Distillation and Boiling Point - Engineering Assignment Help

Assignment Task

Learning Outcomes:

Upon completion of this activity, students will be able to describe the technique of distillation to separate two volatile organic liquids and how to isolate a natural oil.

Introduction:

This activity introduces new techniques used for purification of liquids: simple, fractional, and steam distillation.

Two laws, Dalton’s and Raoult’s Law, that you have encountered in previous chemistry courses are critical to understanding the phenomenon of distillation. You should review these laws as you read more about these types of distillation in your technique manual (pages 143-159).

Liquids can be separated by distillation, a technique of vaporization and condensation. There are several different distillation techniques (simple, fractional, steam, azeotropic, and vacuum distillation). In this activity, we will be exploring the first three types of distillation. You will be watching videos which show how to assemble the glassware for the different distillation techniques and observe the process of simple, fractional, and steam distillation. You will be asked questions about the videos to gain insight into the different distillation processes.

You will then be given actual data from previous students experiments at Miramar College and asked to interpret and draw conclusions from that data. The experiment was performed over three lab periods. In Parts A and B of the experiment, students analyzed the efficiency of two distillation techniques, simple and fractional distillation, at separating a mixture of volatile organic liquids with relatively close boiling points: Cyclohexane (81 oC) and toluene (111 oC). The fractional distillation was performed first followed by duplicating the fraction volumes collected during the simple distillation. The efficiency of separation was analyzed by GCMS (Gas Chromatography – Mass Spectrometry) on collected samples from both distillations.

While no mixture is perfectly ideal, mixtures like cyclohexane and toluene get close to ideal behavior. They are similarly sized non-polar molecules and so have similar van der Waals attractions between them. However, they aren't identical - and so although they get close to being ideal, they aren't actually ideal. For the purposes of this topic, getting close to ideal is good enough!

In the graph above, the vapor pressure of two miscible liquids are plotted against the mole fraction of each component in the mixture. A higher vapor pressure results in a lower boiling point, so B is the component with the lower boiling point (in our case cyclohexane) and A is the component with the higher boiling point (in our case toluene).

The actual vapor pressure of a mixture of miscible liquids differs from the ideal vapor pressure because the intermolecular forces holding A to B are usually weaker than A to A or B to B. The weaker IMF leads to a higher vapor pressure. Notice that the highest vapor pressure anywhere is the vapor pressure of pure B. Cases like this, where the deviation is small, behave just like ideal mixtures as far as distillation is concerned, and we don't need to say anything more about them.

Raoult's Law: (applies to miscible liquid mixtures) The partial vapor pressure of a component in a mixture is equal to the vapor pressure of the pure component at that temperature multiplied by its mole fraction in the mixture (ideal case).

PA = XA x P°A

PB = XB x P°B

In this equation, PA and PB are the partial vapor pressures of the components A and B, XA and XB are the mole fractions of A and B (the fraction of the total number of moles present), P°A and P°B are the vapor pressure of pure A and B.

In any mixture of gases, each gas exerts its own pressure. This is called its partial pressure and is independent of the other gases present. Even if you took all the other gases away, the remaining gas would still be exerting its own partial pressure.

The total vapor pressure of the mixture is equal to the sum of the individual partial pressures, called Dalton’s Law of Partial Pressures.

P Total = PA + PB

Let’s convert the vapor pressure graph to a boiling point composition diagram. We'll start with the boiling points of pure A and B. B has the higher vapor pressure. That means that it will have the lower boiling point.

droop below this. That happens with certain non-ideal mixtures and has consequences that are not explored at this time. To make this diagram really useful we are going to add another line. This second line will show the composition of the vapor over the top of any particular boiling liquid or liquid mixture.

The diagram above shows what happens if you boil a particular mixture of A and B. Notice that the vapor over the top of the boiling liquid has a composition which is much richer in B - the more volatile component. If you repeat this exercise with liquid mixtures of lots of different compositions, you can plot a second curve - a vapor composition line.

Important: Once again, take great care drawing this second curve. No point on the curve must be higher than the boiling temperature of the pure A. If you boil a liquid mixture, you can find out the temperature at which it boils, and the composition of the vapor over the boiling liquid. For example, if you boil a liquid mixture C1, it will boil at a temperature T1 and the vapor over the top of the boiling liquid will have the composition C2.

Notice again that the vapor is much richer in the more volatile component B than the original liquid mixture was. Suppose that you collected and condensed the vapor over the top of the boiling liquid and reboiled it. You would now be boiling a new liquid which had a composition C2. That would boil at a new temperature T2, and the vapor over the top of it would have a composition C3. You can see that we now have a vapor which is getting quite close to being pure B. If you keep on doing this (condensing the vapor, and then re-boiling the liquid produced) you will eventually get pure B. This is the basis for fractional distillation. However, doing it like this would be incredibly tedious. Real fractionating columns (whether in the lab or in industry) automate this condensing and re-boiling process. Boiling stones are used in a fractionating column. The vapor condenses on the stones. The condensed vapor is enriched in the lower boiling component. 

Activity

1. Label the following diagram as a simple or fractional distillation and label each piece of the apparatus with the following terms: water in, water out, thermometer, condenser, receiving flask, distillation flask, distillation adaptor, fractionating column.

As a group watch the youtube video below. Discuss and come up with reasonable answers to the questions below.

2. For simple and fractional distillations should the liquids you are attempting to separate be miscible or immiscible?

3. What are the criteria for deciding whether to use simple versus fractional distillation?

4. How full should your distillation flask be when attempting a distillation?

5. What happens if your distillation flask is not full enough? 6. What if it is too full?

7. Correlate heat transfer to heating mantle size relative to distillation flask size.

8. Does the water flow from top to bottom through the condenser jacket or bottom to top? Why?

9. How fast should the water flow through the condenser? Why?

10. How high should your variac be set? What is meant by a gentle boil?

11. What level should the thermometer tip be set to?

12. What rate of drops per minute should you shoot for during a simple distillation?

13. What can happen if the rate is too fast?

14. What if the rate is too slow?

15. What are the 5 general guidelines when performing a simple distillation? Discuss.

16. When should you use fractional distillation?

17. What is the main piece of glassware used in a fractional distillation that is not used in a simple distillation?

18. How many drops per second shoot you shoot for in a fractional distillation?

19. Why might you need to insulate your fractionating column?

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