Boiling Point Elevation 4-4 Boiling Point Elevation If you dissolve a substance such as ordinary table salt (NCAA) in water, the boiling point of the water will increase relative to the boiling point of the pure water. In this assignment, you will dissolve a sample of NCAA in water and then measure the boiling point elevation for the solution. 1. Start Virtual Chemical and select Boiling Point Elevation from the list of assignments. The lab will open in the Calorimeter laboratory with a calorimeter on the lab bench and a sample of sodium chloride (NCAA) on the balance. 2.
Record the mass of the sodium chloride in the data table. If it is too small to read, click on the Balance area to zoom in, record the reading, and then return to the laboratory. 3. 100 ml of water is already in the calorimeter. Use the density of water at ICC (0. 998 g/ml) to determine the mass from the volume and record it in the data table. Make certain the stirrer is On (you should be able to see the shaft rotating). Click on the green heater light on the control panel to turn on the heater and begin heating the water. Click the clock on the wall labeled Accelerate to accelerate the laboratory time if necessary. Observe the temperature until the first appearance of steam comes from the calorimeter. Immediately click the red light on the heater to turn it off and then record the temperature as the boiling point of pure water in the data table. Letting the water boil will decrease the mass of the water present in the calorimeter. Note that the boiling point may be different than COCO if the atmospheric pressure is not 760 Tour. The current atmospheric pressure for the day can be checked by selecting Pressure on the LED meter on the wall. 5. Drag the weigh paper to the calorimeter and add the NCAA.
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Wait 30 seconds for he salt to dissolve and then turn on the heater. When steam first appears observe and record the temperature in the data table. 6. If you want to repeat the experiment, click on the red disposal bucket to clear the lab, click on the Stockroom, click on the clipboard, and select Preset Experiment #2, Boiling Point Elevation – Niacin. Data Table mass Niacin | 4. 0035 | mass water | 99. 8 | boiling temp of pure water 1 100. ICC I boiling temp to solution The boiling point elevation can be predicted using the equation AT = KGB x m x 7. , where AT is the change in boiling point, I is the number of ions in the solution ere mole of dissolved NCAA (I = 2), m is the impolite of the solution, and KGB is the molar boiling point constant for water which is 0. 51 CO/m. Change in boiling point = KGB * I = 0. 51 * [(4. 0035 158. 5)/0. 0998 2 = 0. 6994 co Calculate the predicted change in boiling point, in co for your solution. 101 ISBN: 0-536-06762-7 Virtual Chemical: General Chemistry, Student Lab Manual/Workbook, V. 2. 5, Third Edition, by Brian F. Whitfield and Matthew C. Capsule. Published by Prentice Hall. Copyright 2006 by Pearson Education, Inc.
Colligating Properties 8. The change in boiling point must be added to the boiling point of pure water in our experiment in order to compare the predicted boiling point with the actual boiling point. What is the calculated boiling point of the solution? Compare this with the actual boiling point. 99. 8 + . 69 = 100. ICC, Actual boiling point = 101. ICC 102 Edition, by Brian F. Whitfield and Matthew C. Capsule. Published by Prentice Bole’s Law 5-1: Bole’s Law: Pressure and Volume Robert Bayle, a philosopher and theologian, studied the properties of gases in the 17th century.
He noticed that gases behave similarly to springs; when compressed or expanded, they tend to ‘spring’ back to their original volume. He published his findings in 1662 in a monograph entitled The Spring of the Air and Its Effects. You will make observations similar to those of Robert Bayle and learn about the relationship between the pressure and volume of an ideal gas. 1. Start Virtual Chemical and select Bole’s Law: Pressure and Volume from the list of assignments. The lab will open in the Gases laboratory. 2. Note that the balloon in the chamber is filled with 0. 00 moles of an ideal gas (MM = 4 g/mol) at a temperature of 298 K, a pressure of 1. 00 ATM, and a volume of 7. 336 L. To the left of the Pressure LCD controller is a lever hat will decrease and increase the pressure as it is moved up or down; the digit changes depending on how far the lever is moved. Digits may also be clicked directly to type in the desired number. You may want to practice adjusting the lever so that you can decrease and increase the pressure accurately. Make sure the moles, temperature, and pressure are returned to their original values before proceeding. 3. Click on the Lab Book to open it.
Back in the laboratory, click on the Save button to start recording P, V, T, and n data to the lab book. Increase the pressure from 1 ATM to 0 ATM one atmosphere at a time. Click Stop to stop recording data, and a blue data link will appear in the lab book. To help keep track of your data links, enter ‘Ideal Gas 1’ next to the link. 4. Zoom Out by clicking the green arrow next to the Save button. Click Return Tank on the gas cylinder. On the table underneath the experimental chamber is a switch to choose Real gases or Ideal gases. Click on the Ideal Gases and choose the cylinder labeled Ideal 8 (Ideal 8 MM = 222 g/mol).