Preparation and Purification of Ethanol PDF

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Fermentation and Purification of Ethanol through Simple Distillation Introduction The process of distillation is a universal application that is most often implicated in the separation of liquid mixtures containing multiple components and in purification techniques. During distillation, two or more miscible liquids that are in solution together are heated until the vapor pressure of one is equal to the external atmospheric pressure—this will be the boiling point of that liquid. For distillation to be successful, the two liquids in the mixture must have distinct boiling points that will ultimately allow for the desired liquid to be extracted and recondensed, while the other liquid remains mostly in the solution (due to a higher boiling point that has not been reached yet). The most well-known application of distillation is the separation of ethyl alcohol (ethanol) from water to produce higher, more desired concentrations of ethanol that can be used in the production of more potent alcoholic beverages (such as liquors and spirits). Although several distillation techniques exist, the two most common applications are simple and fractional distillation. Simple distillation, the one that will be utilized in this experiment, is most commonly used when the boiling points of the two compounds of interest have a difference of approximately 25.0 degrees Celsius or more. This is a favorable choice for separating ethanol (boiling point of 78.4 degrees Celsius) from water (boiling point of 100.0 degrees Celsius). However, if the compounds being separated have a similar boiling point (typically much less than 25.0 degrees Celsius apart) then fractional distillation will be used. The main difference here is that fractional distillation utilizes a fractionating column, in which

the vapor slowly travels up, being re-evaporated and re-condensed, until the resulting vapor is almost completely pure in the lower boiling point compound (in our case, ethanol).

Figure 1: The basic set-up for simple distillation of an impure liquid (2).

The process of producing beer from the use of bacteria (or yeast, saccharomyces cerevisiae) is called fermentation. However, fermentation is only effective up to about 20% alcohol by volume, after which the yeast cannot survive and properly convert the sugars and starches into new constituents. This is why distillation is so important in the mass production of stronger forms of alcohol worldwide; it permits the production of solutions containing more than 20% alcohol by volume. Another problem arises in the separation of water and ethanol, the two miscible substances produce an azeotrope in solution—a mixture of two liquids that contains a constant boiling point. For this reason, ethanol cannot be purified past 95% by volume without chemically removing the remaining water molecules. The potential applications of distillation are practically endless and it will always be necessary for us to be able to make solutions of guaranteed purity. One of the most essential

applications of distillation with the primary objective of purity in our daily lives is that of water. We all want our drinking water to be completely pure, free of any and all contaminants that find their way into it. By boiling water in a distiller, one can free it from up to 99.5% of its original impurities, which include microorganisms (which are mostly inactivated by the boiling water), metals, nitrates, and other dissolved constituents that have boiling points that exceed that of water. However, other techniques exist for removing organic compounds with a boiling point similar to or lower than water. Gas vents and fractional distillation are both effective methods of removing volatile organic compounds (VOCs) that accompanies the water vapor during boiling (1, Dvorak and Skipton). These techniques are incredibly crucial for preventing the contamination of “clean” drinking water that could have harmful effects on other organisms, including humans.

Experimental Fermentation of Sucrose During the first week of this experiment, a solution was created by mixing 2.03 grams of sucrose, 0.35 grams of Pasteur salts (that provide nutrients to yeast and act as a buffer), 2.02 grams of yeast (Saccharomyces cerevisiae), and 195.0 mL of water in a 250-mL Erlenmeyer flask. The flask was stoppered, agitated, covered with parafilm, and store in a cabinet for two weeks to allow fermentation to proceed. Ethanol Purification by Simple Distillation After setting up the simple distillation apparatus (see Figure 1), the fermented solution was decanted into the 250-mL round bottom flask and combined with boiling chips to ensure even boiling. 10 mL of the original ethanol solution was stored in a screw-cap vial for further

analysis. As the Bunsen burner heated up the solution and vaporized the contents, 10-mL increments of the condensed vapors were collected and stored separately with labels distinguishing them from one another. The first 10 mL increment, the fourth increment, and the seventh increment were saved while the remaining 40 mL collected were discarded. The initial and final temperature that corresponded with all seven fractions were recorded, as temperature can have an effect on density and what molecules were being evaporated.

Figure 2: simple distillation apparatus used for this experiment. Purity Determination By measuring out 5.0 mL of all four ethanol samples (including the original sample) and weighing them on a scale that is tared with the container being used, the density of each ethanol solution can be determined. These density values were then compared with values from Steffen’s Chemistry Pages (3) that compares density values with their corresponding concentrations (% by volume).

Results Fraction

Mass (of 0.5 mL)

Density (g/mL)

Purity (% by

Average Head

volume) Temperature Original Mixture 4.61 grams 0.922 g/mL 54.7% N/A 1 (0-10 mL) 4.58 grams 0.916 g/mL 58.0% 92.95 °C 4 (30-40 mL) 4.75 grams 0.950 g/mL 39.6% 98.95 °C 7 (60-70 mL) 4.31 grams 0.862 g/mL 79.5% 99.00 °C Figure 3: data obtained from distillation of three separate fractions and the original sample. Purity values obtained from citation #3, Steffen’s Chemistry Pages.

Ethanol Purity vs Average Temperature 100 90 Ethanol Purity (% by volume)

80

79.5

70 60

58

50 40

39.6

30 20 10 0 92

93

94

95

96

97

98

99

100

Head Temperature (in °C) Ethanol Purity (% by Volume)

Linear (Ethanol Purity (% by Volume))

Figure 4: illustrates relationship between average head temperature during distillation and the purity of the ethanol obtained.

Ethanol Purity vs Average Temperature (Class Average) Ethanol Purity (% by volume)

Linear (Ethanol Purity (% by volume))

60 Ethanol Purity (% by volume)

54.47 50 42.64

40 33.61

30 20 10 0 82

83

84

85

86

87

88

89

90

91

Head Temperature (in °C)

Figure 5: class average relating the average head temperature to density/purity of the ethanol solution

Discussion The original mixture of ethanol procured from the 250-mL Erlenmeyer flask that had been fermenting for two weeks showed a density of 0.922 grams per milliliter, which corresponds to an ethanol purity of about 54.7% by volume. Theoretically, distillation should only help raise the purity of ethanol by effectively separating it from the water leftover in the container. The first 10 mL tested from the distillation process was only about 3.3% more pure than the original sample, which is not a significant change but still displays a difference between the two techniques (microbial fermentation versus simple distillation). The fourth fraction exhibited a dramatic decrease in purity of about 15%, and the final fraction (the last 10 mL obtained from the sample) was nearly 25% purer than the original sample. Therefore, fraction 7, with the lowest density (0.862 g/mL), was the purest solution of ethanol out of the

four test samples. A 79.5% purity is still a good distance away from the theoretical maximum value of 95% ethanol, indicating that many impurities are still contained in fraction 7 or that too much excess water vapor reached the condenser due to a relatively high head temperature of 99.00°C. Because ethanol has a pure density of 0.789 g/mL, density values that are closer to this value are going to contain more ethanol, while values closer to 1.00 g/mL (the density of pure water) will indicate a higher concentration of water than ethanol. Because the average head temperature circulated just under the boiling point of water (100 °Celsius), it is reasonable to assume that a large amount of water vapor evaporated along with the ethanol vapor, but at a slower rate since the boiling point of ethanol is significantly lower (78.4 °Celsius). If the experiment had been performed much more slowly, and with a lower average head temperature, the overall ethanol purity levels would be expected to be much higher since the relative amounts of ethanol vapor versus water vapor would increase. For example, if the experiment had been performed under a constant temperature of about 82-85 degrees Celsius, the liquid ethanol molecules would have sufficient kinetic energy to escape into the vapor phase while water molecules would have a hard time doing the same. The unusually low ethanol concentration of fraction four may have been due to the vapor temperature nearing 100 degrees Celsius and allowing too many water molecules to vaporize and wind up in the collection beaker. Based on collected data from the whole lab section, ethanol purity was achieved at higher concentrations when boiled at a lower temperature. As one can see from Figure 5, ethanol purity averaged around 55% purity at 82.50 °C, while a higher temperature of 90 °C resulted in a purity of 42.64%. This makes sense, because more water molecules are going to escape the solution if the head temperature is closer to its boiling point. The theoretical

azeotropic boiling point to produce 95% ethanol is closer to 78.2 °C, slightly lower than the boiling point of pure ethanol. This method of boiling would ensure that the most ethanol molecules possible will be evaporated and re-condensed in relation to water molecules and other impurities in the fermented solution.

Citations: 1) Dvorak, Bruce I., and Sharon O. Skipton. "Drinking Water Treatment: Distillation." NebGuide. University of Nebraska-Lincoln Extension, Institute of Agriculture and Natural Resources, Dec. 2013. Web. 17 Sept. 2015. . 2) Setup for Distillation. Digital image. Separation Technology. N.p., 19 May 2012. Web. 17 Sept. 2015. . 3) Thomas, Dr. "Ethanol - Water - Mixtures." Steffens Chemistry Pages. WordPress, n.d. Web. 17 Sept. 2015. .

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