Simple and Fractional Distillation of a Binary Mixture

Introduction/Background

Distillation is a process that is present in many common liquids such as perfumes, or even alcoholic beverages. It is through distillation that it is possible to have liquor with high alcohol content. Fruit and plant material produce dilute ethyl alcohol when fermented, and through distillation the ferment is transformed into concentrated and purified ethanol. The process of distillation consists in purifying a compound by separating it from non or less-volatile material. This is done by converting a liquid into vapor, and then condensing it back into liquid form.

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 In fractional distillation, a fractionating column is added to the apparatus in between the stock mixture vial and the distillation head. This fractionating column contains several theoretical plates in which vapors that rise are collected in the condensate pool of each; this allows for all the intermediate residues to drip back into the stock mixture to be distilled again, making the distilled compound purer1. The fractionating column also allows for compounds with a close boiling point to be better separated than with just a simple distillation. The separation of these compounds is only possible due to differences in boiling point and vapor pressure between the two compounds in the binary mixture. Once a liquid is heated, this substance only starts to evaporate, or reach its boiling point, once the compound’s vapor pressure becomes equal to the external pressure acting on the liquid.  The compound with the lowest boiling point starts to evaporate first, and collecting in the distillation head, its vapor molecules are then passed through the condenser, where upon cooling, the molecules condensate back into liquid form, therefore leaving a purified liquid compound.

Boiling point is determined by strength of intermolecular forces, molecular weight, surface area and hydrogen bonds, therefore giving each compound its own boiling point. Moreover, to fully understand distillation, it’s important to understand that mixtures obey Dalton’s Law and Raoult’s Law. Dalton’s Law2 states that the vapor pressure of a mixture is equal to the sum of the partial vapor pressures of each compound in that mixture, and Raoult’s Law3 states that the partial pressure of a compound in an ideal mixture is equal to the vapor pressure of the pure compound at the same temperature, multiplied by the mole fraction. Another important concept to know when it comes to distillations is the term azeotropes. Azeotropes refer to a certain class of liquid mixtures that have a constant boiling point which may be higher or lower than the pure components and the liquid does not change composition when evaporating4.

The aim of this experiment was to separate and purify cyclohexane and toluene from a 1:1 binary mixture using a fractional distillation, and to create a plot of the process using a temperature versus volume graph.

Experimental section

 

 

 

Chemicals

Chemical

Chemical formula

Molecular weight

Melting point

Boiling point

Organic structure

IUPAC name

Chemicalproperties

Cyclohexane

C6H12

84.16 g/mol

6.47°C

81°C

Cyclohexane

Colorless, flammable liquid with a mild, sweet odor

Toluene

C6H5CH3 or C7H8

92.14 g/mol

-94.9°C

111°C

              CH3

Toluene

Colorless, flammable liquid with a benzene like odor

 

 

 

 

Results

Fractional distillation: Cyclohexane

Time

Temperature in ℃

Volume present in vial

0 – 1h25m

20℃

0

1h26m – 1h28m

20℃ – 30℃

0

1h29m

41℃

.3mL

1h30m

45℃

1mL

1h30m -1h35m

45℃

2mL

1h40m

44℃

2.87mL

1h42m

43℃

3mL

1h43m

39℃

3.06ml

1h45m

37℃

3.13mL

Fractional Distillation (after changing collecting vials): Toluene

Time

Temperature in ℃

Volume in collecting vial

0

32℃

0

1m8s

30℃

0

2m14s

28℃

0.03mL

1m8s – 22m

30℃ – 20℃

0.03mL

33m

25℃

0.03mL

34m

40℃

0.17mL

35m30s

60℃

1mL

36m

62℃

1.33mL

37m30s

70℃

2mL

38m20s

76℃

2.5mL

39m

77℃

3mL

SAMPLE CALCULATIONS:

Percent Recovery for Cyclohexane During Simple Distillation:

Percent Recovery =(Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent Recovery for Toluene During Simple Distillation:

Percent Recovery =( Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

SAMPLE CALCULATIONS:

Percent Recovery for Cyclohexane During Simple Distillation:

Percent Recovery =(Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent Recovery for Toluene During Simple Distillation:

Percent Recovery =( Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

SAMPLE CALCULATIONS:

Percent Recovery for Cyclohexane During Simple Distillation:

Percent Recovery =(Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent Recovery for Toluene During Simple Distillation:

Percent Recovery =( Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

SAMPLE CALCULATIONS:

Percent Recovery for Cyclohexane During Simple Distillation:

Percent Recovery =(Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent Recovery for Toluene During Simple Distillation:

Percent Recovery =( Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

SAMPLE CALCULATIONS:

Percent Recovery for Cyclohexane During Simple Distillation:

Percent Recovery =(Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent Recovery for Toluene During Simple Distillation:

Percent Recovery =( Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

SAMPLE CALCULATIONS:

Percent Recovery for Cyclohexane During Simple Distillation:

Percent Recovery =(Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent Recovery for Toluene During Simple Distillation:

Percent Recovery =( Actual Yield/Theoretical Yield)x100%

Percent Recovery = (3.0 mL/3.5 mL)x100%

Percent Recovery =86%

Percent recovery calculations from fractional distillationCyclohexane:Percent recovery = (Actual Yield/Theoretical yield) x 100%Percent recovery = (3mL/3.5mL) x 100 = 85.7%

Toluene: Percent recovery: Percent recovery = (Actual Yield/Theoretical yield) x 100%Percent recovery: (2.6mL/3.5mL) x 100 = 74.2%

To have a standard measure in mL, the number of drops to equal 1mL were counted. It was concluded that 1mL =  30 drops.

Distillation curve for fractional distillation

Distillation curve for fractional distillation

Discussion:

In this lab, the main purpose was to separate a binary mixture of cyclohexane and toluene into their pure compounds. This was done successfully, but the amount of cyclohexane and toluene retrieved at the end did not equal the amount that the original mixture had. This is referred to as percent recovery, the percentage of an original substance recovered after the distillation has been completed.  The original binary mixture had 3.5mL of cyclohexane and 3.5mL of toluene, after distillation only 3mL of cyclohexane, and 2.6mL of toluene were recovered.  Meaning that the percent recovery for cyclohexane was 85.7% and for toluene 74.2%. The percent recovery for cyclohexane presents a good number and indicates that the experiment went as desired in the first part. For toluene, the percent recovery is lower than 80%, which indicates that there must have been an error during the experiment. A plausible explanation as to why the percent recovery for toluene was much lower than for cyclohexane might be due to adding aluminum foil to the distillation head when waiting for the toluene drops to start. This was done due to the experiment taking abnormally long, 33 minutes to be exact, to start distilling toluene after already collecting cyclohexane. This made the temperature rise rapidly from 25℃ to 70℃ in only 2 minutes, and was the only thing done differently from the first part of the experiment, which can explain the difference in percent recovery. Another factor into play was that the experiment was stopped due to the stock mixture being close to burning out. This might have been due to the rapid rise in temperature, which may have caused particles to escape the apparatus due to a faulty connection, or similar.

The graph in the lab manual combines both the cyclohexane and toluene distillation as we as combining simple and fractional distillation into one graph. This differs from this report’s graphs due to not obtaining the data from the simple distillation, and therefore only having fractional distillation graphs, as well as the data not being recorded in regular increments. This was due to a distraction when cyclohexane first started dropping into the receiver, therefore not capturing the start of it. While the toluene distilled incredibly fast, in about a third of the time from the cyclohexane, making it hard to take values at regular intervals. The shapes of the graph mainly match the graph from the lab manual, with exception of the temperature dropping at the end of the cyclohexane, and in the very beginning of the toluene distillation.

Conclusion

 The aim of this experiment was to separate and purify cyclohexane and toluene from a 1:1 binary mixture using a fractional distillation. This was proved to be mainly successful as in the end a mixture of cyclohexane and toluene were obtained, although the percent recovery for cyclohexane was higher than for toluene. The information from the data revealed that toluene indeed has a much higher boiling point than cyclohexane, as the distillation of toluene was only concluded at much higher temperatures than for cyclohexane. Furthermore, the information from the data obtained proved how the difference in boiling points between compounds is a key component for distillation to happen.

The techniques used in this lab are quite essential in the real world, especially in the oil industry. It is by fractional distillation that crude oil is broken down into gasoline, which powers the majority of commercial vehicles5. The understanding of this technique can help improve the amount and purity of distillate collected. In turn, this can decrease the amount of fossil fuels used, both by obtaining higher amounts of gasoline from a crude oil barrel, as well by decreasing the energy used to perform such distillations, as there are more than 40000 distillation units across the country.

Overall, the lab accomplished what it was set out to do, demonstrating the relationship between difference in boiling points and the distillation process.

References

Nichols, Lisa. “Organic Chemistry Laboratory Techniques”, LibreTexts, 2019. https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Book%3A_Organic_Chemistry_Lab_Techniques_(Nichols)/5%3A_Distillation/5.3%3A_Fractional_Distillation

Dutton, F.B. “Dalton’s Law of Partial Pressures.”, Journal of Chemical Education Aug 1961. http://pubs.acs.org/doi/pdfplus/10.1021/ed038pA545.1

Kugel, Roger W. “Raoult’s Law: Binary Liquid-Vapor Phase Diagrams” Journal of Chemical Education 75.9. 1998

Clark, Jim. “Fractional Distillation of Non-ideal Mixtures (Azeotropes)”, LibreTexts, 2019. https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Physical_Equilibria/Fractional_Distillation_of_Non-ideal_Mixtures_(Azeotropes)

 Distillation and Filtration -Real life applications. Science Clarified, 2017. http://www.scienceclarified.com/everyday/Real-Life-Chemistry-Vol-2/Distillation-and-Filtration-Real-life-applications.html

Weldegirma, Solomon. Experimental Organic Chemistry: Laboratory Manual for CHM 2210L and CHM 2211L. University of South Florida. 2018