B. Charles's Law Consider you have the set-up as shown Figure 2, where you 1. Half-fill a 400-mL beaker with water and you 3. Place the stoppered Erlenmeyer flask into a 400-mL beaker half-filled with water. 4. Heat and allow the water to boil for 5 minutes. Figure 2 6. Record the temperature of the boiling water. 7. Cover the hole of the stopper with your finger. Remove the flask from the boiling water and submerge it in an inverted position into a water trough filled with water. 8. Remove your finger to allow water to get into the flask. 9. Take the temperature of water in the trough. 10. Make the level of water inside and outside the flask equal. 11. Cover the hole of the stopper with your finger and remove the flask from the water trough. 12. Measure the volume of water inside the flask by using a graduated cylinder. 13. Fill the flask with water and stopper it. Carefully remove the stopper and measure the total volume of water. Record this as volume of air at higher temperature.

B. Charles's Law Consider you have the set-up as shown Figure 2, where you 1. Half-fill a 400-mL beaker with water and you 3. Place the stoppered Erlenmeyer flask into a 400-mL beaker half-filled with water. 4. Heat and allow the water to boil for 5 minutes. Figure 2 6. Record the temperature of the boiling water. 7. Cover the hole of the stopper with your finger. Remove the flask from the boiling water and submerge it in an inverted position into a water trough filled with water. 8. Remove your finger to allow water to get into the flask. 9. Take the temperature of water in the trough. 10. Make the level of water inside and outside the flask equal. 11. Cover the hole of the stopper with your finger and remove the flask from the water trough. 12. Measure the volume of water inside the flask by using a graduated cylinder. 13. Fill the flask with water and stopper it. Carefully remove the stopper and measure the total volume of water. Record this as volume of air at higher temperature.

Introductory Chemistry: A Foundation

9th Edition

ISBN:9781337399425

Author:Steven S. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Donald J. DeCoste

Chapter15: Solutions

Section: Chapter Questions

Problem 17CR: Are changes in state physical or chemical changes? Explain. What type of forces must be overcome to...

See similar textbooks

Similar questions

Refer to Figure 11.12 to answer these questions: (a) You heat some water to 60 C in a lightweight plastic bottle and seal the top very tightly so gas cannot enter or leave the carton. What happens when the water cools? (b) If you put a few drops of liquid diethyl ether on your hand, does it evaporate completely or remain a liquid? Figure 11.12 Vapor pressure curves for diethyl ether [(C2H3)2O], ethanol (C2H5OH), and water. Each curve represents conditions of T and P of which the two phases, liquid and vapor, are in equilibrium. These compounds exist as liquids for temperatures and pressures to the left of the curve and as gases under conditions to the right of the curve. (See Appendix G for vapor pressures for water of various temperatures.)
Cooking A cook prepares a solution for boiling by adding12.5 g of NaCl to a pot holding 0.750 L of water. Atwhat temperature should the solution in the pot boil?Use Table 14.5 for needed data.
A 0.500-g sample of KCl is added to 50.0 g of water in a calorimeter (Figure 5.12). If the temperature decreases by 1.05 °C, what is the approximate amount of heat involved in the dissolution of the KCl, assuming the specific heat of the resulting solution is 4.18 J/g °C? Is the reaction exothermic or endothermic?
Are changes in state physical or chemical changes? Explain. What type of forces must be overcome to melt or vaporize a substance (are these forces intramolecular or intermolecular)? Define the molar heat of fusion and molar heat of vaporization. Why is the molar heat of vaporization of water so much larger than its molar heat of fusion? Why does the boiling point of a liquid vary with altitude?
Dissolving 3.0 g of CaCl2(s) in 150.0 g of water in a calorimeter (Figure 5.12) at 22.4 °C causes the temperature to rise to 25.8 °C. What is the approximate amount of heat involved in the dissolution, assuming the specific heat of the resulting solution is 4.18 J/g °C? Is the reaction exothermic or endothermic?
How does the process described in the previous item relate to the system shown in Figure 16.4?
Hiking Imagine that on a cold day you are planning to take a thermos of hot soup with you on a hike. Explain why you might fill the thermos with hot water first before filling it with the hot soup.