# Practice Problems Unit 4 Intermolecular Forces Which gas, SO2 or

```1.
Practice Problems
Unit 4
Intermolecular Forces
Which gas, SO2 or CO2, should be least ideal at STP? Explain
2.
Why do gases under high pressure deviate from ideal behavior?
3.
Why do gases at temperatures near their boiling point deviate from ideal behavior?
4.
Which letter illustrates the types of molecular forces?
Dipole-Dipole
5.
Dispersion
H-bond
ion-dipole
For each pair, circle the molecule with the higher boiling point and then justify your choice.
Pair
Justification
H2O & H2S
Ne & Kr
Cl2 & SO2
6.
7.
Explain the boiling points for the two isomers.
In addition to dispersion, what type of force would you expect between the following molecules?
H2
H2S
CHF3
NH3
6.
For which substance would you predict the highest heat of vaporization? Explain.
a) F2
b) H2O c) HF d) NaCl e) Br2
7.
A sample of water is heated from a liquid at 40oC to a gas at 110oC.
a.
b.
On the heating curve diagram provided above, label each of the following regions:
Liquid, only ; Gas, only; Phase change
For section QR of the graph, state what is happening to the water molecules as heat is added.
c.
For section RS of the graph, state what is happening to the water molecules as heat is added.
1
8.
The graph below represents the heating curve of a substance that starts as a solid below its freezing point.
a.
b.
c.
d.
e.
What is the melting point of this substance?
At what temperature does the substance go from a liquid to a gas?
Neatly label the diagram using the following terms: melting, freezing, vaporizing, and condensation.
How much energy does this substance absorb as it goes from 65°C to 90°C? The Cp is 0.565, and the mass is 2.0 g.
How much energy is released as the substance cools from 90°C to 70°C
Base your answers to questions 9-11 on the heating curve below, which represents a substance starting as a solid below its melting
point and being heated at a constant rate over a period of time.
9.
What is happening to the average kinetic energy of the particles during segment BC?
10. What is happening to the average KE of the particles during segment EF?
11. How does this heating curve illustrate that the heat of vaporization is greater than the heat of fusion?
12. What is the name of the process by which:
a. a solid changes directly into a gas?
b.
a liquid changes into a gas?
c.
a gas changing into a liquid?
d.
a liquid changing into a solid?
e.
a gas changing into a solid?
f.
a solid changing into a liquid?
12. Referring to the figure shown here, what is the phase of water at point 1?
13. How many phases are there at points 2, 4, and 8?
14. What changes is indicated by the line from point 1, 2, 3, 4, and 5?
15. What is changing as indicated by the points 6, 4, 7, 8, and 9?
16. At the conditions indicated by point 3 and 6, what is the stable phase of water?
17. In this sketch, what is the temperature of the triple point of water with respect to 0°C?
18. In the diagram above, what is the melting point of ice at exactly 1 atm?
2
19. What is significant about the critical point?
Graph for
Question 25
20. Consider the phase diagram in your notes.
a. How does melting point change when pressure increases?
b. How would the diagram differ for most substances?
Answer the questions based on the phase diagram below.
Temperature (oC)
21.
Can liquid CO2 exist at room pressure?
What happens to CO2(s) at -78.5oC?
Which is the most dense phase for CO2?
What is the triple point pressure for CO2?
What is the critical temperature for CO2?
22. Explain how pure water can boil at room temperature when placed in an evacuated bell jar.
23. Explain why baking takes longer at high elevations.
24. What is the total amount of energy absorbed by 10.0 g of H20 that starts at an initial temperature of -5ºC and has a final
temperature of 105 ºC?
25. The graph shown at the top of the page shows vapor pressure curves for two substances, A and B:
a. What is the vapor pressure of A at 35C?
b. What is the vapor pressure of B at 35C?
c. At what temperature is the vapor pressure of A equal to 106.6 kPa?
d. What is the vapor pressure of B at this temperature?
e. At what temperature is the vapor pressure of B equal to 106.6 kPa?
f. What is meant by “normal boiling point”?
g. What is the normal boiling point of A?
h. What is the normal boiling point of B?
i. At what temperature would A boil if atmospheric pressure were 93.3 kPa?
j. What would the atmospheric pressure have to be in order for B to boil at the temperature you gave as your answer to
Question i ?
26. Which liquid has the strongest intermolecular bonding?
27. Which liquid is the most volatile?
28. What would happen if you slowly heated a beaker filled with
the three liquids from 20oC to 100oC?
3
29. Trend in Halogens
Graph the molar enthalpy of vaporization of F2, Cl2, and Br2.
Halogen
Molar Mass
Hvap (kJ/mol)
F2
38.0
6.6
Cl2
70.9
20.4
Br2
159.8
30.0
I2
253.8
31. Use the information in the table below to respond to
the statements and questions that follow. Your
answers should be in terms of principles of molecular
structure and intermolecular forces.
Molar Mass
a)
b)
c)
Predict the value of I2: ____ kJ/mol
What type of IMF is exhibited by these molecules
Why does the Hvap increase for these molecules as
the molar mass increases? Explain the trend in terms of
IMF's.
30. Graph the molar enthalpy of vaporization of H2O,
H2S, H2Se, and H2Te.
Molecule
Molar Mass
Hvap (kJ/mol)
H2O
18.0
40.7
H2S
34.1
18.7
H2Se
81.0
19.9
H2Te
129.6
23.8
(a)
List the type of intermolecular force associated with each of
the above compounds.
(b)
Energy is required to boil ethanol. Consider the statement
“As ethanol boils, energy goes into breaking
CC bonds, CH bonds, and OH bonds.” Is the statement
(c)
Ethanol is completely soluble in water, whereas ethanethiol
has limited solubility in water. Account for the difference in
solubilities between the two compounds in terms of
intermolecular forces.
Molar Mass
a) Identify the type of IMF for each molecule.
Molecule
Strongest IMF
H2O
H2S
H2Se
H2Te
b) Explain the general shape of this graph.
c) Account for the trend in the larger three molecules in terms of
IMF’s.
4
32. Use a check mark and indicate the strongest IMF holding together crystals of the following:
Molecular Crystal
London
forces
Dipoledipole
attractions
Hydrogen
Bonds
Metal
Ionic Crystal
Network Solid
Metallic
Bonds
Ionic
Bonds
Covalent
Bonds
NH3
Kr
HCl
F2
KMnO4
NaCl
SO2
CO2
C3H8
CH4
CH3Cl
HF
C6H6
NO
H2SO4
WC
Si
SiO2
C(graphite)
N2
CH3OH
Ag
(C2H5)2NH
NaOH
Al
PCl3
33. Complete the chart for each type of solid.
Metallic
Covalent Network
Molecular
Structural Unit
Bond name
Bond strength
Melting point
Solubility
Conductivity
Malleability
34. Explain the following observations. You must discuss both of the substances in your explanation.
a) SO2 melts at 201 K and SiO2 melts at 1,883 K.
b)
Cl2 boils at 238 K and HCl boils at 188 K.
c)
KCl melts at is 776oC and NaCl melts at 801oC.
d)
Si melts at 1,410oC and Cl2 melts at -101oC.
Ionic
5
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