Why distill under reduced pressure

Boiling stones cannot be used with vacuum distillation as air trapped in the stone's pores is rapidly removed under vacuum, causing the stones to fail to produce bubbles. Although greasing is somewhat of a personal choice with simple and fractional distillations, all joints must be greased in vacuum distillations or the system will leak and fail to achieve a low pressure Figures 5. Begin assembly of the apparatus near the vacuum source. If using a water aspirator, test to be sure that the aspirator works well as some are more functional than others.

To test an aspirator, apply thick vacuum hosing to the nub on the aspirator, turn on the water and feel for suction at the end of the hose with your finger Figure 5.

A Claisen adapter should be included in the apparatus as solutions under vacuum tend to bump violently a Claisen adapter is labeled in Figure 5. Attach thick-walled tubing to the vacuum adapter on the distillation apparatus Figure 5. A trap suitable for a water aspirator is shown in Figure 5. Connect the trap to the vacuum source aspirator or vacuum pump.

It is best to not bend or strain the tubing as much as is practical, as this may create a leak in the system. Insert a wood block Figure 5. Begin the Distillation Before heating , turn on the vacuum source to begin reducing pressure inside the apparatus.

There should not be a hissing sound or else there is a leak in the system. The purpose of reducing the pressure before heating is for removal of very low-boiling liquids e. If the system were heated at the same time, the low-boiling liquids might boil violently in the flask.

If a manometer is available, take note of the pressure inside the apparatus. This may be used to predict the boiling point of the sample. When confident that the apparatus is adequately evacuated and any low-boiling compounds have been removed, begin heating the sample Figure 5. If it is difficult to achieve more than a reflux, the Claisen and three-way adapter can be insulated by wrapping them tightly with glass wool then aluminum foil Figure 5. Insulation allows the column to maintain heat and the sample to remain in the gas phase longer.

A small gap should be left in the insulation near the distilling flask to "peek in" and make sure the stirring mechanism continues to work properly. Record the temperature over which material is collected, making sure the value corresponds to a temperature when the thermometer bulb is fully immersed in vapors.

If a manometer is used, also record the pressure. If no manometer is used, record the vacuum source e. Pure liquids do not always distill at a constant temperature when under vacuum, as variations in pressure so easily occur and affect the boiling temperature.

This is especially true when the vacuum source is a water aspirator, where variations in water flow alter the pressure. If more than one fraction of distillate is desired, the distillation must be stopped before changing the receiving flask see the next section for how. If available, a "cow" receiving flask can be used to collect different fractions without ceasing the vacuum Figure 5. Stop the Distillation To stop the distillation, first remove the heat source, cool the flask to room temperature then further cool in a tap water bath Figure 5.

Slowly reinstate the atmospheric pressure into the flask by opening the pinch clamp at the vacuum trap Figure 5. You will know the system is open to the atmosphere when there is an increase in water flow at the aspirator, or if a hissing sound is heard. Then turn off the vacuum source. Let's begin by discussing the vapor pressure of a pure substance and how it varies with temperature. Vapor pressure is an equilibrium property. If we return to that hot windy day at the beach and consider the relative humidity in the air, the cooling effect of the wind would be most effective if the relative humidity was low.

Everyone in St. Louis has experienced how long it takes to dry off on a hot humid day. At equilibrium, the process of vaporization is compensated by an equal amount of condensation. Incidentally, if vaporization is an endothermic process i.

Now consider how vapor pressure varies with temperature. Figure 1 illustrates that vapor pressure is a very sensitive function of temperature.

It does not increase linearly but in fact increases exponentially with temperature. If we follow the temperature dependence of vapor pressure for a substance like water left out in an open container, we would find that the equilibrium vapor pressure of water would increase until it reached 1 atmosphere or Pa At this temperature and pressure, the water would begin to boil and would continue to do so.

It is not possible to achieve a vapor pressure greater than 1 atmosphere in a container left open to the atmosphere. Of course, if we put a lid on the container, the vapor pressure of water or any other substance for that matter would continue to. Elevation of the boiling point with increase in external pressure is the principle behind the use of a pressure cooker. Elevation of the boiling point with an increase in external pressure, while important in cooking and sterilizing food or utensils, is less important in distillation.

However, it illustrates an important principle that is used in the distillation of many materials. If the boiling point of water is increased when the external pressure is increased, then decreasing the external pressure should decrease the boiling point. While this is not particularly important for the purification of water, this principle is used in the process of freeze drying, an important commercial process.

In addition, many compounds cannot be distilled at atmospheric pressure because their boiling points are so high. At their normal boiling points, the compounds decompose. Some of these materials can be distilled under reduced pressure however, because the required temperature to boil the substance can be lowered significantly. A nomograph is a useful device that can be used to estimate the boiling point of a liquid under reduced pressure under any conditions provide either the normal boiling point or the boiling.

Figure 2. A nomograph used to estimate boiling points at reduced pressures. To use, place a straight edge on two of the three known properties and read out the third. Column c is in mm of mercury. An atmosphere is also equivalent to To use the nomograph given the normal boiling point, simply place a straight edge at on the temperature in the central column of the nomograph b. Rotating the straight edge about this temperature will afford the expected boiling point for any number of external pressures.

Simply read the temperature and the corresponding pressure from where the straight edge intersects the first and third columns. Using the nomograph in Figure 2 and this temperature for reference, rotating the straight edge about this temperature will afford a continuous range of expected boiling points and the required external pressures necessary to achieve the desired boiling point.

Although all of us have brought water to a boil many times, some of us may have not realized that the temperature of pure boiling water does not change as it distills.

This is why vigorous boiling does not cook food any faster than a slow gentle boil. The observation that the boiling point of a pure material does not change during the course of distillation is an important property of a pure material. The boiling point and boiling point range have been used as criteria in confirming both the identity and purity of a substance.

Of course, additional criteria must also be satisfied before the identity and purity of the liquid are known with certainty. You will use both of these properties later in the semester to identity an unknown liquid. Occasionally, mixtures of liquids called azeotropes can be encountered that mimic the boiling behavior of pure liquids.

These mixtures when present at specific concentrations usually distill at a constant boiling temperature and can not be separated by distillation. The azeotropic composition sometimes boils lower the than boiling point of its components and sometimes higher.

Mixtures of these substances at compositions other than those given above behave as mixtures. Returning to our discussion of boiling water, if we were making a syrup by the addition of sugar to boiling water, we would find that the boiling point of the syrup would increase as the syrup begins to thicken and the sugar concentration becomes significant. Unlike pure materials, the boiling point of an impure liquid will change and this change is a reflection of the change in the composition of the liquid.

In fact it is this dependence of boiling point on composition that forms the basis of using distillation for purifying liquids. We will begin our discussion of distillation by introducing Raoult's Law, which treats liquids in a simple and ideal, but extremely useful manner. Figure 3. The apparatus used in a simple distillation.

Note the position of the thermometer bulb in the distillation head and the arrangement of the flow of the cooling water. This relationship as defined is capable of describing the boiling point behavior of compound A in a mixture of compounds under a variety of different circumstances. Although this equation treats mixtures of compounds in an oversimplified fashion and is not applicable to azeotropic compositions, it does give a good representation of the behavior of many mixtures.

Let's first consider a binary system 2 components in which only one of the two components is appreciably volatile. Raoult's law states that the observed vapor pressure of water is simply equal to the product of the mole fraction of the water present and the vapor pressure of pure water at the temperature of the mixture. Once the sugar-water mixture starts to boil, and continues to boil, we know that the observed vapor pressure of the water must equal one atmosphere.

Water is the only component that is distilling. Since the mole fraction of water in a mixture of sugar-water must be less than 1, in order for the observed vapor pressure of water to equal one atmosphere, must be greater than one atmosphere.

As the distillation of water continues, the mole fraction of the water continues to decrease thereby causing the temperature of the mixture to increase. Remember, heat is constantly being added. If at some point the composition of the solution becomes saturated with regards to sugar and the sugar begins to crystallize out of solution, the composition of the solution will become constant; removal of any additional water will simply result in the deposit of more sugar crystals.

During the course of the distillation, the water vapor which distilled was initially at the temperature of the solution. Suspending a thermometer above this solution will record the temperature of the escaping vapor.

Cooling below this temperature will cause most of the vapor to condense to a liquid. This is why the distillate is frequently chilled in an ice bath during the distillation. The distillation of a volatile material from non-volatile is one practical use of distillation which works very well.

However, often there may be other components present that although they may differ in relative volatility, are nevertheless volatile themselves. Let's now consider the two component system you will be using in the distillations you will perform in the laboratory, cyclohexane and methylcyclohexane. The vapor pressures of these two materials in pure form are given in Table 1. As you can see from this table, although cyclohexane is more volatile than methylcyclohexane, the difference in volatility between the two at a given temperature is not very great.

This means that both materials will contribute substantially to the total vapor pressure exhibited by the solution if the distillation is carried out at 1 atmosphere. The total pressure, P T , exerted by the solution against the atmosphere according to Dalton's Law of partial pressures, equation 2, is simply the sum of the observed vapor pressures of cyclohexane, , and methylcyclohexane, :. As before, boiling will occur when the total pressure, P T , equals an atmosphere.

However since we have two components contributing to the total pressure, we need to determine the relative contributions of each. Again we can use Raoult's Law but we need more information about the system before we can do so. In particular we need to know the composition of the solution of cyclohexane and methylcyclohexane. For ease of calculation, let's assume that our original solution has equal molar amounts of the two components. What we would like to determine is whether it would be possible to separate cyclohexane from methylcyclohexane by distillation.

By separation, we would like to determine if it would be possible to end up with two receiver flasks at the end of the experiment that would contain mainly cyclohexane in one and mainly methylcyclohexane in the other. It is clear that at some point we will need to intervene in this.

Table 1. Vapor pressures of cyclohexane and methyl cyclohexane as a function of temperature. Otherwise, if we were to collect the entire contents of the original distilling flask, called the pot, into one receiver flask, we would end up with the same composition as we started.

Initially the mole fractions of both cyclohexane and methylcyclohexane are 0. From Raoult's Law equation 1 , Dalton's Law equation 2 and the information in Table 1, we can estimate that boiling will occur at approximately K when the total pressure of the two components equals one atmosphere or The first thing that we should note is that the initial boiling point is higher than the lowest boiling component and lower than the highest boiling component.

Next, we should inquire about the composition of the vapor. Is the composition of the vapor the same as the initial composition of the pot or is it enriched in the more volatile component? If the composition of the vapor is the same as that of the original mixture, then distillation will not be successful in separating the two components. However, we should ask, "What is the composition of the vapor? First we note that:. If the total vapor can be treated as an ideal gas, then according to Dalton's Law, so can each of the components.

Since the two components are in thermal contact and are distilling together, we can expect them to be at the same temperature. We don't necessarily know the volume of the container, but since it is assumed that the volumes of the molecules are very small in comparison to the total volume the gas occupies, whatever the value of V, it is the same for both components. This means we can establish the following ratio:.

If we use the experimental values found in Table 1, we conclude that the composition of the vapor is 1. This simple treatment allows us to understand the principles behind distillation. However it is important to point out that distillation is far more complex than our simple calculation indicates. For example, we just calculated the composition of the vapor as soon as the solution begins to boil and we have correctly determined that the vapor will be enriched in the more volatile component.