Exp. 7:  Hydrogen Bonding, Solubility, and Polarity - An Approach to Bond Classification

It has been found that the energy required for the vaporization of a liquid is an indication of the nature of the forces of attraction between the molecules of the liquid. If these forces are strong, as evidenced by vaporization energies, a given liquid can be considered to be polar. A relative degree of polarity could be established for a group of liquids on the basis of their heats of vaporization.

These attractive forces of a liquid are also an important factor in the solution process. If the attractive forces of the solute are compatible with those of the solvent, solution results. In order for this to occur, it can be assumed that the solute-solvent attractive forces must be greater than the solute-solute and the solvent-solvent attractive forces. By determining the solubility of an unknown substance in solvents of known polarity, it should be possible to learn something of the polar characteristics of the unknown.

Hydrogen bonding exists in many molecules, and accounts for deviations from the expected, such as in higher melting or boiling points for those liquids. It also can occur between molecules of different substances mixed together. By mixing polar and nonpolar liquids in solvents of different polarity one can compare the degree to which hydrogen bonding affects solubility.

Procedure:

Part A.

Put about half a test tube amount of each of the three solvents in separate test tubes. Repeat two more times until you have a total of nine test tubes (3 with each solvent). You will be given three unknowns. Be sure to record their identifying letter. Then add 10 drops of liquid unknowns or very small amount of solid unknown into each solvent. Stir each for at least 5 minutes or untill the solute has completely dissolved. Discard the three solvents in the appropriate waste containers.

In addition to your observations concerning the solubilities of each solute, use the following scale to indicate solubilities and place the information in the data chart in the spreadsheet on the computer at the workstation. Save when you have finished entering your data.

0 ="insoluble"; 1 = "partially soluble"; 2 = "soluble".

When all the data has been entered and is agreed upon we will print a good copy and then run off copies for each person. Trim and place in Procedure/Data/Observations section.

Part B.

Prepare three medium-sized test tubes as follows:

• FIRST pour 5 mL of distilled water into the 1st test tube, 5 mL of ethylene glycol into the 2nd, and 5 mL of glycerol into the 3rd test tube. Pour 5 mL of C₆H₁₂ into each test tube and note whether C₆H₁₂ forms the top or bottom phase. Watch carefully as you pour the C₆H₁₂ into the test tube. Measure the height of each layer. Note the viscosity of this and the other liquids.

• SECOND, stopper and shake each test tube. Two phases will settle out after shaking. Measure the depth of each phase.

• THIRD, without shaking, add 5 mL ethanol to each test tube. Note whether ethanol enters top or bottom phase. Stopper, and shake the three liquids together. Note the number of phases formed after shaking, and measure the depth of each phase. Discard in appropriate containers.

Questions:

Part A.

1. Using class data as well as your own personal observations, arrange the solutes in order of increasing polarity.

2. Draw Lewis structures for: d, e, f, h, j, k, l which are, in that order, 1-butanol, acetone, cyclohexane, potassium chloride, 2-propanol, methylene chloride, and amyl alcohol.

3. Potassium chloride, (h), is more soluble than cyclohexane (f) in water, while the reverse is true regarding their solubilities in hexane. Calculate the percent ionic character of the bonds in each substance and from their structures (which you chould also draw) explain the solubilities in these solvents.

4. Draw the Lewis structures for water, ethanol, propanol, acetone, and hexane. Discuss as fully as you can the polarity (or lack of polarity) of each and the physical implications of that polarity.

Part B.

1. Why are the viscosities of the four liquids so different? Draw the Lewis structures for ethylene glycol and glycerol (glycerine). Compare the structures with those of water and cyclohexane in your explainations.

2. Why is C₆H₁₂ immiscible with water, with ethylene glycol, and with glycerol?

3. In all three test tubes, the third liquid added is ethanol, and it always enters the C₆H₁₂ phase before shaking. Account for any change in this arrangement after shaking.

4. Suppose you add 5 mL of a third liquid to a test tube containing 5 mL ethanol and 5 mL C₆H₁₂. What information will help you predict whether the liquid will be more soluble in the ethanol or in the C₆H₁₂?

Part B is based on an experiment in LABS (Laboratory Assessment Builds Success) published by the Lawrence School of Science, University of California, Berkley.