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Chem 162 Experiment 5 Determination of the Enthalpy of Vaporization of Water and the Exploration of Intermolecular Forces

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Chem 162 Experiment 5 Determination of the Enthalpy of Vaporization of Water and the Exploration of Intermolecular Forces Chem 162 Experiment 5 Determination of the Enthalpy of Vaporization of Water and the Exploration of Intermolecular Forces Pre-lab questions: 1. In terms of kinetic energy (KE), explain why vapor pressure increases with increasing temperature. 2. Provide a definition for what is meant by the normal boiling point for a liquid. 3. Consider the Clausius-Clapeyron equation on the next page. The data you will collect in this lab will be vapor pressures of water at various temperatures. What should be the trend in the data? Based on the literature value for ∆Hvap of H2O, make your prediction. 4. Calculate the normal boiling point of a compound that has a vapor pressure of 500 torr at 20ºC and requires 32 kJ to vaporize 1 mole? 5. Complete the table in Part II of the experiment. The question wants you to determine what volume of each of those liquids contains the same number of molecules as 1 microliter (µL) of water. List the volumes here as part of your pre lab answers. 6. Rank the compounds in the Part II table according to increasing vapor pressure at constant temperature. Justify your answers. Be sure to have all of your data tables prepared in your lab notebook. Introduction The goal of this laboratory is to determine the vapor pressure of water over a limited temperature range. We will the use this data to find the heat of vaporization of water (∆Hvap) using the Clausius- Clapeyron equation. Vapor pressure is measured as the equilibrium pressure of a vapor above a solvent. In a closed container, solvent molecules will escape the liquid phase and enter the vapor phase. Similarly, some molecules in the vapor phase will condense back into the liquid phase. After some time, the rate of evaporation will equal the rate of condensation. This is the point at which the equilibrium vapor pressure can be measured. In general, as intermolecular attractions (IMA) increase between solvent molecules, the vapor pressure decreases. The logic is as follows: • If intermolecular attractions are strong then the substance is likely to have a strong surface tension (cohesion among molecules at the liquids surface). • In order for a liquid to vaporize, molecules must break the surface tension to escape from the liquid. • Thus, if the surface tension is strong, solvent molecules are unable to break free from the liquid surface, the amount of vapor above the liquid is low, and thus vapor pressure is low. For low boiling liquids, the case is just the opposite. Weak IMA lead to an easier transition from liquid to the vapor due to high vapor pressures (or larger amounts of solvent already in the vapor phase). It should also be apparent that vapor pressure for a liquid will increase with increasing temperature, regardless of the IMA present in solution. This relationship is not linear, but using the inverse of temperature vs. the ln P, given by the Clausius-Clapeyron equation below, shows a direct relationship: P is the vapor pressure of the solvent, ∆Hvap is the heat of vaporization, R is the ideal gas constant 8.314 , T is temperature and C is a constant (not related to heat capacity). In this form, the equation resembles y = mx + b, where ln P is on the y-axis, 1/T is on the x-axis, the slope of the line (m) is -∆Hvap/R. Thus, taking simple laboratory measurements we can now easily find the ∆Hvap for any liquid. Part I: Enthalpy of Vaporization (∆Hvap) A sample of air is trapped in an inverted 10.0 mL graduated cylinder in a beaker of water (ask the instructor to demonstrate the set-up). The water is heated to 80ºC and the trapped air becomes saturated with water vapor. The pressure of this gas then equals the pressure of the air plus the pressure of the vapor. The bath is allowed to cool and temperature readings are taken at 5ºC intervals. The pressure in the graduated cylinder, Pcyl, is equal to the atmospheric pressure, Patm. This pressure, Pcyl, is due to BOTH the pressure of the water vapor, Pvap, and the air pressure, Pair. The pressure of the vapor can be determined by difference if we know Pair. The Pair at each temperature can be determined using the ideal gas law, PV=nRT. What is left unknown is the moles of gas in the cylinder. If we cool the water down to 5ºC, the Pvapor in the cylinder is negligible and the Pair at 5ºC will equal Patm. If you observe the volume of the air at 5ºC, then you have all the pieces you need to solve PV=nRT to find the number of moles of air, n. NOTE: Because the meniscus is reversed on the graduate cylinder scale, a small systematic error will be introduced at the gas-water interface. Compensate for this by subtracting 0.20 mL from each volume reading. Procedure 1. Make sure that your 10.0 mL graduated cylinder is marked with 0.1 mL divisions. Fill the cylinder about 2/3 full with water. Cover the top with your finger and invert the cylinder into a beaker filled with water. There should be about 4 mL of air trapped in the cylinder. 2. Add water to the beaker to ensure that the trapped air is surrounded by water (do this the best that you can!). Heat the apparatus to about 85ºC, or until the air expands beyond the scale on the cylinder. Remove the Bunsen burner and allow the water to cool. 3. When the gas in the cylinder contracts enough to allow the scale to be read, record the volume of the air to the nearest 0.01 mL and the temperature to 0.1oC. 4. Measure T and V every 5ºC, until the temperature reaches 50ºC. While you are waiting for the system to cool, you should perform Part II of the experiment, remembering to take readings every 5ºC or so. 5. Once cooled, next move the apparatus to a trough that is filled with ice water and put ice into the beaker until the temperature is below 5ºC. WHEN THE TEMPERATURE STABILIZES record the volume and temperature. This is the data you can use to calculate moles of air trapped in the cylinder. Part II Rate of Evaporation Perform this part of the experiment while waiting for Part I to cool down. Remember to take temperature readings in Part I every 5 degrees! To measure relative vapor pressure, we will determine the amount of time required for the same number of molecules of each liquid to evaporate. Before carrying out the experiment, it is necessary for you to determine what volumes of ethanol, dichloromethane, and acetone contain the same number of molecules as are contained in 1 microliter of water. The data in the table will enable you to make the calculations. Volume equivalent Density, Molar mass, Substance to g/mL g/mole 1 µL of water Water, H2O 1.00 18.015 1.00 Ethanol, CH3CH2OH 0.7893 46.07 Acetone, (CH3)2CO 0.7899 58.08 Dichloromethane, 1.3266 84.93 CH2Cl2 Procedure 1. Obtain a small amount (1 µL) of each liquid from the table on the next page. 2. Rinse a 5-µL syringe three times with one of the 3 liquids, then draw up exactly the volume YOU calculated above. 3. Dispense the syringe contents onto a watch glass, and measure the time required for the liquid to evaporate. Your goal is to achieve reproducibility in your technique, so that the precision in measured time is less than 5 percent (error of 1 second in 20 seconds total). 4. Make repeated ejections of the first liquid until you are confident that you are ejecting and measuring the time exactly the same way every time. 5. Rinse the syringe with a second liquid and measure the time required for the calculated volume to evaporate completely. NOTE: WATER AND DICHLOROMETHANE ARE IMMISCIBLE. FOLLOW THE SOLVENT ORDER IN THE TABLE TO AVOID MISCIBILITY PROBLEMS. 6. Make at least five measurements for each liquid and determine the average evaporation time for each liquid. Calculations and Results Note: Be sure to COMPLETE all experiments before you attempt the calculations! 1. Calculate the moles of air trapped in the cylinder using PV = nRT. See step 5. Part I. 2. Calculate the pressure of air (Pair) at EACH temperature in Part One. 3. Calculate the pressure of water (Pwater) at each temp. Remember Patm = Pair + Pwater 4. Plot ln P vs. 1/T and analyze the data using linear regression for the relationship between vapor pressure and temperature. 5. From the graph, calculate the experimental value for enthalpy of vaporization, ΔHvap for water. NOTE: BE UNIT CONSCIOUS WHEN CALCULATING ΔHVAP. Post Lab Questions 1. What is the literature value for the ΔHvap of H2O? How does your calculated value compare? Determine a percent error and discuss any potential sources of error. o 2. Using the linear regression on your graph, estimate the vapor pressure of H2O at 90 C. 3. Rank the relative vapor pressures of the solutions you measured in Part II according to your data. Is this the same ranking you predicted in your pre-lab? Explain any discrepancies that arise. .
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