Experiment 14 Intermolecular Forces rev 1/12 GAL: We will examine connections between molecular structure, intermolecular forces, and physical properties. BAKGRUND: Physical properties such as solubility, melting point, and boiling point are determined by a substance s intermolecular forces. Non-polar molecules have the weak intermolecular forces known as London dispersion forces caused by slight, temporary asymmetries in their electron clouds. Although London dispersion forces are relatively weak, as the number of electrons in a molecule increases so does the strength of the London dispersion forces. So, if we are comparing several non-polar molecules to each other, we expect the largest ones (with the most electrons) to have strongest London dispersion forces. Polar molecules have London dispersion forces in addition to dipole forces. Dipole forces are attractions between + and - charges on the molecules. Dipole forces are generally stronger than London dispersion forces because the charges are permanent, not temporary, and the charges are stronger. Both dipole forces and London dispersion forces figure into the total strength of intermolecular forces for polar compounds. Suppose we had two molecules with very similar London dispersion forces but one was polar and one was not. The presence of the dipole forces will give the polar molecule stronger intermolecular forces overall. - + ydrogen bonds are the strongest intermolecular forces found in covalent compounds. They are formed by the attraction between a hydrogen atom on one molecule and a nonbonding pair of electrons on a second molecule. ydrogen bonds are found only in molecules where a hydrogen atom is bonded to fluorine, oxygen, or nitrogen. These molecules also have dipole and London dispersion forces, and all three types of forces contribute to the total strength of intermolecular forces. The presence of hydrogen bonding makes the total intermolecular forces very strong and leads to high melting and boiling points. aving multiple groups capable of hydrogen bonding makes melting points even higher. Two molecules with exactly the same types of intermolecular forces may still differ from each other in the strength and impact of these forces due to general molecule shape or the placement of hydrogen bonding groups which are very particular about geometry. Linear and planar molecules often exhibit stronger intermolecular forces because their shape allows the molecules to pack closely and maximize the interactions. The combined effects of all intermolecular forces determine physical properties. The stronger the intermolecular forces a compound has, the more energy will be required to overcome those attractions. Thus, strong intermolecular forces correspond to high boiling points and less volatile compounds, i.e., ones that will evaporate less. In Part 1, you will observe several liquids as they evaporate. All liquids 2012 Intermolecular Forces 15-1
cool as they evaporate because they must use energy to overcome intermolecular forces and move into the gas phase. The liquids that show the most cooling have lost the most molecules due to evaporation. This indicates a compound with weaker intermolecular forces. In Part 2, you will measure the melting points of several compounds. In Part 3, you ll also use Spartan Student to calculate their dipole moments (polarities). In Part 4, you will use physical models to look at intermolecular forces in water and aqueous solutions. Table 1: Names and Structures of rganic ompounds Used in this Experiment Name Structure Name Structure methanol ethanol 1-propanol 1-butanol pentane hexane benzoic acid 2-hydroxybenzoic acid or o-salicylic acid 4-hydroxybenzoic acid or p-salicylic acid naphthalene 15-2 Intermolecular Forces 2012
PRE-LAB ASSIGNMENT: Two of the molecules you will study this week are methanol and ethanol. The structures of these molecules are given on page 15-2. opy the structures in your notebook. List all the types of intermolecular forces these molecules have. AZARDS: Several of the liquids that you will be working with are highly flammable, so no flames are allowed in the lab. While none of the chemicals used in this experiment pose unusual health hazards, your hands may be exposed to these chemicals. Wear disposable gloves to limit your exposure. Keep the test tubes containing the liquids closed when not in use. Part 1 is done in the hood. Look up the MSDS for methanol (http://hazard.com/msds/index.php, or Trexler 464). Record the following information in your notebook hazards section: appearance, odor, and boiling point (see Physical & hemical Properties) a summary of potential health effects (see azards Identification) a summary of potential effects if released to the environment (see Ecological Information) LABRATRY DATA AND BSERVATINS: The in-lab portion of this experiment will be done in teams of two students, but the lab reports will be done separately. This means each student will need a full set of notebook entries. Note the name of your in-lab partner in your lab notebook. Remember to record both what you do and what you observe in your notebook. The four parts of this experiment may be done in any order. Your lab instructor will assign you to start with a given part in order to avoid congestion. PART 1 PREDURE: Evaporation 1. Your numerical data for Part 1 will be most easily recorded in a table with four columns labeled: liquid, maximum temperature, minimum temperature, and. You will need one line for each of the six liquids tested. 2. Use one of the computers that are set up in a hood. If it has not already been done for you, prepare the computer for data collection by opening Experiment 13 from the Roanoke Experiments folder in Logger Pro. n the Graph window, the vertical axis has temperature scaled from 5 to 30 o. The horizontal axis has time scaled from 0 to 250 seconds. 3. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by wire. Roll the filter paper around the probe tip in the shape of a cylinder. The paper should be even with the probe end. 4. You will be testing 6 liquids. The four alcohols are methanol, ethanol, 1-propanol, and 1-butanol. The two alkanes are pentane and hexane. These have been pre-measured into test tubes. You may do the liquids in any order. Get two liquids at a time. When finished, return the liquids to the supply bench so that others may use them. Keep the liquids stoppered when they are not in use. 2012 Intermolecular Forces 15-3
5. Wear gloves for the rest of Part 1. Stand Probe 1 in the first liquid container and Probe 2 in the second liquid container. Record which probe goes into which liquid. Make sure the containers do not tip over. 6. Prepare 2 pieces of masking tape, each about 10-cm long, to be used to tape the probes in position during Step 7. 7. After the probes have been in the liquids for at least 30 seconds, begin data collection by clicking. Monitor the temperature for 15 seconds to establish the initial (maximum) temperature of each liquid. Then simultaneously remove the probes from the liquids and tape them to the metal flashing at the front edge of the hood so that the probe tips extend 5 cm over the edge and into the hood. The probe tip should not touch the working surface of the hood. Pull the hood sash down as far as you can. Stopper the test tubes to minimize the loss of these liquids. 8. When both temperatures have reached minimums and have begun to increase, click to end data collection. Two statistics boxes should appear on the screen. You may need to drag the boxes apart since the computer sometimes places them on top of each other. Record the maximum and minimum values for Temperature Probe 1 (first liquid) and Temperature Probe 2 (second liquid) in your notebook. 9. For each liquid, subtract the minimum temperature from the maximum temperature to determine T, the temperature change during evaporation. Record this in your table. 10. Roll the rubber band up the probe shaft and dispose of the filter paper in the waste container in the hood. 11. Return your first two liquids to the supply bench. Be sure they are stoppered. Get two different liquids. When you choose ollect for your second set of liquids, you will be prompted to discard your prior set of data. Discard it. Repeat the procedure above first for liquids 3 & 4, and then for liquids 5 & 6. 12. Be sure that both you and your lab partner have a full set of data before leaving lab. Remember that you will write up your reports separately. PART 2 PREDURE: Melting points of organic solids 1. ne member of your team should prepare melting point capillaries of benzoic acid and naphthalene. The other team member should prepare melting point capillaries of 2- hydroxybenzoic acid (or o-salicylic acid) and 4-hydroxybenzoic acid (or p-salicylic acid). Remember that you need just one crystal of each in your capillary. Set each capillary on a labeled piece of paper so that you don t mix them up. 2. Measure the melting point range of each solid, starting with the temperature where you first see softening to the temperature where the entire sample has melted. Put both capillaries in the MelTemp at the same time. Make sure you know which is which. You will heat rapidly until you near the melting point for each compound and then slow the rate of heating to just 1 o every 10-15 seconds. 15-4 Intermolecular Forces 2012
3. If you are doing the benzoic acid/naphthalene set, heat quickly to 70 o then adjust so that temperature increases just 1 o every 10-15 seconds until your first solid melts. Now heat quickly to 110 o and then again adjust so that temperature increases just 1 o every 10-15 seconds until your second solid melts. For each solid, you should record a melting point range of just a degree or two in which the solid softened and melted. 4. If you are doing the 2-hydroxybenzoic acid/4-hydroxybenzoic set, heat quickly to 145 o then adjust so that temperature increases just 1 o every 10-15 seconds until your first solid melts. Now heat quickly to 205 o and then again adjust so that temperature increases just 1 o every 10-15 seconds until your second solid melts. For each solid, you should record a melting point range of just a degree or two in which the solid softened and melted. 5. As soon as your second solid melts, get the apparatus cooling back down using the compressed air. Dispose of the capillaries in the broken glass box. PART 3 PREDURE: Using Spartan Student to calculate dipole moments of organic solids 1. Use Spartan Student to draw models of the four molecules from Part 2: benzoic acid, 2- hydroxybenzoic acid, 4-hydroxybenzoic acid, and naphthalene. Follow the structures which are shown on page 2. ave these drawings handy to follow. Recall that you used Spartan Student for Exp 11 in EM 111. Begin by opening Spartan Student on one of the lab computers. 2. pen a new file by clicking the button. n the righthand side of the screen, you should see a collection of molecular fragments that you can use to build the molecules. 3. Start by building benzoic acid. Add a benzene ring to the main drawing area. Now click on the Groups dropdown menu. hoose arboxylic Acid. Add this to an open bond in your drawing. 4. Spartan should recognize your drawing as benzoic acid. Look at the bottom right of the screen. lick the up arrow next to the name. hoose Replace. 5. Now click View: lick and drag to rotate the molecule. 6. From the dropdown menu at the top of the screen, choose Display, Properties. Record the displayed dipole moment in your notebook. 7. You can easily modify this benzoic acid structure into 2-hydroxybenzoic acid, also known as salicylic acid. lick Add Fragment: 8. hoose the molecular fragment at right that is an oxygen atom with two single bonds. Add this to the appropriate place on your molecule. Spartan should recognize your molecule. Follow steps 4-6 above to find your dipole moment. 9. Again modify your structure, this time to form 4-hydroxybenzoic acid. lick on Add Fragment as in Step 7. Now delete the group from your molecule using the Delete Button: Switch back to Add Fragment, and add the new group on the correct spot. 2012 Intermolecular Forces 15-5
10. Again, Spartan should recognize your molecule as p-salicylic acid. Follow steps similar to Steps 4-6 above to find the dipole moment. 11. The last molecule you need to make is naphthalene. It is sufficiently different from the others that you should close this file and open a new one using the buttons across the top of the screen. You do not need to save a copy of the current file. 12. nce you have a blank drawing area, choose the Rings dropdown menu and select Naphthalene. You do not need to modify this structure! Simply follow Steps 4-6 above to record the dipole moment. 13. nce you have dipole moments for your four compounds, simply use the lose File button but leave the software open for the next group of students. PART 4 PREDURE: Intermolecular forces in water and aqueous solutions 1. Spread out your up of Water models. You should have 12 2 molecules, one Na 1+ ion, one l 1- ion, one hydroxyl group, and one 2 6 (ethane) molecule. Small magnets in the models will help us simulate intermolecular forces. Record observations in your notebook as you follow the instructions below. 2. Select two 2 molecules. ow are they most attracted to each other: to, to, to, or some combination of these? Which two types of intermolecular forces are being represented here? Record your observations. 3. hoose one 2 molecule to be your central molecule. ow many other water molecules can you attach directly to this one central molecule? Draw a structure similar to the one of water molecules on page 13-1 but showing how all the molecules attach to the central water molecule. 4. In solid 2 (ice), the water molecules arrange themselves so that each water molecule can hydrogen bond with as many others as possible. Arrange all 12 of your water molecules in the most compact structure that also maximizes hydrogen bonding. Use a sketch and written description to record this structure in your notebook. Describe how much open space is trapped inside your solid structure. 5. In liquid 2, the water molecules are constantly forming and breaking intermolecular hydrogen bonds as the individual molecules move. Take your model of solid ice in your hands and compact it like you would a snow ball to simulate this. When compared to your solid structure before, how much open space is now trapped inside the liquid structure? 6. Now let s consider how solutes interact with water molecules. Find the Nal. The smaller ion is Na 1+. Note the strength of its attraction for a l 1- ion. ow many l 1- could fit around a single Na 1+? 7. Examine the Nal model on the instructor s bench. Use a sketch and written description to record this structure in your notebook. 15-6 Intermolecular Forces 2012
8. ow are water molecules attracted to the Na 1+ : through the or through the? Fit the maximum number of water molecules on the Na 1+ possible. Use a sketch and written description to record this structure in your notebook. 9. ow are water molecules attracted to the l 1- : through the or through the? Fit the maximum number of water molecules on the l 1- possible. Use a sketch and written description to record this structure in your notebook. 10. You should now have a Na 1+ surrounded by water molecules and a l 1- surrounded by water molecules. This is how they actually exit in water. We say that each ion is surrounded by a hydration sphere. old the hydrated Na 1+ next to the hydrated l 1-. With the hydration spheres intact, will the ions stick together? 11. Pull the water off the ions and examine the attraction between the water molecules and the ethane ( 2 6 ) molecule. ow strong is its attraction for water? Does it form a hydration sphere? 12. ne of the hydrogen atoms on the ethane molecule is marked with a small colored spot and small raised dots on the carbon atom at its base. arefully remove this hydrogen and set it aside in a safe place. Insert the, hydroxyl group, in its place. You now have 2 5, ethanol. ow strong is its attraction for water? Use a sketch and written description to record the structure/interaction between ethanol and water. 13. Remove the from your ethanol and replace the that you previously set aside. arefully count the models as you return them to the original cup. Be sure that you have 12 2 molecules, one Na 1+ ion, one l 1- ion, and one hydroxyl group, and one 2 6 ethane molecule. RESULTS: For Part 1, prepare a Results table with 6 columns: liquid name, formula, structure, molar mass, intermolecular forces, and T. Structures for the liquids are given in the Introduction. List all the intermolecular forces expected for each compound. The Introduction or your textbook will provide help. ombine Parts 2 and 3, preparing a Results table with 7 columns: compound name, structure, molar mass, experimental melting point, literature melting point, dipole moment, and intermolecular forces. Structures for the compounds are given in the Introduction (draw them in by hand). Experimental melting points are the ones you found in lab. Literature melting points are accepted values that you find in chemical literature. Look up your compounds in either the R (see the Table of Physical Properties of rganic ompounds) or the Merck Index. List the source of your data just below your Results table. In the final column, list all the intermolecular forces expected for each compound. QUESTINS: Questions 1-4 refer to Part 1 only: 1. Which of the alcohols studied in Part 1 has the strongest intermolecular forces of attraction? The weakest intermolecular forces? Explain how the results of this experiment show this. Explain how this result could be predicted from looking at the structures or formulas. 2012 Intermolecular Forces 15-7
2. Which of the alkanes studied in Part 1 has the stronger intermolecular forces of attraction? The weaker intermolecular forces? Explain how the results of this experiment show this. Explain how this result could be predicted from looking at the structures or formulas. 3. Two of the liquids, pentane (an alkane) and 1-butanol (an alcohol), have nearly the same molar masses, but significantly different t values. Explain the difference in t values of these substances, based on their intermolecular forces. 4. Diethyl ether ( 3 2 2 3 ) has a molar mass similar to that of pentane and 1-butanol (see Question 3). Draw the Lewis structure of diethyl ether. Is it polar or non-polar? What types of intermolecular forces are expected in diethyl ether? What t do you predict for diethyl ether? (Think about the t you measured for pentane and 1-butanol) Explain your reasoning. Questions 5-9 refer to Parts 2 and 3 only: 5. ow well do your experimental melting points correspond to the literature melting points? What might cause an experimental melting point to deviate from an accepted value? 6. Rank your four compounds in order of increasing molar mass. Do your melting points correlate well with molar mass (that is, does increasing molar mass always cause a similarly sized increase in melting point)? ite examples from your data to support your statement. 7. Rank your four compounds in order of increasing dipole moment. Do your melting points correlate well with dipole moment (that is, does increasing dipole moment always cause a similarly sized increase in melting point)? ite examples from your data to support your statement. 8. What general connection should exist between melting points and the strength of intermolecular forces? Give a general explanation of why this is true. (Talk about changes that happen when a solid melts.) See the Introduction or your textbook for help. 9. Polar molecules should have a measurable dipole moment. Do the dipole moments that Spartan calculated correspond to your polarity predictions from just looking at the structures? ite examples from your data to support your statement. Questions 10-13 refer to Part 4 only: 10. List all the types of intermolecular forces present in a sample of pure water. Describe each in a sentence or two. Which of these were simulated by our models? 11. Use what you observed about the structures of solid and liquid water to explain why ice floats on liquid water. (int: think about the effect of that open space you noticed) 12. Nal dissolves well in water. Describe what you observed with the models to explain why and how it dissolves. 13. Describe the interactions between water and the two related molecules of ethane and ethanol. Which will be more soluble in water? Why? 15-8 Intermolecular Forces 2012