling point is lower than that of any other combination of the compounds in the mixture. This is caused by the attraction between the molecules of the mixture, and is not predicted by the ideal gas laws. When an azeotropic mixture vaporizes, the vapor has exactly the same concentrations as the liquid, which is why you cannot form a more concentrated mixture by distilling.

On the left, we've attempted to combine two graphs into one to explain the phenomenon.

Fig. 8-7

The first graph plots the mol fraction Y of ethanol vapor you get from a mixture with mol fraction X of ethanol. This is the curved line running from bottom left to top right. Where it crosses the straight line joining the two ends, the mol fraction of ethanol in the vapor is the same as the mol fraction of ethanol in the liquid. Below that point the vapor is always richer in ethanol than the liquid it came from, and above that point the vapor is always leaner in ethanol than the liquid it came from. This is the azeotropic point.

The second graph plots the values of X and Y for all the temperatures from 78°C to 100°C. It is the pair of curved lines running down from the top left. These lines meet again before the temperature reaches 78°C, then run up and to the right. They finally meet again when X = 1. The lowest point of this graph matches the crossover point of the first graph, when the mol fraction of ethanol in the liquid equals the mol fraction in the vapor. These are the Azeotropic points and XA = YA. This shows why you can't get a higher percentage if ethanol than around 96% by distillation alone - at atmospheric pressure. Significantly lowering the pressure in the still (vacuum distillation) will produce a higher strength product, but is not worth the cost and danger.

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