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LABORATORY EXPERIMENT #04: DISTILLATION CHEM 12A GISELA GÓMEZ-CRUZ

Introduction: The purpose of this experiment was to build an empirical basis of simple and fraction al distillation and then using the empirical evidence, determining a theoretical explanation for it all/ Dalton’s Law of pressure as well as Raoult’s law were also reviewed to further apply them in distillation. Gas chromatography was also used in order to determine the number of theoretical plates generated in both types of distillation. These theoretical plates would then assist in proving that one method is actually a superior method. [idk if to add the last objective bc idk if we actually did that…] Experimental: 1.)Simple Distillation For this type of distillation, we first used automatic pipets to add equal amounts (2.0ml) of hexane and of toluene into a 5.0ml conical vial along with a boiling stone inside. This conical vial was then attached to our distillation with Hickman-Hinkel still head apparatus (figure 1). We made sure the bulb of the thermometer went below the reservoir (just past the side arm) so not to touch any of glass surface. The bottom of the conical vial was submerged into the sand bath while not touching the bottom of the sand bath (too close to hot plate). We then covered the vial and the outside thermometer with a piece of aluminum foil right below the side arm to shield the Hickman-Hinkel still head from overheating since it is so close to the hot plate. The san bath was then heated at a rate of 5  C/min to reach a temperature of 90-100  C. Once the boiling started, we adjusted the heating rate to 2  C/min. Once the temperature of the vapor reached 67-83 C (bath temperature ~95-105 C), we sued a Pasteur pipet to collect the condensate through the side arm and transferred it to a perfume vial until it was ~3/4 full. Capped the vial immediately to prevent it from escaping and labeled it to save for later. The Pasteur pipet was left inside of the side arm the entire time to help slow down the escape of vapor from the Hickman-Hinkel still head. We collected a second and third fraction at vapor pressure of 83-103  C and 103-112  C respectively. We capped and labeled all of the perfume vials which we also stored in the refrigerator until the next class to analyze by gas-liquid chromatography. Once the apparatus had cooled down, we disassembled it and cleaned everything up. 2.) Fractional distillation using a spinning band column For this distillation, we placed a teflon spinning band in the 5.0ml conical vial and added the equal portions (2.0ml) of hexane and toluene. The conical vial was then attached to the distillation apparatus with condenser piece added (figure 2). We had to turn on the magnetic stirrer so that the band would spin. Just like we did in the simple distillation experiment, we heated the solution at the same rates and collected three fractions at the specific vapor pressure temperatures. The spinning had to be controlled carefully ensuring that it was slow at first and then faster during boiling and the fastest once when the vapor entered the condenser column. The three fractions were collected in perfume vials each ~3/4 full and labeled to be stored for further gas-liquid chromatography analyzation. Once the apparatus had cooled down, we disassembled it and cleaned everything up.

Figure 1- Simple distillation apparatus with Hickman-Hinkel Still Head

Figure 2- Fractional distillation apparatus with Teflon spinning band.

Calculation 1- Formula used to calculate the mole fractions of toluene for each of the three fractions taken for both simple and fractional distillation.

mole fractiontoluene=

Area of toluene ( Areaof toluene+ Area of hexane)

Simple DistillationFraction #01: mole fractiontoluene=

( 469155989 ) ( 469155989 + 1736301348 ) ¿ 0.2127250349 ¿ 0.213 → 21.3 %

Fraction #02: mole fractiontoluene=

Fraction #03:

618979674 (618979674+1632285777 ) ¿ 0.2749474407 ¿ 0.275 →27.5 %

mole fractiontoluene=

1717664804 ( 1717664804 +812086533) ¿ 0.6789856295 ¿ 0.679 → 67.9 %

Fractional distillationmole fractiontoluene=

Area of toluene ( Areaof toluene+ Area of hexane)

Fraction #01: mole fractionof toluene=

390024183 (390024183 + 1962259800 ) ¿ 0.1658065877 ¿ 0.166 →16.6 %

Fraction #02: mole fractionof toluene=

767954180 (767954180 + 1732579309 ) ¿ 0.3071161348 ¿ 0.307 →30.7 %

Fraction #03: mole fractionof toluene=

2729392460 (2729392460+ 92037806 ) ¿ 0.9673790251 ¿ 0.967 → 96.7 %

Results and discussion: This experiment helped us learn more about separating a mixture of liquids based on their boiling points. Looking at Graph 1 and Graph 2 we can see the difference between using simple vs. fractional distillation. We see that the number of theoretical plates is greater for fractional distillation as opposed to simple distillation. Prior to doing the experiment we had hypothesized this based on the fact that we were adding a condenser for the fractional distillation. In order to prove this hypothesis, we then performed both types of distillation and took three fractions of each which we ran in the GC. The GC gave us different area values for both hexane and toluene which we then used to calculate the mole fraction of toluene for each fraction taken from both simple and fractional distillation. These mole fraction values for toluene were then used to produce these graphs from which the theoretical plates were calculated. We knew that the more theoretical plates, the better the separation therefore we know that fractional distillation did a better job at separating hexane and toluene in the mixture. During distillation, the mixture of liquids began to boil and since hexane has lower boiling point than toluene, hexane began to evaporate and condensed first. This means that in the first fraction we took, there was more

hexane found than toluene. By the last fraction taken, there was a lot more toluene than hexane because we were reaching the boiling point of toluene. The graphs are in relation to mole fraction of toluene because it goes from 0-100%, meaning that at the lowest temperature we should have 100% hexane and 0% toluene. As the temperature increased and we reached the boiling point of toluene, the closer we got to there being 100% toluene and 0% hexane. The boiling points for each fraction were important because these were the points used in order to make the graphs. We were able to correspond the temperatures with the areas given by the GC to produce the graphs which consequently lead to determining the number of theoretical plates. This was the mathematical approach we used to further prove that fractional distillation is better than simple distillation when trying to separate two liquids in a mixture.

Trial # Temp. of sand ( C) Temp of vapor ( C) Temp. at removal ( C) 1 156.8 68.9-72.7 68.9 2 163.0 83.0-84.4 83.3 3 171.2 104.0-105.2 104.2 Table 1- Shows the different temperatures during simple distillation.

Trial # Temp. of sand ( C) Temp of vapor ( C) Temp. at removal ( C) 1 111.0 73.5-77.0 73.5 2 128.0 83.7-92.9 83.7 3 105.0-106.5 105.0 Table 2- Shows the different temperatures during fractional distillation.

115.00

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105.00

100.00

Boiling point (  C)

95.00

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85.00

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75.00

70.00

65.00 0.00

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Percent Toluene

Graph 1- Phase diagram used to calculate theoretical plates via simple distillation.

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115.00

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105.00

100.00

Boiling point (  C)

95.00

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75.00

70.00

65.00 0.00

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50.00

Percent toluene Graph 2- Phase diagram used to calculate theoretical plates via fractional distillation.

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100.00

Conclusion: If this experiment could be redone I think we would opt out of using a boiling stone and the teflon spin vane for the simple distillation. We noticed that the spinning band was not able to spin as well with the boiling stone in it. We were initially told to leave it in there which is what we did but that caused the spinning band to get stuck to some degree. I think we would also figure out a way to help prevent the boiling to go upwards into the bottom collection area of the Hickman-Hinkel still head. We did have the aluminum foil there to help shield it from over heating but maybe try some more different ways of applying the foil to help prevent this issue since it happened for both simple and fractional distillation. References: Chem 12A Lab Manual; Dr. Chi Hwang, Dr. Jose Valentin. 2017. Clark, Jim. Raoult's Law and Ideal Mixtures of Liquids, Feb. 2014, www.chemguide.co.uk/physical/phaseeqia/idealpd.html. Ian. “Difference Between.” Difference Between Similar Terms and Objects, 29 Aug. 2017, www.differencebetween.net/science/difference-between-fractional-and-simple-distillation/. Questions: 1. The placement of the thermometer has a major influence on the efficiency of measuring the boiling points of the collected mixtures. Explain. -The thermometer has to placed right where the side arm is on the Hickman-Hinkel still head because this is the part of the apparatus in which the evaporated liquid condenses and is collected for the fractions to be taken. If the thermometer is placed too low or too high, the temperature is inaccurate because it would be the temperature before it begins to condensate. 2. What is the purpose of using a boiling stone during heating? -The purpose of adding a boiling stone during heating is to increase surface area. They can also help in preventing superheating of the liquid that is being distilled while giving a more controlled boil. Also, to help eliminate the possibility that the liquid in the distillation flask will bump into the condenser. 3. What is the effect of the following on the separation of the mixture? a. A soluble, nonvolatile impurity b. A decrease in barometric pressure- If the barometric pressure decreased then that would give a false boiling point for the liquids being separated making it challenging to identify the liquids being distilled. c. An increase in barometric pressure- If the barometric pressure increased. It would be similar to it decreasing. It would give a false boiling point for the liquids being separated. d. Rapid heating rate- if the heating rate is too fast or too slow the separation of the liquids will not occur at the actual temperatures they are supposed to happen.

4. An ideal solution contains 40 mole percent of A and 60 mole percent of B at 80  C. The vapor pressure of pure A and B at this temperature are 480 torr and 890 torr, respectively. Calculate the vapor pressure of the solution. PT = PA+PB = 480 torr+ 890 torr= 1370 torr...