Ethanol Determination by Distillation

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Fermentation

In this lab, you will use distillation as a means to study the stoichiometry of the chemical reaction involved in alcohol fermentation as performed by yeast.

Many microorganisms, notably yeasts and bacteria, extract energy from their food (glucose) by fermentation. The overall chemical reaction for this is:

C₆H₁₂O₆ (glucose) → 2CO₂ + 2CH₃CH₂OH (ethyl alcohol)

or, starting from sucrose or maltose,

C₁₂H₂₂O₁₁ + H₂O → 4CO₂ + 4CH₃CH₂OH (ethyl alcohol)

It was this process which Louis Pasteur studied, leading to a biochemical understanding of biological processes. Humans have known about and utilized the process of fermentation for many thousands of years. CO₂ liberated by yeast cells doing alcohol fermentation causes bread to rise. The Egyptians and many subsequent civilizations have fermented grains such as barley to break the starch down to malt (maltose), then glucose, and finally alcohol. For at least that long, people have known that various fruits, especially grapes, could also be fermented to produce alcoholic beverages. Thus, alcoholic fermentation is the process which is responsible for the production of wine, beer, and other fermented products. It is the toxic nature of ethanol which acts to preserve these brews, and which leads to intoxication upon consumption. In fact, yeasts cannot generally survive in alcohol concentrations in excess of approximately 12 to 14%. In this lab, we will be studying the process of fermentation as performed by the yeast, Saccharomyces cerevisiae (sacchar = sugar; myces = fungus; Ceres = goddess of grain; vis = to see; -ia = state of, condition of, disease).

Stoichiometry of Fermentation

Today’s lab can be used to demonstrate a number of important principles relating to alcoholic fermentation including the stoichiometry (the study of the definite proportions in which chemicals will interact with each other — stoichi = an element, metr = measure) of the conversion of sugar to alcohol. The following conversion factors will be important, here:

Molecular Weight of Sucrose/Maltose	342.17 g/m
Moles of EtOH per Mole of Sucrose/Maltose	4 m
Molecular Weight of Ethanol	46.05 g/m
Grams of EtOH per Mole of Sucrose/Maltose	4 × 46.05 = 184.20 g
Grams of EtOH per Gram of Sucrose/Maltose	184.20 ÷ 342.17 = 0.538 g
Density of Pure EtOH at 20° C	0.789 g/mL
Density of Pure H2O at 20° C	0.99823 g/mL
Weight of One Can of Malt Syrup	1.5 kg (3.3 lb)
Percentage of Weight of Maltose in Malt Syrup*	80%
Percentage of Weight of H2O in Malt Syrup*	20%
Weight of 1 C of Granulated Sucrose†	213.97 g
Density of Malt Syrup§	1.426 g/mL
*These percentages were found on another Web site †This was weighed here in our lab §This number was found on another Web site as 0.084 gal/lb, and since 0.084 gal     = 318.04 mL and 1 lb = 453.59 g, therefore 453.59 g/318.04 mL = 1.426 g/mL

Thus, if the fermentation reaction is 100% efficient, for every gram of sucrose or maltose that is put in, 0.538 g (= 0.682 mL) of ethanol should be produced.

In this lab, we may be testing samples of “lab brew” and/or root beer made in previous labs. Thus, it would be of interest to reexamine the recipes from those two labs and calculate the hypothetical stoichiometry involved.

When we made our root beer, the recipe called for 1 C of sugar (sucrose) per 2-L bottle. While most of that sugar was included to make the root beer taste sweet, and not all was intended for the yeast to ferment, if hypothetically, 100% of that would be converted to alcohol, that would be

1 C sucrose	×	213.97 g sucrose	×	0.538 g EtOH	=	115.12 g EtOH
2-L bottle		1 C		1 g sucrose		2-L bottle

115.12 g EtOH	÷ 2 =	57.56 g EtOH
2-L bottle		L of root beer

57.56 g EtOH	×	1 mL EtOH	=	72.95 mL EtOH
L of root beer		0.789 g EtOH		L of root beer

For our lab brew, the amount of sugar going in depended on which recipe was followed. The original, Fankhauser recipe calls for 6 C of sucrose + 1⅓ C of canned malt syrup. Alternatively, we sometimes make the lab brew with one, whole, 1.5-kg can of malt, and no added sucrose. As you, hopefully, recall, the recipe makes 5 gal, or about 10 2-L bottles of lab brew. Let’s start with the calculations for the latter situation, first, because they are easier:

1500 g syrup	×	0.80 g maltose	×	0.538 g EtOH	=	32.28 g EtOH
20 L of lab brew		1 g syrup		1 g maltose		L of lab brew

32.28 g EtOH	×	1 mL EtOH	=	40.91 mL EtOH
L of lab brew		0.789 g EtOH		L of lab brew

Then, for the sucrose in the original, Fankhauser recipe,

6 C sucrose	×	213.97 g sucrose	×	0.538 g EtOH	=	34.53 g EtOH
20 L of lab brew		1 C sucrose		1 g sucrose		L of lab brew

and for the maltose,

1⅓ C syrup	×	236.64 mL syrup	×	1.426 g maltose	×	0.538 g EtOH	=	12.10 g EtOH
20 L		1 C syrup		1 mL syrup		1 g maltose		L of lab brew

which can be combined and converted to milliliters:

34.53 + 12.10 g EtOH	×	1 mL EtOH	=	59.10 mL EtOH
L of lab brew		0.789 g EtOH		L of lab brew

Expressing Concentrations of Ethanol Solutions

The concentration of ethanol solutions can be expressed in several ways. Often, molarity is not used, but rather concentration is expressed as either percent-by-weight (%w/w) or percent-by-volume (%v/v). Wine bottle labels show the amount of alcohol in their contents in terms of %v/v. The density of ethanol at 20° C is 0.789 g/mL, so pure ethanol is lighter than water (density of 0.99823 g/mL), but despite knowing that, in the case of alcohol solutions, it’s not easy to convert %w/w to/from %v/v. As you may have learned in chemistry, alcohol-water mixtures take up significantly less volume than the total of the components. For example, if you mix 100 mL of ethanol with 100 mL of water, you’ll end up with only about 192 mL of solution. However, since that 100 mL of ethanol weighs 78.9 g and that 100 mL of water weighs 99.8 g, even though you end up with “only” 192 mL of solution, that much would still weigh 78.9 + 99.8 = 178.7 g — it’s all there, but it takes up less space. In contrast, if you mix 100 g (= 126 mL) of ethanol with 100 g (= 100.18 mL) of water, you’ll end up with 200 g of solution. However, since the decrease in volume of various ethanol-water mixtures varies with the amounts of each included, there’s no easy way to predict what the final volume of that mixture will be.

However, in spite of how complicated all that seems to be, a lot of work has been done by analytical chemists studying the relationship between the specific gravity of ethanol solutions and their concentrations, expressed as either %w/w or %v/v, and these relationships have been published in various forms in a number of chemistry handbooks. Thus, similar to the Sugar in Soft Drinks lab you did, here also, we can weigh 100 mL of an ethanol solution and 100 mL of dH₂O, then calculate the specific gravity of that solution, and look that number up on a chart to determine either %w/w or %v/v of that solution. Since you may be more familiar with the %v/v used on alcoholic beverages, we will use a chart that correlates specific gravity with those numbers. Thus, the table for percentage of alcohol in a solution (excerpted from a larger one in one of the chemistry handbooks) which is included here and in your protocol is based on percentage by volume, in other words, how many milliliters of ethanol per 100 mL of solution.

Background on Distillation

The technique of distillation as a method of separating ethanol from the ferment will be demonstrated in this lab exercise. Distillation is the process of heating a solution to its boiling point, passing the vapors through a cooling device called a condenser, and collecting the liquid which condenses. Because the boiling point of ethanol is 78.5° C, so considerably lower than that of water (100° C), almost all of the alcohol will boil off first, followed by the water. Thus if you start with 100 mL of solution, by the time just under 100 mL has distilled, you will have collected all of the alcohol and almost all of the water. If you then q. s. the resulting distillate to exactly 100 mL with dH₂O, you will restore it to its original volume and concentration. Because of the alcohol it contains, the distillate will have a specific gravity lower than that of distilled water. By the use of the table provided, the percentage of alcohol in a solution can be determined by knowing the specific gravity of the solution, assuming that no interfering substances co-distill with ethanol which would affect the specific gravity of the distillate.

Interestingly, since the boiling point of ethanol is so far below the boiling point of water, the ethanol can be concentrated by the process of fractional distillation (collecting the distillate in aliquots rather than the whole solution). This is the source of liquors with higher alcohol contents, but is not what we will be doing in this lab.

A liquid must be brought just slightly above its boiling point before bubble initiation can begin to start it boiling. As a bubble of vapor appears within the liquid, it may do one of two things: if it is below a minimum size, it will collapse because of the surface tension of the liquid, or if it is larger than the critical size, grow and rise to the surface of the liquid. If a liquid, which is free of solid impurities or dissolved gases, is heated slowly, a temperature much above the boiling point can be reached without any boiling actually taking place. This superheating occurs because extra energy is required before bubble formation is initiated. If a bubble should start to form in such a superheated solution, it might suddenly grow with almost explosive violence enough to shatter the container. This problem, called bumping, can be overcome by adding a boiling chip, a piece of porous material, to the liquid before it is heated to the boiling point. The pores act as built-in bubbles so that a liquid cannot superheat. As the distillation proceeds, the air in these pores is replaced by vapor of the distilling material, but this vapor cannot condense because the temperature of the liquid is just slightly above its boiling point. However, whenever a distillation is stopped and then started again, a new boiling chip will be needed because the vapor in the pores will have condensed and filled the pores with liquid.

(This, by the way, has direct practical application. There have been numerous cases in which someone tried to heat water to boiling in a clean glass or china container placed in a microwave. When the person checked on the water, it wasn’t boiling, yet, so the person “nuked” it for a longer time, and then again, and again, without it ever starting to boil. Then, in frustration, the person removed the container from the microwave, jostling the water in the process, and simultaneously looked into it. However, the process of jostling the water initiated bubble formation in this superheated water, causing it to bump and explode out of the container, severely scalding the person’s face. The “moral of the story” is, if you need to heat water in a microwave, always make sure to include some porous substance such as a toothpick, teabag, or loose tea leaves, etc. that will serve as a site of bubble initiation, thereby allowing the water to boil without bumping.)

Safety Considerations

• Remember that while Bunsen burners are in use, long hair should be tied back (anything long enough to go in a rubberband counts). Secure all loose clothing, sleeves, etc.
• Because, as explained above, possible bumping could cause a flask to explode, goggles should be worn while working near a running distillation apparatus.
• While many of them have broken and been replaced, some of the thermometers in the stoppers we’re using may contain mercury (the “silver stuff”), which is quite toxic to humans. Your group should try to obtain a stopper with a (red) spirit thermometer in it, and only use a mercury-containing thermometer if there are not enough spirit thermometers for everyone. If you are using one of the remaining mercury thermometers, handle it VERY carefully! If one of these thermometers does break, notify your instructor and/or the lab manager AT ONCE! Spilled mercury must be cleaned up using a special mercury spill kit. Students are NOT allowed to do this — only trained lab personnel are to clean up mercury spills.

Distillation of Ethanol Solutions

1. Work in groups for this lab.
2. Locate a spot for your group to set up your apparatus. This needs to be near a gas outlet and a sink.
3. Each group should obtain the equipment needed to set up your distillation apparatus:
   • gas outlet, cold water tap, and sink in close proximity to each other
   • a pre-assembled “lid” consisting of
     • #10 two-holed rubber stopper
     • 15-cm glass tubing, bent 60 inserted into one hole
     • #4 one-holed rubber stopper attached to other end of glass tubing
     • thermometer,-10 to 110° C inserted into the other hole
   • 2 ring stands
   • 1 universal clamp
   • 1 large ring
   • 1 wire gauze (screen)
   • 2 approx. 1-m lengths of rubber tubing
   • 1 Bunsen burner with rubber tubing attached, plus striker
   • 1 Liebig condenser
   • 500-mL wide-mouth Erlenmeyer flask
   • 100-mL volumetric flask, clean & dry
   • funnel that will fit into volumetric flask without tipping over
4. Your group will also need these supplies as you do the lab:
   • access to a balance
   • “regular” thermometer
   • 100 mL of dH₂O
   • ice bath
   • 125 to 150 mL of decanted lab brew (possibly root beer or hard cider if available?)
   • 100-mL grad. cylinder
   • 250-mL Erlenmeyer flask
   • several boiling chips
   • several drops of vegetable oil
   • refractometer (share with the rest of the class)


5. While others are setting up the distillation apparatus as described below, one person from your group should weigh the DRY volumetric flask to the nearest 0.01 g (you may use the electronic balance), and everyone in the group record the weight in his/her notebook (be careful — this in your only chance to get the dry weight!).
6. Then, that person should fill the flask with dH₂O to the bottom of the neck (don’t adjust it to exactly 100 mL, yet). Then, with a thermometer in the water, use the ice bath (as per the Sugar in Soft Drinks lab) to adjust the temperature to exactly 20° C (remember to remove the flask from the ice bath when the water is a little above 20° because the temperature will continue to drop a bit).
7. Then, q. s. to exactly 100 mL, completely dry the outside of the flask, make sure there are no air bubbles in the water, and weigh the flask-water combination, again to the nearest 0.01 g. Everyone should record the flask + water weight. As in previous labs, subtract to determine the weight of the dH₂O (record this in your notebooks, too). Place the “used” water in the designated container.


8. Another person from the group should pour about 125 mL of the decanted lab brew (or root beer) into a 250-mL Erlenmeyer flask and shake vigorously to degas. (Note that because the boiling point of ethanol is so low, we cannot heat the solution to degas it like we did with our soft drinks.)


9. While one person is weighing the flask and water and another person is shaking the lab brew, others in the group should set up a distillation assembly as pictured here and as demonstrated by your instructor.
10. 1. Attach the Bunsen burner to the gascock.
    2. Place the ring on one of the ringstands, about 5 cm above the mouth of the burner, and make sure the ring is horizontally straight. If the ring slopes, you may need to turn it over to get it level.
    3. Then, set the wire gauze, asbestos pad down, on the ring. (The asbestos pad under the flask moderates the amount of heat transferred to the flask so that control of distillation can be maintained.)
    4. Use one piece of tubing to attach the lower inlet on the condenser to the cold water source.
    5. Attach one end of the other piece of tubing to the upper outlet on the condenser and the other end should run into the sink. Check the tightness of the fittings. Tubing should be attached to the condenser such that water is flowing from the BOTTOM to the TOP.
    6. Attach the “lid” (thermometer, etc.) assembly lightly on the 500-mL flask and set that on the wire gauze. If possible, set things up so that you can read the thermometer (may or may not be possible depending on location of sink, etc.) when the smaller stopper is pointing toward the condenser. Thermometers cannot be turned around within the stopper (don’t try — it’ll break and cut you!), so if there are enough extras, you may exchange the “lid assembly” for a different one if you need one facing the other direction (supplies are limited, so this may not be possible).
    7. Place the funnel into the 100-mL volumetric flask.
    8. Attach the clamp to the second ringstand, and center the condenser in the clamp. The condenser should be oriented such that the cut angle of the condenser “spout” is pointing down and the inlet and outlet (attached to the tubing) are pointing upward.
    9. Adjust the height and angle of the clamp so that the wide (top) end of the condenser will connect securely with the “lid” assembly, and the lower end of the condenser is just above the funnel.


11. Use the 100-mL graduated cylinder to measure exactly 100 mL of the degased lab brew (if you measure carefully, the percentage of error from using the cylinder rather than a volumetric flask will be negligible). Pour this into the 500-mL wide-mouth Erlenmeyer flask, being careful to get it all. Rinse the graduated cylinder with 50 mL of dH₂O and add this to the 500-mL flask. That will do two things: get any remaining lab brew that was in the graduated cylinder, and add some “extra” water to what will be boiled so it doesn’t boil dry.
12. Add 2 or 3 boiling chips to the 500-mL flask. Then, add two or three drops of vegetable oil to help prevent foaming. Replace the 500-mL flask on the ring/wire gauze and firmly attach the “lid” assembly. Attach the 1-hole stopper firmly to the condenser. Make sure the funnel and volumetric flask are under the spout of the condenser.
13. Gently turn on the cold water supply for the condenser. Adjust the flow to about 100 mL/10 sec (time how long it takes to fill a 100-mL beaker). The backflow preventers on the faucets may prevent you from adjusting the flow to exactly 100 ml/10 sec, but come as close as you can. A tiny bit faster is better than too slow. Double check that all stoppers are snugly seated into their joints.
14. When you have everything set up and ready to go, the water flow adjusted, etc., then ask your instructor to check out the assembly and verify that it’s all OK.
15. Once you have the OK from your instructor, light the Bunsen burner, properly adjust the air and gas for a hot flame,, and place the burner under the boiling flask. Monitor the temperature of the contents of the boiling flask. When the temperature of the liquid reaches 80 to 85° C, reduce the flame so the liquid boils slowly. BEWARE OF FOAMING! If the liquid starts foaming or boiling too vigorously, be ready to pull the burner out from under the flask until you reduce the flame.
16. Once the foam has subsided, the size of the flame may be carefully increased. Once drops of distillate are dripping from the condenser into the volumetric flask, adjust the flame such that the distillation rate is 2 drops/sec.
17. Monitor the volume of distillate in the volumetric flask. Turn off the burner when the fluid level reaches the lower end of the neck of the volumetric flask. Because EtOH has a lower boiling point than water, this will get all the EtOH and most of the H₂O from the lab brew. When the volume of distillate in the volumetric flask has reached this level, turn off the gas and remove the volumetric flask.
18. As was done with the dH₂O, use a thermometer and ice bath to adjust the temperature of the flask and contents to exactly 20° C (remember to remove the flask from the ice bath when the distillate is a little above 20° because the temperature will continue to drop a bit).
19. Q.s. with 20° C dH₂O to 100.00 mL (thus restoring the original volume and concentration of EtOH). If there are droplets of distillate clinging to the thermometer, you should use a couple drops of the water you’re adding to rinse them into the flask. Dry off the outside of the flask, then weigh flask + distillate and record the weight in your lab notebook. Subtract to determine the weight of 100.00 mL of distillate.


20. As was done in the Sugar in Soft Drinks lab, determine the specific gravity of the distillate by dividing the weight of the distillate by the weight of the dH₂O.

    wt distillate @ 20° C	× sp gr of distillate
    wt dH2O @ 20° C

    Using the table below or in your protocol, locate the closest specific gravity and read the corresponding percentage by volume of EtOH.
21. Hold a piece of Parafilm over the end of your flask with a finger or thumb, and turn the flask upside down several times to make sure the contents are evenly mixed.
22. Put a couple drops of dH₂O in the refractometer. With a bright light source shining on the refractometer, look through it and focus it by turning the black band around the eyepiece. You should see a blue area and a white area. Notice where the boundary between these two crosses the scale on the left-hand side of your view. With plain dH₂O in the refractometer, this boundary should be at the bottom of the left-hand scale, at 1.3330. If it is not, there is a metal screw that can be turned to adjust it.
23. Use a Kimwipe to wipe the water off the refractometer. Then, place a few drops of your distillate in the refractometer, and close the lid. Notice where the blue-white boundary crosses the left-hand scale, now (1.33??). Record this number in your lab notebook and use the right-hand columns on the provided table to determine the corresponding percentage by volume of EtOH. (Be very careful not to bump the screw on the top — it is used to set the refractive index of dH₂O to 1.3330, and if it gets moved, your reading will be off.) Remember to interpolate.
24. When you are done, wipe the distillate off the refractometer with a Kimwipe. Please leave the refractometer clean and dry when you are done with it.
25. Clean up! Rinse out your glassware, but do not let the boiling chips go down the drain. Place them in the trash. Place all wet glassware, tubing, etc. in the designated locations.
26. Someone from your group should submit your group’s data online.
27. For an extra mathematical challenge: how does the amount of alcohol in the fermented lab brew or root beer relate to the starting quantity of sugar? Try to figure out how much sugar must have been present in the original batch of brew to make that percentage of alcohol. Conversely, knowing how much sugar was added, what percentage of it was actually used/fermented?

Table of Specific Gravity and Percent Alcohol

Other Things to Include in Your Notebook

Make sure you have all of the following in your lab notebook:

• all handout pages (in notebook or separate protocol book)
• all notes you take during the introductory mini-lecture
• all notes and data you gather as you perform the experiment
• all requested calculations based on those data
• print-out of class data
• drawing (yours!) of the distillation set-up with all key features labeled and annotated as needed
• drawings (yours!) of a refractometer and of the scale inside the refractometer
• answers to all discussion questions, a summary/conclusion in your own words, and any suggestions you may have
• any returned, graded pop quiz