This article incorporates, in modified form, material from Illustrated Guide to Home Chemistry Experiments: All Lab, No Lecture.
Distillation is the oldest method used for separating mixtures of liquids. Distillation exploits the fact that different liquids have different boiling points. When a mixture of liquids is heated, the liquid with the lower (or lowest) boiling point vaporizes first. That vapor is routed through a condenser, which cools the vapor and causes it to condense as a liquid; the liquid is then collected in a receiving vessel. As the original liquid mixture continues being heated, eventually, some or all of the lower-boiling liquid is driven off, leaving only the higher-boiling liquid or liquids in the distillation vessel.
I say “some or all” because distillation is an imperfect method for separating mixtures of liquids that form azeotropes. An azeotrope, also called a constant boiling mixture, is a mixture of two or more liquids at a specific ratio, whose composition cannot be altered by simple distillation. Every azeotrope has a characteristic boiling point, which may be lower (a positive azeotrope or minimum-boiling mixture) or higher (a negative azeotrope or maximum-boiling mixture) than the boiling points of the individual liquids that make up the azeotrope.
For example, ethanol forms a positive azeotrope with water. The boiling point of a mixture of 95.6% ethanol (by weight) with 4.4% water is 78.1 °C, which is lower than the boiling point of pure water (100 °C) or pure ethanol (78.4 °C). Because the azeotropic mixture boils at a lower temperature, it’s impossible to use simple distillation to produce ethanol at concentrations higher than 95.6%. (More concentrated ethanol solutions can be produced by using drying agents such as anhydrous calcium chloride that physically absorb the water from a 95.6% solution of ethanol. These solutions must be stored and handled carefully, because otherwise they absorb water vapor from the air until they reach the 95.6% azeotropic concentration.)
Ethanol also forms azeotropes with many other liquids, including some that are poisonous or taste bad. This allows production of denatured ethanol, which is toxic, cannot be drunk and so can be sold cheaply without cannibalizing sales of (and taxes on) much more expensive potable ethanol, such as vodka and other distilled beverages.
High Boiling Point Azeotropes
Hydrochloric acid is one familiar example of a negative (high-boiling) azeotrope. Pure hydrogen chloride has a boiling point of -84 °C and pure water a boiling point of 100 °C. A solution of 20.2% hydrogen chloride (by weight) in water has a boiling point of 110 °C, higher than the boiling point of either component. This means that boiling a solution of hydrochloric acid of any concentration eventually produces a solution of exactly 20.2% hydrogen chloride by weight. If the starting solution is more dilute, water is driven off until the solution reaches 20.2% concentration. If the starting solution is more concentrated, hydrogen chloride gas is driven off until the solution reaches 20.2% concentration.
In this laboratory, we’ll use distillation to increase the concentration of an ethanol solution. At 25 °C, pure water has a density of 0.99704 g/mL and pure ethanol a density of 0.78522 g/mL. Solutions of ethanol and water have densities between these figures. If you add the densities of the pure liquids and divide by two, you get 0.89113 g/mL, which you might assume is the density of a 50/50 ethanol-water mixture. As it turns out, that’s not true. Ethanol and water do not mix volumetrically; that is, if you mix 100 mL of pure ethanol with 100 mL of pure water, you do not get 200 mL of solution, for the same reason that dissolving 100 mL of sucrose in 100 mL of water does not yield 200 mL of solution.
Nonetheless, it’s possible to determine the concentration of ethanol by measuring the density of the solution and comparing that value to ethanol-water density tables available in the CRC handbook and similar publications. We’ll measure the density of the starting solution and the resulting distillate and compare those values to published values to determine the ethanol concentrations of the two solutions.
Required Equipment and Supplies
- goggles, gloves, and protective clothing
- balance (optional, but strongly recommended)
- cylinder, graduated, 100 mL
- flask, Erlenmeyer, 250 mL
- thermometer (optional)
- stopper, 2-hole (to fit Erlenmeyer flask)
- glass tubing, 75 mm (2)
- flexible tubing, ~ 50 cm
- pipette, disposable
- hotplate or kitchen stove burner
- coffee can, pickle jar, or similar tall wide-mouth container
- glycerol (see Substitutions and Modifications)
- ethanol, 70% (100 mL)
- tap water
All of the specialty lab equipment and chemicals needed for this and other
lab sessions are available individually from Maker Shed or other laboratory
supplies vendors. Maker Shed also offers customized laboratory kits at special
prices, including the Basic Laboratory Equipment Kit, the Laboratory Hardware Kit, the Volumetric Glassware Kit, the Core Chemicals Kit, and the
Supplemental Chemicals Kit.
Ethanol solutions of 45% or greater concentration are flammable. Ethanol vapor is extremely flammable. Handle the hot liquids used in this experiment with extreme care and have a fire extinguisher handy. The distillate produced in this experiment is not safe to drink. Ethanol is denatured with compounds that form azeotropes with ethanol, specifically to ensure that denatured ethanol cannot be undenatured by distillation. Wear splash goggles, gloves, and protective clothing.
Substitutions and Modifications
- If you do not have a balance, you can still do this experiment to observe the distillation process, but you will not be able to determine the ethanol concentration of the distillate.
- If you do not have a thermometer, plug the second hole in the stopper with a small piece of duct tape or simply by dropping a disposable pipette into the hole. (Not much pressure is generated because the ethanol vapor can escape from the flask through the glass and flexible tubing.)
- If you do not have glycerol (sold in drugstores under that name, or as glycerin), you can use vegetable oil or mineral oil to lubricate the stopper before you insert the thermometer and glass tubing. If you use oil, wash the stopper thoroughly with soap or detergent as soon as possible after you complete the experiment. Allowing any type of oil to remain in contact with the stopper for too long can damage it.
- Ethanol is sold in drugstores, sometimes by name, and sometimes as ethyl alcohol, ethyl rubbing alcohol, or simply rubbing alcohol. (If the bottle is labeled as rubbing alcohol, make sure it’s ethanol rather than isopropanol.) Most drugstore ethanol is 70% by volume (about 64.7% by weight), but 90% and 95% concentrations are also common. If you start with more concentrated ethanol, dilute it with water to 70% before you begin the distillation.
Figure 6-2 shows the distillation apparatus I originally used for this lab session. Note that both flasks are securely clamped to prevent tipping, and that the stopper in the receiving flask (right) has an open hole to vent the flask and prevent pressure buildup. The receiving flask sits in a large container of ice water. As the alcohol vapor passes through the tubing and into the receiving flask, it condenses immediately when it contacts the cold air in the receiving flask. For this modified version of the lab session, we’ll use a simplified distillation apparatus, substituting a 100 mL graduated cylinder for the second Erlenmeyer flask as the receiving vessel. We’ll also use a thermometer in the distillation flask to provide a visual indication of when all of ethanol has distilled over to the receiving vessel, leaving only water in the distillation flask.
Figure 6-2. The distillation apparatus I used for the lab session in the book
- If you have not already done so, put on your splash goggles, gloves, and protective clothing.
- Weigh the clean, dry 100 mL graduated cylinder on your balance and record the mass on Line A in Table 6-2.
- Transfer 100.0 mL of ethanol to the graduated cylinder, using the disposable pipette to add the last few drops to bring up the volume to 100.0. Reweigh the graduated cylinder and record the mass on Line B in Table 6-2.
- Subtract Line A from Line B and record the result (the mass of the 100 mL of ethanol) on Line C. Divide Line C by 100 to give the density of the ethanol in g/mL and record this value on Line D.
- Transfer the 100 mL of ethanol from the graduated cylinder into the 250 mL Erlenmeyer flask and set the flask aside. Dry the inside of the graduated cylinder thoroughly.
- Place a drop of glycerol or other lubricant on the bulb end of the thermometer and carefully slide the thermometer into one hole of the two-hole stopper. Wear gloves to protect your hands if the thermometer breaks, and press gently with a twisting motion. Do not force it. You want the bulb end of the thermometer to protrude a few cm below the bottom of the stopper, making sure the range from about 70 °C to 100 °C is visible above the stopper.
- Lubricate one of the 75 mm glass tubes and insert it into the other hole in the rubber stopper as described above. (If the ends of the glass tube are not already fire-polished, fire-polish both ends and allow the tube to cool before proceeding.) Insert the glass tube until just a few mm protrude below the bottom of the stopper, leaving as much as possible of the tube protruding from the top of the stopper, where you will connect the flexible tubing.
- Connect one end of the flexible tubing to the glass tube protruding from the stopper, and insert the other (fire-polished) glass tube 1 cm or so into the other end of the flexible tubing. (You can lubricate the glass tubes slightly with glycerol where they connect to the flexible tubing, but that’s generally not necessary for a temporary setup. If you leave the glass tubes inserted into the flexible tubing, they may eventually bind to and split the flexible tubing.)
- Insert the rubber stopper assembly carefully into the neck of the 250 mL flask. Press down gently to ensure a seal, but don’t force the stopper too far into the neck of the flask.
- Place the 100 mL graduated cylinder in the coffee can or other container and fill the container with ice water. You needn’t use a lot of ice, but make sure that there are still at least a few chunks of ice floating in the water after it has cooled down completely. You want the water level to be as high as possible against the outside of the graduated cylinder, but not so high that there’s any risk of any water being transferred into the graduated cylinder.
- Place the 250 mL Erlenmeyer flask on the hotplate, position the coffee can as necessary, and place the free end of the glass tube on the flexible tubing inside the graduated cylinder, making sure that the tip is well below the level of the ice water but not so far into the graduated cylinder that it will be submerged by distillate as it forms.
- Turn on the hotplate to its lowest setting and observe the liquid in the flask. As the liquid begins to boil gently, the thermometer reading increases to indicate the temperature of the vapor that fills the flask. That temperature is the boiling point of the fraction that is being vaporized, which is an ethanol/water mixture with an ethanol concentration higher than 70% (and approaching the 95% concentration of the low-boiling azeotrope).
- Continue boiling the contents of the flask gently, adjusting the heat as necessary, to maintain the liquid at a gentle boil. As the more concentrated ethanol boils off, the vapor moves through the flexible tubing and into the graduated cylinder, where it cools and condenses into liquid.
- Observe the thermometer carefully. As long as any ethanol remains in the distillation vessel, the temperature of the vapor remains at the boiling point of the ethanol/water azeotrope. When all of the ethanol has boiled off, you’ll see a sudden increase in temperature, which indicates that you are now boiling off pure water as steam. At that point, immediately turn off the hotplate and carefully remove the end of the (hot) tubing from the graduated cylinder. Place the tubing end aside in a beaker or other safe location, as some steam will continue to come over until the distillation vessel has cooled.
- Remove the graduated cylinder from the cold water bath, being careful not to spill any of the contents or to allow any liquid from the water bath to enter the graduated cylinder. Dry the exterior of the graduated cylinder completely, and place it aside to allow it to equilibrate to room temperature.
- After the graduated cylinder has reached room temperature, make sure the exterior is completely dry and weigh the graduated cylinder with its contents. Record the combined mass on Line E of Table 6-2.
- Subtract the empty mass of the graduated cylinder (Line A) to determine the mass of the distillate. Enter that mass on Line F of Table 6-1.
- Determine the volume of distillate in the graduated cylinder and record that value as accurately as possible on Line G of Table 6-1.
- Determine the density of the distillate by dividing its mass by its volume. Enter the value you calculate on Line H of Table 6-1.
- From the density you calculated in the previous step, determine the concentration of ethanol in the distillate by looking up density values for aqueous ethanol solutions in a handbook or other source. (One convenient source is the Wikipedia Ethanol data page.)
Table 6-2. Distillation of ethanol – observed and calculated data
|A. mass of 100 mL graduated cylinder||___.___ g|
|B. mass of 100 mL graduated cylinder + 100.0 mL 70% ethanol||___.___ g|
|C. mass of 100.0 mL 70% ethanol (B – A)||___.___ g|
|D. density of 70% ethanol (C / 100)||___.___ g/mL|
|E. mass of graduated cylinder with distillate||___.___ g|
|F. mass of distillate (E – A)||___.___ g|
|G. volume of distillate||___.___ mL|
|H. density of distilled ethanol (F / H)||___.___ g/mL|
You can retain the distilled ethanol and use it as fuel for your alcohol lamp or for any other purpose that requires denatured ethanol where concentration is not critical, or you can safely pour it down the drain. If you retain the distillate, label it properly.
If you have time and the required materials, consider performing these optional activities:
- Repeat the distillation, transferring the original distillate to the distillation vessel for re-distillation.
- Run the distillation a third time, and determine the density (and ethanol concentration) of the third distillate.
Q1: Would you expect the density of the distilled ethanol solution to be lower or higher than the density of the original ethanol solution? Why?
Q2: Using values for the densities of various concentrations of ethanol and water that you obtain from the CRC handbook or a reliable on-line source, estimate the ethanol concentrations in the original solution and the distillate.
Q3: A drugstore offers denatured ethanol in concentrations of 70%, 95%, and 99% by weight. The 70% and 95% solutions are relatively inexpensive, but the 99% solution is very costly. Why?
Q4: Distilling wine produces not colorless pure ethanol as you might expect, but brandy, which is deeply colored and contains complex flavors. Why?
August 24, 2009
49 thoughts on “Laboratory 6.2: Distillation: Purify Ethanol”
A few points:
– though calcium chloride is a drying agent, it would not normally be used to dry ethanol or other alcohols because it is soluble in many of them (including ethanol). A sulfate drying agent (magnesium sulfate, sodium sulfate, copper sulfate, calcium sulfate…) would be more appropriate to this example.
– in your list of equipment for this example, it would be helpful to specify the diameters needed for the glass tubing and flexible tubing, as well as perhaps specifying the material for the flexible tube (not everything is going to resist ethanol vapor, for example clear vinyl might not be a great choice). Of course these are going to depend on the size of the holes in the stoppers; typically we might see a 5mm diameter hole, which would be appropriate for a 6mm OD or 1/4″ OD glass tube, and then you’d want a flexible tube with a slightly smaller or equal internal diameter. Simpler, however, is to get a somewhat stiff flexible tube whose outer diameter matches (or rather slightly exceeds) the stopper hole, and stuff that into the stopper directly. This eliminates the need for glass tubing.
– your point #14 is actively wrong. There is no sudden discontinuity in vapor temperature ‘when all the ethanol has boiled off’, nor will the vapor normally match the azeotropic composition exactly. Rather, the vapor that comes from the boiling liquid *will always be much closer* to the azeotropic composition than that of the liquid. Thus boiling a liquid that is 40% ethanol might get you vapors that are 80% ethanol, while boiling beer would get you vapors that might be 25% ethanol. These numbers are wild guesses on my part but you can look up a distillation chart for ethanol that would give you exact values; distillers have been studying this for hundreds of years.
By extension, the vapor composition gradually becomes more water-rich (and hotter) as the distillation goes on, just as the boiling liquid does. But it’s not discontinuous.
– as a practical matter, your setup has poor heat transfer from the receiver/condenser to the vapor. If boiling is rapid, vapors will tend to escape through the pressure relief. A great improvement can be achieved just by using a longer flexible tube and looping some of it through a saucepan or other suitable container of cold water. This also makes the use of ice superfluous.
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Assuming that you’re making the alcohol then adding water to dilute it and drink it, I’d be worried about the impurities companies put in the drug store ethanol to begin with!! aka isopropyl alcohol, acetone, methyl ethyl ketone, methyl isobutyl ketone, and denatonium.
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