8.E: Attractive Forces

The following questions are related to the material covered in this chapter. For additional discussion on each topic, also check the links included in each heading.

8.2: Polar Bonds

1. Give an example of a nonpolar covalent bond. How do you know it is nonpolar?
2. Give an example of a polar covalent bond. How do you know it is polar?
3. How do you know which side of a polar bond has the partial negative charge? Identify the negatively charged side of each polar bond.
   1. H–Cl
   2. H–S
4. How do you know which side of a polar bond has the partial positive charge? Identify the positively charged side of each polar bond.
   1. H–Cl
   2. N-F
5. Label the bond between the given atoms as nonpolar covalent, slightly polar covalent, definitely polar covalent, or likely ionic.
   1. H and C
   2. C and F
   3. K and F
6. Label the bond between the given atoms as nonpolar covalent, slightly polar covalent, definitely polar covalent, or likely ionic.
   1. S and Cl
   2. P and O
   3. Cs and O

8.3: Molecular Shapes

1. What is the basic premise behind VSEPR?
2. What is the difference between the electron group geometry and the molecular geometry?
3. Identify the electron group geometry and the molecular geometry of each molecule.
   1. H₂S
   2. POCl₃
4. Identify the electron group geometry and the molecular geometry of each molecule.
   1. CS₂
   2. H₂S
5. Identify the electron group geometry and the molecular geometry of each molecule.
   1. HCN
   2. CCl₄
6. Identify the electron group geometry and the molecular geometry of each molecule.
   1. BI₃
   2. PH₃
7. What is the geometry of each species?
   1. CN⁻
   2. PO₄³⁻
8. What is the geometry of each species?
   1. PO₃³⁻
   2. NO₃⁻
9. What is the geometry of each species?
   1. COF₂
   2. C₂Cl₂ (both C atoms are central atoms and are bonded to each other)
10. What is the geometry of each species?
    1. CO₃²⁻
    2. N₂H₄ (both N atoms are central atoms and are bonded to each other)

Answers

1. Electron pairs repel each other.
2.  
3. 1. electron group geometry: tetrahedral; molecular geometry: bent
   2. electron group geometry: tetrahedral; molecular geometry: tetrahedral
4.  
5. 1. electron group geometry: linear; molecular geometry: linear
   2. electron group geometry: tetrahedral; molecular geometry: tetrahedral
6.  
7. 1. linear
   2. tetrahedral
8.  
9. 1. trigonal planar
   2. linear and linear about each central atom

8.5: Intermolecular Forces

1. What type of intermolecular force do all substances have?
2. What is necessary for a molecule to experience dipole-dipole interactions?
3. What is necessary for a molecule to experience hydrogen bonding?
4. How does varying the temperature change the preferred phase of a substance?
5. Identify the strongest intermolecular force present in each substance.
   1. He
   2. CHCl₃
   3. HOF
6. Identify the strongest intermolecular force present in each substance.
   1. CH₃OH
   2. (CH₃)₂CO
   3. N₂
7. Identify the strongest intermolecular force present in each substance.
   1. HBr
   2. C₆H₅NH₂
   3. CH₄
8. Identify the strongest intermolecular force present in each substance.
   1. C₁₀H₂₂
   2. HF
   3. glucose

Answers

1. dispersion force
2.  
3. An H atom must be bonded to an N, O, or F atom.
4.  
5. 1. dispersion forces
   2. dipole-dipole interactions
   3. hydrogen bonding
6.  
7. 1. dipole-dipole interactions
   2. hydrogen bonding
   3. dispersion forces

Q8.4.1

Explain the differences between the microscopic and the macroscopic properties of matter. Is the boiling point of a compound a microscopic or macroscopic property? molecular mass? Why?

Q8.4.2

How do the microscopic properties of matter influence the macroscopic properties? Can you relate molecular mass to boiling point? Why or why not?

Q8.4.3

​​For a substance that has gas, liquid, and solid phases, arrange these phases in order of increasing density.

1. strength of intermolecular interactions.
2. compressibility.
3. molecular motion.
4. order in the arrangement of the molecules or atoms.

10.3: Phase Transitions - Melting, Boiling, and Subliming

1. What is the difference between melting and solidification?
2. What is the difference between boiling and condensation?
3. Describe the molecular changes when a solid becomes a liquid.
4. Describe the molecular changes when a liquid becomes a gas.
5. What is the energy change when 78.0 g of Hg melt at −38.8°C?
6. What is the energy change when 30.8 g of Al solidify at 660°C?
7. What is the energy change when 111 g of Br₂ boil at 59.5°C?
8. What is the energy change when 98.6 g of H₂O condense at 100°C?
9. Each of the following statements is incorrect. Rewrite them so they are correct.
   1. Temperature changes during a phase change.
   2. The process of a liquid becoming a gas is called sublimation.
10. Each of the following statements is incorrect. Rewrite them so they are correct.
    1. The volume of a gas contains only about 10% matter, with the rest being empty space.
    2. ΔH_sub is equal to ΔH_vap.
11. Write the chemical equation for the melting of elemental sodium.
12. Write the chemical equation for the solidification of benzene (C₆H₆).
13. Write the chemical equation for the sublimation of CO₂.
14. Write the chemical equation for the boiling of propanol (C₃H₇OH).
15. What is the ΔH_sub of H₂O? (Hint: see Table 10.3.1 "Enthalpies of Fusion for Various Substances" and Table 10.3.2 "Enthalpies of Vaporization for Various Substances".)
16. The ΔH_sub of I₂ is 60.46 kJ/mol, while its ΔH_vap is 41.71 kJ/mol. What is the ΔH_fus of I₂?

Answers

1. Melting is the phase change from a solid to a liquid, whereas solidification is the phase change from a liquid to a solid.
2.  
3. The molecules have enough energy to move about each other but not enough to completely separate from each other.
4.  
5. 890 J
6.  
7. 10.7 kJ
8.  
9. 1. Temperature does not change during a phase change.
   2. The process of a liquid becoming a gas is called boiling; the process of a solid becoming a gas is called sublimation.
10.  
11. Na(s) → Na(ℓ)
12.  
13. CO₂(s) → CO₂(g)
14.  
15. 46.69 kJ/mol

8.8: General Properties of Aqueous Solutions

1. What are the advantages to carrying out a reaction in solution rather than simply mixing the pure reactants?
2. What types of compounds dissolve in polar solvents?
3. Describe the charge distribution in liquid water. How does this distribution affect its physical properties?
4. Must a molecule have an asymmetric charge distribution to be polar? Explain your answer.
5. Why are many ionic substances soluble in water?
6. Explain the phrase like dissolves like.
7. What kinds of covalent compounds are soluble in water?
8. Why do most aromatic hydrocarbons have only limited solubility in water? Would you expect their solubility to be higher, lower, or the same in ethanol compared with water? Why?
9. Predict whether each compound will dissolve in water and explain why.
   1. toluene
   2. acetic acid
   3. sodium acetate
   4. butanol
   5. pentanoic acid
10. Predict whether each compound will dissolve in water and explain why.
    1. ammonium chloride
    2. 2-propanol
    3. heptane
    4. potassium dichromate
    5. 2-octanol
11. Given water and toluene, predict which is the better solvent for each compound and explain your reasoning.
    1. sodium cyanide
    2. benzene
    3. acetic acid
    4. sodium ethoxide (CH₃CH₂ONa)
12. Of water and toluene, predict which is the better solvent for each compound and explain your reasoning.
    1. t-butanol
    2. calcium chloride
    3. sucrose
    4. cyclohexene
13. Compound A is divided into three equal samples. The first sample does not dissolve in water, the second sample dissolves only slightly in ethanol, and the third sample dissolves completely in toluene. What does this suggest about the polarity of A?
14. You are given a mixture of three solid compounds—A, B, and C—and are told that A is a polar compound, B is slightly polar, and C is nonpolar. Suggest a method for separating these three compounds.
15. A laboratory technician is given a sample that contains only sodium chloride, sucrose, and cyclodecanone (a ketone). You must tell the technician how to separate these three compounds from the mixture. What would you suggest?
16. Many over-the-counter drugs are sold as ethanol/water solutions rather than as purely aqueous solutions. Give a plausible reason for this practice.
17. Use a list of solubility rules to determine if each of these will dissolve in water.
    1. NaCl
    2. AgCl
    3. CaCl₂
    4. PbCl₂
18. Use a list of solubility rules to determine if each of these will dissolve in water.
    1. Na₂CO₃
    2. CaCO₃
    3. FeCl₃
    4. Fe(OH)₃
19. What distinguishes a weak electrolyte from a strong electrolyte?
20. Which organic groups result in aqueous solutions that conduct electricity?
21. It is considered highly dangerous to splash barefoot in puddles during a lightning storm. Why?
22. Which solution(s) would you expect to conduct electricity well? Explain your reasoning.
    1. an aqueous solution of sodium chloride
    2. a solution of ethanol in water
    3. a solution of calcium chloride in water
    4. a solution of sucrose in water
23. Which solution(s) would you expect to conduct electricity well? Explain your reasoning.
    1. an aqueous solution of acetic acid
    2. an aqueous solution of potassium hydroxide
    3. a solution of ethylene glycol in water
    4. a solution of ammonium chloride in water
24. Which of the following is a strong electrolyte, a weak electrolyte, or a nonelectrolyte in an aqueous solution? Explain your reasoning.
    1. potassium hydroxide
    2. ammonia
    3. calcium chloride
    4. butanoic acid
25. Which of the following is a strong electrolyte, a weak electrolyte, or a nonelectrolyte in an aqueous solution? Explain your reasoning.
    1. magnesium hydroxide
    2. butanol
    3. ammonium bromide
    4. pentanoic acid
26. Which of the following is a strong electrolyte, a weak electrolyte, or a nonelectrolyte in aqueous solution? Explain your reasoning.
    1. H₂SO₄
    2. diethylamine
    3. 2-propanol
    4. ammonium chloride
    5. propanoic acid