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the microscale laboratory Microscale Distillation-Calculations and Comparisons Marietta H. Schwartz University of Massachusetts at Boston 100 Morrissey Bbd. Boston, MA 02125
--nd.Ny13045
Experimental Vapor-Liquid Equilibrium Data
for 2-Propanol/MethanolSystem at 760 mm Mercurya
Temperature ('C)
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SUNY-Coniand
Mole Percent of Methanol
Most laboratorv texts present the theorv of distillation but usually mclude experiments that focus mainly on the performance ofthe technique with minimal analysis ofthe results. The standard volime versus temperature plot is used widely, but this does little to demonstrate the efficiency of a particular distillation setup to the student; the major thrust of the experiment tends to be learning to asA distillation experiment cgn have semble the marc of an impact on students if they perform some of the calculations (such .~~ as determinine thenumber of theoretical plates in a given setup), thereby proving to themselves that, for example, a fractional distillation is more efficient than a simple distillation. Microscale distillations can be performed two ways, using either the traditional distillation setup with micmscale elassware or some version of the Hickman still. Most of thelaboratory texts available use variations of the traditional elassware. ( 1 3 )occasionallv mentionine the Hickman stzl as an option. Others pref& some versLn of the Hickman still (41, and some present the theory of distillation without a dedicated experiment (5).We prefer the traditional setup,' because it prepares students to work on a larger scale if necessary. It is similar to distillations performed in research laboratories and provides an easier method for collection of samples for analysis. In the distillation experiment currently used in our organic teaching laboratory, each student performs both a simple and a fractional distillation using the micmscale version of traditional glassware. The distillate is analyzed using refractometry (although gas chromatography would be equally useful) and the results are used to calculate the number oftheoretical plates in each setup.2Half ofthe students use an air condenser as a fractionating column and half use a condenser packed with steel wool. The simple ~
ARDENP. ZIP;
Vapor
66.22 95.35 67.94 89.10 70.22 80.00 72.67 68.50 74.76 57.00 77.06 42.85 78.94 29.60 13.20 81 .OO 'Derived from the data given in ref. 6.
Liquid 90.10 79.00 66.05 52.20 40.80 29.30 19.50 8.10
distillation typically gives an average of 1.5 theoretical plates (instructor average 2.01, the fractional distillation with an air condenser an average of 2.0 theoretical plates (instructor average 3.01, and the fractional distillation with a packed fractionating column an average of 3.5 theoretical plates (instructor average 4.0). Interestingly, the spinning band Hickman-Hinckle still (tested multiple times by two instructors) did not perform significantlybetter than the packed fractionating column for this particular distillation, giving an average of 3.0 theoretical plates for this solvent system. Details of this distillation experiment are provided in the Experimental section. Brief Experimental Procedure A mixture of 1.5 mL of methanol and 6 mL of isopmpano13 is used for each distillation. A small sample of the initial mixture is retained for analysis, and the rest is subjected to either simple or fractional distillation, with the first distillate also being retained for analysis. The refractive index of each sample as well as that of pure methanol and isopropanol is obtained and data analysis is performed as shown below. Data Analysis Analysis by Refractive Index
Most glassware companies are quite willing to modify their microscale glassware kits to suit individual preferences. The difference in cost between a kit containing the Hickman distillation glassware and the microscale version of traditional glassware is negligible. A theoretical plate corresponds to a vaporization-mndensation cycle within the distillation apparatus. The more of these cycles that occur during the course of a distillation,the more efficient the separation of the two components will be. Equal volumes of the two alcohols are not used as the proportions oiven convert to ao~roximatelv30 mole Dercent methanol. a laroe enouoh to fiasible but small enouoh zo ..-fraction -~- iake the - distillations ~ they are no1 trivia. hole in the sample calc~latonthat tne 1ract;onal d~stlllallonImproves Ins to approximately 90 mole percent metnanol. The major ass~mpfion here 1s mat the voldmes of the two liq~os are additive. This is not always true, but in many systems (including that in this experiment)it is close enough within the limits of this analysis. ~
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Assuming that the refractive index is a linear function of the com~ositionof the s a m ~ l ethe . ~ volume oercentaees of experimkntal samples can Le calculated using eq 1,where VA and Vg represent the volumes of methanol and isopropanol, respectively, and Vt is the total volume.
The same information can be obtained more easily from a graph of refractive index versus the volume percentage of one of the components, prepared by plotting the refractive indexes of pure A (methanol) and pure B (isopropanol) and connecting them with a straight line as shown in Figure 1.The volume percent composition of a sample mixture can be obtained from its measured refractive index using Volume 69 Number 4 April 1992
A127
the microscale laboratory
'!Ki ..
g
5 -
1
a? -
vopor
Liquid
50 40
I 30
stort
20 10 0 64
66
68
70
72 74
76 78 8 0 82 84
Temperature (degrees C)
Figure 1. Conversion graph to translate refractiveindexes intovolume percent methanol, with the data for a simple distillation indicated on the graph.
Figure 2. Vapor-liquid equilibrium diagram generated from the data in the table, with the theoretical plate markings for a simple indication indicated on the diagram.
this graph. This can be converted to mole percent composition using the appropriate conversion factors as described below.
show that this simple distillation contained approximately two theoretical plates. A fractional distillation using approximately the same amounts of methanol and isopropanol gave refractive indexes of 1.3665 for the starting mixture and 1.3345 for the first distillate. This converts, using the calculations described above, into 33.6 mole percent and 90.9 mole percent of methanol, respectively Figure 2 shows that this fractional distillation contained approximately four theoretical plates.
Theoretical Plate Calculation To calculate the number of theoretical plates for the distillation, the starting and ending mole percent compositions are determined and marked on a vapor-liquid equilibrium graph as shown in Figure 2. Then, starting from the beginning composition, straight lines representing vaporization-condensation cycles are drawn until the ending composition is reached. The general procedure for the calculation is as follows. First, the refractive index of the sample is measured. Next, a graph such as that shown in Figure 1 is used to convert refractive index to volume percent. Volume percent is converted to mole percent, assuming a total volume of 100 mL and using the density and molecular weight of isopropanol and methanol. Finally, the mole percent numbers and the vapor-liquid equilibrium graph are used to count the number of theoretical plates. For example, the samples obtained from the simple distillation described here gave refractive indexes of 1.3669 for the initial mixture and of 1.3511 for the first distillate. Figure 1 was used to convert these to 19.4 vol % methanol for the initial mixture and 52.1 vol % methanolfor the first distillate. The volume percentages were then converted to mole percentages as shown below (assume 100 mL total volume). 19.4 mL MeOH x 0.791 g/mL x 1 moW32.04g = 0.479 mol MeOH 80.6 mLi-PrOH x 0.785 g/mL x 1 moW60.10 g = 1.053 mol i-PrOH
Mole percent methanol = moles methanol divided by total moles = 31.3 mole percent for the starting mixture. The same calculations for the fust distillate of the simple distillation show that 52.1 volume percent methanol becomes 67.3 mole percent methanol. This gives the starting and ending points for the theoretical plate calculation. These numbers, when plotted on Figure 2 and connected by the "stair steps" of vaporization-condensation cycles, A128
Journal of Chemical Education
Acknowledgment
The roots of this experiment lie in the in-house laboratory manual written by Paul Schatz at the University of Wisconsin-Madison. The author would like to thank Ken Cemy for assisting in testing the various distillation setups and Robert Carter for helpful discussions. Literature Cited 1. Wileor, C. P. Jr.Erpeimenla1 OgonY Ckmiahy (ASmoU-SenbAppmhJ; MacMillan hbliahing Company: New York 1988; Chapter 2. 2. Wfiliamsrm. K L M m s c n l o and Micmsmlo Orzonic Exmtimnts: D.C. Heath and
4. Mayo, D.w.: Pik%R.M :Butcher, S. S. Mi~mpml.0gonrcInbormory;John Wi14. & Sona:New York, 1989: pp 54-69. 5. Pavia, D. L;Lmpman, G M.; Krir, G.S.; Engel, R. G.Intnxlufion lo OgonicLabomlorv Tkhnwues-A MicmscoleA.~.m o e hSaundem : Colleec Publishine: Chiearn. 19s0; ~eehniduea8 and 10. 6. Bdlard, L. H.:Van Winkle, M. I n d U 8 f d ?zndEngimtiwCkmisf?y 1952,44,2451.
The Synthesis of Urea: An Undergraduate Laboratory Experiment Stephanie Tanski, Janeen Petro, and David W.
Cleveland State University Cleveland, OH 44115
all'
In this paper we present a laboratory experiment for the production of urea from silver cyanate and ammonium chloride. The chemical equation for this synthesis is
