Fall Organic Chemistry Experiment #5
Distillation

Suggested Reading:

Digital Lab Techniques Video #15 (Distillation: Simple and Fractional)

Lehman, The Student's Lab Companion: Laboratory Techniques for Organic Chemistry

Operation #30Simple Distillation pages 182-195
Operation #32Fractional Distillationpages 208-221

Introduction to Distillation

Most of us have some idea of what distillation is all about and are familiar with Jack Daniel's whiskey and other hard liquors. In fact, the process of purifying ethanol for alcoholic beverages is one of the main applications of distillation. Another common usage is in the purification of crude oil for gasoline. Notice, by the way, that both of these processes involve the purification of mixtures of LIQUIDS. As organic chemists,distillation is the method primarily used to purify mixtures of liquids.

So what underlying physical phenomena allow us to purify liquids by distillation? Actually, you should have a pretty good understanding of the process from what you learned in general chemistry with regard to vapor pressure equilibrium. However, in order to employ this technique (distillation), you will need to understand how vapor pressures of mixtures are dependent upon the structures of the components. In addition, you will need to develop an understanding of how vapor pressure affects the behavior of the components during the distillation and how it facilitates their separation.

So, imagine that you have a flask in front of you at room temperature that contains some liquid (say about 100 mL). Now, we all know that the molecules that make up the liquid are dynamic. That is, they are in constant motion and have a certain degree of kinetic energy associated with them. Some of these molecules may contain enough internal kinetic energy so that they can actually escape from the liquid phase to the gas phase. We expect that this is occurring to a certain degree at any temperature. Over time a certain amount of the liquid molecules will escape into the gas phase and exert a pressure (vapor pressure) on the liquid below. Although vapor pressure will differ according to the identity of a compound, vapor pressure always increases as the temperature rises.

As the temperature of our system increases (say, by external heating of the flask), the percentage of liquid molecules with enough energy to escape to the gas phase will increase. In addition, the vapor pressure will also increase. Ultimately, we expect that as we raise the temperature, we will be able to increase the vapor pressure of our system to the point where it equals any external pressure (atmospheric pressure at liquid surface + hydrostatic pressure) on the system. At this point (the boiling point), the thermal energy has risen enough for molecules in the interior of the liquid to enter the gas phase and form bubbles that can rise to the surface of the liquid. The result is the observed phenomenon of boiling. It is this process of heating a liquid to the boiling point that allows mixtures of liquids to be separated from one another. The process of distillation consists of generating the vapor by boiling a liquid and subsequently condensing the "purified" vapor and collecting it in a separate container. The apparatus for this process can vary considerably depending upon the type of separation that you desire. There are basically two types of distillation set-ups -- simple and fractional. Please read the assigned operations (#30 and #32 Simple and Fractional Distillation) from the Lehman text in the Suggested Reading above. These sections will provide you with a comprehensive look at the physical principles and experimental operations involved in distillation.

A note on azeotropes

Not all liquids conform to Raoult's law and form ideal solutions. An example is the mixture of water and ethanol. Because of certain intermolecular interactions, a unique mixture (an azeotrope) of 95.5% ethanol and 4.5% water forms that boils below (78.1 ˚C) the boiling point of ethanol (78.3 ˚C). Thus, we say that ethanol/water mixtures form a low-boiling azeotrope because the boiling point of the azeotrope is lower than the boiling points of either single component. No matter how efficient is our distilling apparatus, pure 100% ethanol cannot be obtained by the distillation of a mixture of ethanol/water. You need to be careful when dealing with azeotropic systems, because a constant boiling point is NOT necessarily an indication of a single component system. In fact, azeotropes represent an example of a system that consists of two (or even three) components that boil at a constant temperature that is different than the boiling point of either pure component.

Procedure for Distillation

The purpose of this experiment is to illustrate the use of distillation to separate a mixture of two volatile liquids with different boiling points and to analyze the separation via gas chromatography. Each mixture will be issued to you and consist of a combination of two of the following liquids.

Components
cyclohexane
toluene
pentane
heptane
diethyl ether
1-butanol
Isobutyl alcohol
2-propanol
ethyl acetate
2-butanone
acetone
2-butanol

The liquids in the mixture will be separated by FRACTIONAL distillation (See Figure E23 on Page 214 of the Lehman text). The results will be analyzed in a subsequent experiment by gas chromatography-mass spectrometry (GC-MS), infrared (IR) and nuclear magnetic resonance (NMR) spectroscopy. You will also construct a graph of the distillation temperature versus the total volume of distillate collected. This graph will enable you to determine the boiling points of the two liquids and to evaluate the overall efficiency of the two methods.

Each team will acquire an unknown sample mixture of approximately 50 mL. The appropriate apparatus (MACROSCALE) should be assembled according to the Figure E23 in Lehman (page 214). Carefully notice the position of the thermometer bulb in these diagrams. It must be positioned correctly for optimum results. Also, note that we will be packing the fractional distillation column with glass beads. Add your unknown mixture to the boiling flask along with a stir bar. You’ll heat the flask over a stirrer plate with a heating mantle powered by a Variac.

Start circulating the cooling water in the condenser and adjust the heat so that INITIALLY the liquid boils rapidly. You will want to maintain a gentle boiling as you are heating, so be advised that you may have to adjust the setting on your Variac. Do not heat up too quickly otherwise you will not achieve a good separation of components. We want to keep the rate to about 10 drops of distillate per minute. You are to collect your distillate in a graduated cylinder keeping track of the volume collected over time. Be sure to record the temperature after every 2-3 mL. At some point during the distillation, you should collect two aliquotsin small vials and label them “early and late”. YOU NEED TO DETERMINE THE TIMING OF COLLECTION and you must record the temperature at the moment that you collect these aliquots. Collect approximately 2 mL of each aliquot. Once you have collected a sufficient amount of your high boiling component, you can cease distillation. DO NOT DISTILL TO DRYNESS. Once your distillation is complete, you should save your vials in the refrigerator for subsequent analyses (Exp. #7). Clean up your work area and all glassware. Dispose of your liquids in the appropriate waste bottles in the fume hood. As usual, keep a good, accurate accounting of the experimental procedure and observations in your notebook.