INTERMOLECULAR FORCES AND THE LIQUID-VAPOR EQUILIBRIUM 1

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Pearl Daniels
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1 Experiment 12Z FV 7/10/17 INTERMLEULAR FRES AND TE LIQUID-VAPR EQUILIBRIUM 1 MATERIALS: 150 ml beaker, 6 ml graduated plastic syringe sealed at the tip, digital thermometer, hot plate, plastic bin, 2 metal washers; organic samples; injector needles PURPSE: The purpose of this experiment is to measure the effect of temperature on the vapor pressure of several liquids. The data will be analyzed to extract values for the heat of vaporization, Δ vap, for each of the various liquids and these will be interpreted in terms of intermolecular forces. LEARNING BJETIVES: By the end of this experiment, the student should be able to demonstrate the following proficiencies: 1. Use a spreadsheet program for data manipulation, graphing, and regression analysis. 2. Describe the effects of changes in temperature on the vapor pressure of a pure substance. 3. Describe how intermolecular forces influence the relative vapor pressure of a pure substance. 4. Understand the use of graphical methods to extract thermodynamic information from experimental pressure and temperature data. 5. Utilize Dalton s Law of Partial Pressures, and the Ideal Gas Law, to relate experimental data to properties of the test substance. DISUSSIN: The molecules of a gas move freely throughout the entire volume of the container, the individual molecules staying widely separated and experiencing little or no interaction with other molecules. Molecules in a liquid, while free to move throughout the volume of the sample, are constrained by intermolecular forces to remain in contact with their neighbors. The strength of such intermolecular forces and the energy of motion available to the sample (based on the temperature), together dictate the physical state of a substance. Evaporation is the process of converting a substance from the liquid phase to the gas phase. It is an endothermic process, since energy is required to overcome the attraction that a liquid molecule feels for its neighbors. The molar enthalpy of vaporization, Δ vap, is the energy required to evaporate one mole of a substance at constant temperature and pressure. (This quantity is often simply called the heat of vaporization.) The magnitude of Δ vap is thus a measure of the strength of the intermolecular forces in a pure substance. The molecules in a liquid will have a distribution of energies at any temperature, as do the molecules of a gas. If a liquid is placed in an evacuated, closed container, some of the molecules of the liquid (those in the higher energy range) will have sufficient energy to escape to the gas phase. Thus the pressure in the container will rise. Some of the gas phase molecules will hit the liquid surface and be unable to escape the attractions for their new neighbors; these (lower energy molecules) have undergone condensation and become part of the liquid. As more molecules accumulate in the gas phase (via evaporation), the rate of condensation will also increase. Eventually, the rate of evaporation and the rate of condensation will become equal, and the pressure in the container will level off at some constant value. The system is said to be in equilibrium, and the pressure of gas that exists over the liquid is called the equilibrium vapor pressure of the liquid. The equilibrium vapor pressure depends on the temperature of the sample (since a higher temperature gives a larger fraction of high-energy molecules), and on the strength of the intermolecular forces holding molecules of the liquid together. So, a comparison of the equilibrium vapor pressures of a number of substances at the same temperature allows one to rank the relative strengths of the intermolecular forces of those substances. For any individual substance the variation of vapor pressure (P) with temperature allows a determination of the enthalpy of vaporization of that substance, as given by the lausius-lapeyron equation (assuming a constant Δ vap): 1 Based on Levinson, G.S., Journal of hemical Education, 59, 337 (1982), and adapted From USNA Exp.12E by MIDN 1/ Jonathan abarrus, 16 & 1/ Michael Brown, 16. E12Z-1

2 ln PP = Δ vvvvvv 1 + (1) RR TT where R is the gas constant (8.314 J/mole K), T is the absolute temperature (in K), and is a constant. As seen in this equation, liquids with a large positive value of Δ vap will have a low equilibrium vapor pressure at any temperature. As the temperature increases, ln P, and thus P, also increases. In this experiment, the volume of a gas mixture of air and sample vapor will be measured at several temperatures and at atmospheric pressure. ne such measurement will be made near the freezing point of water, when the vapor pressure of water or the sample organic compound is nearly zero. This allows a determination of the (constant) number of moles of air trapped in the cylinder and thus the partial pressure of air in the mixture at any temperature. From this and the barometric pressure, the partial pressure of the liquid at any temperature can be determined by their differences. Then the application of Eq. (1) allows the determination of the molar heat of vaporization Δ vap by graphical methods. E12Z-2

PREDURE: (Work in pairs.) 1. btain the barometric pressure and record it in the attached data sheet. Take the small plastic bin from your student drawer and fill it with ice. 2.

arefully cover the top of the syringe with your finger, invert it, and place it in the beaker.

RGANI STEP: See your Instructor for a sample loading syringe containing your assigned organic liquid. (If assigned water as a sample, just continue to step 4.

3 PREDURE: (Work in pairs.) 1. btain the barometric pressure and record it in the attached data sheet. Take the small plastic bin from your student drawer and fill it with ice. 2. Fill a 150 ml beaker with distilled water. Place about 6 ml of distilled water in the prepared plastic syringe. arefully cover the top of the syringe with your finger, invert it, and place it in the beaker. Do not release the syringe until you have placed a metal washer atop the syringe to hold it under the water. The syringe must be completely immersed throughout the experiment. 3. RGANI STEP: See your Instructor for a sample loading syringe containing your assigned organic liquid. (If assigned water as a sample, just continue to step 4.) You or your Instructor should load the sample by carefully placing the opening of the injector tube underneath the opening of the submerged syringe. Slowly inject about 1 ml of organic liquid; it will rise to the top inside the submerged syringe, under the air bubble. 4. Record the identity of your sample liquid and its normal boiling point on the data sheet. (See the table on p. E12Z-6. You should have this value in mind as you work through the procedure.) Place the newly prepared apparatus on the heating mantle and turn the heat dial to the highest setting. 5. Place a ring stand with clamp adjacent to the hot plate, and use it to loosely dangle your thermometer into the beaker, with the tip about mid-height. Bend the wide blade of your metal spatula to a shallow angle (roughly 30 o ). Using the bent spatula, continuously stir the water while it heats. This is done to ensure that the temperature of the water bath is uniform throughout. The bend will help lift and distribute the hotter water from the bottom 6. While continuing to stir, monitor the volume of the gas bubble within the submerged syringe. As the temperature rises, the gas bubble will expand exponentially. Should it begin to expand too quickly, remove the beaker from heat, wait about a minute, and begin heating again. (If any of the gas bubble escapes, you will have to start over - see your Instructor.) ontinue to stir and heat until the gas bubble passes the 4 ml mark on the syringe. Typically, you need to heat the water bath to a temperature ~10 o below the boiling point of your organic liquid, but focus on the size of the bubble, not this temperature guideline. 7. As the volume of the gas bubble passes the 4 ml syringe mark while expanding, remove the beaker from the heat (use ot ands) and allow it to cool slowly. ontinue to stir throughout the cooling process. 8. The bubble will continue to expand slightly after the beaker is removed from the heat source, but as the water bath cools, the gas bubble will begin to contract. nce the gas bubble again reaches the 4.0 ml mark as it contracts, begin to record the bath temperature and bubble volume at every 0.2 ml of volume change. Record this data to the nearest 0.1 o. ontinue to stir throughout the measurement process. Make sure the tip of the thermometer is in the middle of the beaker, not at the bottom, when you measure temperature. It will also be helpful to hold the syringe in a constant position by pressing down on the top of the syringe while stirring. 9. After temperature values have been recorded for the contracting gas bubble for volumes between 4.0 and 2.0 ml, collect two additional data points between 10 and 20 o by adding small amounts of ice to the beaker and stirring vigorously until all the ice has melted. Wait a minute or two for the system to reach equilibrium before recording the bubble volume and bath temperature. 10. Finally, cool the beaker to below 5º. To do this, place your 150 ml beaker and syringe in the plastic bin of ice, without losing the air trapped in the syringe. You can also add ice directly to the beaker. ontinue stirring and allow the bath to cool completely before recording the temperature and gas bubble volume. lean-up 1. Remove your syringe from the 150 ml beaker, empty any liquid remaining into the beaker, and place the syringe in the container labeled for the sample it contained. 2. Empty the water from your 150 ml beaker into the waste container in the Instructor s hood, NT IN TE SINKS. The water is contaminated with a small amount of organic liquid and must be treated separately. E12Z-3

4 Name Lab Partner Section Date DATA SETIN Experiment 12Z Sample Liquid Sample Boiling T ( o ) Barometric P (mm g) Gas Bubble Volume (ml) Temperature (º) Low Temperature (ice bath) Measurement: Gas Bubble Volume (ml) Temperature (º) E12Z-4

5 DATA TREATMENT Experiment 12Z 1. Set up an Excel spreadsheet, creating columns for the Volume (ml) and Temperature ( ) data you collected, as well as additional columns for Volume (in L) and Absolute Temperature (in K), Insert your volume and temperature data for all temperatures except the lowest (ice bath) temperature into the spreadsheet. (We will treat the lowest temperature data separately so there is no need to include it in the table.) 2. Use Excel formulas to convert all volumes into Liters and all elsius temperatures into absolute (Kelvin) temperatures. 3. Determine the number of moles of air, n air, trapped in the gas bubble using your gas volume from the lowest temperature (ice bath) measurement. Treat the gas bubble as an ideal gas: n air = P barv/rt, where the volume and temperature values come from your ice bath data and the pressure is the barometric pressure P bar. (NTE: At the low temperature used for this calculation, the vapor pressure of water or organic sample is negligible. Thus, the gas bubble consists essentially of air alone at this temperature. The pressure of air in the bubble is also assumed to be the same as the atmospheric pressure in the room.) Show your calculation of n air. n air = mol air 4. Add a column to your spreadsheet for the number of moles of air. Enter the value just calculated into all the cells in the column. (Since the number of moles of air remains constant, these values will all be the same.) 5. At all but the lowest temperature, the gas bubble is a mixture of air plus sample vapor. The partial pressure of the air in the bubble will change because the temperature and volume change, even though the amount of air (the number of moles of air) is constant. Add a column P air to your spreadsheet and for each of the data pairs tabulated in the spreadsheet, set up an Excel formula to calculate the partial pressure of air in the gas bubble: P air = n airrt/v, recognizing that the number of moles of air does not change from step 3 above. Show a sample calculation of P air for your highest temperature data point. P air = atm air 6. By Dalton s Law, the total pressure of a gas mixture is the sum of the partial pressures of the components. Your mixture contains air plus the test substance, and the total pressure of the mixture is the measured barometric pressure. Add a column P sample to your spreadsheet and for each of the data pairs tabulated in the spreadsheet, use Dalton s Law to calculate the partial pressure of the test sample, P sample in the gas bubble: P sample = P bar - P air. Show a sample calculation for your highest temperature data point. P sample = atm 7. Add columns for the natural logarithm of the sample pressure, ln(p sample), and the inverse Kelvin temperature (1/T(K)) to your spreadsheet. Use Excel formulas to calculate these quantities from the other entries. E12Z-5

6 8. Plot the vapor pressure curve P sample vs. T(K) for your test substance. Fit the data with an exponential trendline. Show the trendline equation and R 2 value on the plot, as usual. (Use Format Trendline Label to set the number values to scientific notation.) Enter your trendline equation and R 2 values here. Equation of trendline: R 2 for trendline: 9. The normal boiling point of a substance is the temperature at which the vapor pressure of the substance equals 1 atm. Use the trendline equation just determined to find the normal boiling point of your sample. alculate the % error from the literature value (see table). Show your work. sample measured normal boiling point K % error 10. Select only the points where the gas bubble was between 4.0 ml and 2.0 ml. Plot Ln(P sample) vs. 1/T for these points. Perform a linear trendline analysis on this plot, showing the trendline equation and R 2 value on the plot. (Use Format Trendline Label to set the number values to scientific notation.) Enter your trendline equation and R 2 values here. (The reason for limiting the selection is that the lowest temperature points are typically not well equilibrated.) Equation of trendline: Slope (with units): R 2 for trendline: 11. According to equation (1), the slope of your plot is related to the heat of vaporization of the sample. From the slope of your regression line, determine the molar enthalpy (heat) of vaporization, Δ vap, and compare it to the accepted value of your test sample (see table). alculate the percent error. Show your work. sample measured Δ vap kj/mol % error 12. Enter your result in the lass Data table. Record data for the other samples from the lass Data table, so that you have experimental Δ vap values for all samples used. Literature Values: Molar Mass Normal Boiling eat of Vaporization Sample Formula (g/mol) Point 1 ( o ) vap (kj/mol) Water yclopentane yclohexane exane Ethyl Acetate Methyl t-butyl Ether eptane ctane eptanol Glushko Thermocenter, Entropy and eat apacity of rganic ompounds in NIST hemistry WebBook, NIST Standard Reference Database Number 69, Eds. P.J. Linstrom and W.G. Mallard, National Institute of Standards and Technology, Gaithersburg MD, 20899, (retrieved April 8, 2016). 2 hikos, J.S. and Acree Jr., W.E, J. Phys. hem, Ref. Data, 32, 519 (2003), averages from data sets over temperature range ~280K-350K E12Z-6

7 Sample water LASS DATA Experiment 12Z eat of Vaporization ( vap) (kj/mol) Average (kj/mol) E12Z-7

8 QUESTINS Experiment 12Z 1. The table shows the structures of some substances which may be used in this experiment Name hemical Formula ondensed Formula water 2 2 cyclopentane ethyl acetate methyl t-butyl ether 5 12 ( 3) 3 3 hexane ( 2) 4 3 heptane ( 2) heptanol ( 2) 5 2 octane ( 2) 6 3 Structural Formula 2 2 Using the structures provided, circle all of the intermolecular forces present in each pure liquid. Water - Dipole-Dipole ydrogen Bonding London Dispersion Forces yclopentane - Dipole-Dipole ydrogen Bonding London Dispersion Forces Ethyl acetate - Dipole-Dipole ydrogen Bonding London Dispersion Forces Methyl t-butyl ether - Dipole-Dipole ydrogen Bonding London Dispersion Forces exane - Dipole-Dipole ydrogen Bonding London Dispersion Forces eptane - Dipole-Dipole ydrogen Bonding London Dispersion Forces 1-eptanol - Dipole-Dipole ydrogen Bonding London Dispersion Forces ctane - Dipole-Dipole ydrogen Bonding London Dispersion Forces 2. For a clear comparison, focus on the substances of question 1 that can utilize NLY London dispersion forces. What trend would be expected for the normal boiling points and heats of vaporization of these substances? Is that trend observed in the literature (not experimental) data? (see table on p. E12Z-6). EXPLAIN YUR ANSWER E12Z-8

9 3. For that same set of substances that utilize NLY London dispersion forces, what trend would you expect for the value(s) of the equilibrium vapor pressure(s) of the substance(s) at 25 o? EXPLAIN YUR ANSWER. 4. onsider the liquids actually used in your experiment, and tabulated in the lass Data table. Use the averages of the measured heats of vaporization, and rank the substances from greatest to least Δ vap highest Δ vap lowest Δ vap 5. onsider the substances actually measured, and the experimental rankings of question 4. ow well do the measured values match your expectation based on the intermolecular forces you identified for the substances? EXPLAIN YUR ANSWER. E12Z-9

10 Name Section PRE-LAB Questions Experiment 12Z Date Use the following table to consider the effect of structure on intermolecular forces. methyl t-butyl ether pentane pentanoic acid pentanol Which of the following pure liquids would be expected to utilize NLY London Dispersion Forces? ircle ALL that apply.) a) methyl t-butyl ether ( 5 12) b) pentane ( 5 12) c) pentanoic acid ( ) d) pentanol ( 5 12) 2. Which of the following pure liquids would be expected to utilize Dipole-Dipole forces, but NT ydrogen Bonding? (ircle ALL that apply.) a) methyl t-butyl ether ( 5 12) b) pentane ( 5 12) c) pentanoic acid ( ) d) pentanol ( 5 12) 3. Which of the following pure liquids would be expected to utilize ydrogen Bonding? (ircle ALL that apply.) a) methyl t-butyl ether ( 5 12) b) pentane ( 5 12) c) pentanoic acid ( ) d) pentanol ( 5 12) 4. Which of the following alkanes will have the highest normal boiling point? a) 2 6 b) 4 10 c) 6 14 d) 8 18 e) Which of the following alkanes will have the highest vapor pressure at room temperature? a) 2 6 b) 4 10 c) 6 14 d) 8 18 e) For a substance studied in Exp. 12Z, a plot of ln P vs 1/T (in K) had the trendline equation y = x Based on Equation 1 in the lab, use the slope of the line to determine the value of Δ vap for the substance, in kj/mol. 7. The normal boiling point of a substance is the temperature at which the vapor pressure of the substance equals 1 atm. If the vapor pressure curve (P in atm vs. T in K) follows the equation P = 1.13x10-7 e T, what is the temperature (in K) of the normal boiling point? E12Z-10

EXPERIMENT 7 - Distillation Separation of a Mixture

EXPERIMENT 7 - Distillation Separation of a Mixture Purpose: a) To purify a compound by separating it from a non-volatile or less-volatile material. b) To separate a mixture of two miscible liquids (liquids