Chapter 7

Based on the strength of their molecular dipole moments, which compound should have the higher boiling point, HF or HBr? Explain.

Based on the strength of their molecular dipole moments, which compound should have the higher boiling point, \(\mathrm{H}_{2} \mathrm{~S}\) or \(\mathrm{H}_{2} \mathrm{O}\) ? Explain.

The compound ethylene, \(\mathrm{H}_{2} \mathrm{C}=\mathrm{CH}_{2}\), is a gas at room temperature, but the compound polyethylene, made from a large number of ethylene units bonded together, is a solid at room temperature. Explain why.

(a) Draw a dot diagram for \(\mathrm{NH}_{3}\) and one for \(\mathrm{PH}_{3}\). (b) Is either molecule polar? (The electronegativities of \(\mathrm{N}, \mathrm{P}\), and \(\mathrm{H}\) are \(3.0,2.1\), and 2.1, respectively.) (c) In which substance are the London forces stronger? Explain. (d) Basing your answer solely on the London forces in the two substances, which substance would you expect to have the higher boiling point? Explain.

Draw some molecules of ammonia, \(\mathrm{NH}_{3}\), and use dotted lines to show the hydrogen bonding between them.

For the hydrogen halides, the order of boiling points is \(\mathrm{HF}>\mathrm{HI}>\mathrm{HBr}>\mathrm{HCl}\). (a) Why does HF have the highest boiling point? (b) Why is the boiling point of HI greater than that of \(\mathrm{HBr}\) and \(\mathrm{HCl}\) ?

Sodium oxide, \(\mathrm{Na}_{2} \mathrm{O}\), is a white-gray powder that sublimes (goes directly from solid to gas) at \(1275{ }^{\circ} \mathrm{C}\). Is sodium oxide a molecular or nonmolecular solid?

Elemental sulfur is a yellow solid that melts at \(113{ }^{\circ} \mathrm{C}\). Studies show sulfur to be a molecular solid consisting of \(S_{8}\) molecules held together in a lattice. What forces must be overcome to melt the solid?

The following are solids at room temperature (about \(25^{\circ} \mathrm{C}\) ). Classify them as molecular, ionic, network, or metallic: (a) Zirconium, \(\mathrm{Zr}, \mathrm{mp}=1852{ }^{\circ} \mathrm{C}\) (b) Lead, \(\mathrm{Pb}, \mathrm{mp}=328^{\circ} \mathrm{C}\) (c) Calcium nitride, \(\mathrm{Ca}_{3} \mathrm{~N}_{2}, \mathrm{mp}=1195^{\circ} \mathrm{C}\) (d) Graphite form of carbon, \(\mathrm{C}\), sublimes at \(3652{ }^{\circ} \mathrm{C}\) (e) Yellow phosphorus, \(\mathrm{P}_{4}, \mathrm{mp}=44^{\circ} \mathrm{C}\)

Consider cooling a gas so that it gets colder and colder. (a) Explain why this would eventually cause the gas to condense into a liquid. (b) Explain what would happen to this gas if you could somehow turn off all intermolecular forces.

What does the kinetic energy of molecules have to do with changing phases?

Is it incorrect to say that molecules are motionless in the liquid phase? Explain.

Explain in molecular terms how heating causes a liquid to change to the gas phase.

Water vapor liquefies when cooled below \(100^{\circ} \mathrm{C}\). Gaseous nitrogen liquefies when cooled below \(-196^{\circ} \mathrm{C}\). What does this information tell you about the relative strengths of the intermolecular forces for these molecules?

Why does a gas expand to fill the container it is in, but a liquid and a solid do not?

Why does a gas expand to fill the container it is in, but a liquid and a solid do not?

On the molecular level, describe each phase of matter with respect to the amount of order present.

Draw a picture that shows how three polar \(\mathrm{HBr}\) molecules in the gas phase would attract one another. What kind of intermolecular force is involved?

Explain what gives rise to London forces and when they occur.

Would you expect \(\mathrm{CCl}_{4}\) or \(\mathrm{CBr}_{4}\) to have the higher boiling point? Explain your answer.

Chloromethane \(\left(\mathrm{CH}_{3} \mathrm{Cl}\right)\) has a much higher boiling point than methane \(\left(\mathrm{CH}_{4}\right)\). Give two reasons for this.

Propane \(\left(\mathrm{C}_{3} \mathrm{H}_{8}\right)\) is a gas at room temperature, whereas octane \(\left(\mathrm{C}_{8} \mathrm{H}_{18}\right)\) is a liquid. Explain why this is so.

What is wrong with the statement "London forces are always weaker than dipole- dipole forces"?

Consider the molecules HF and HCl. The electronegativities of the atoms involved are \(\mathrm{H}, 2.1 ; \mathrm{Cl}, 3.0 ; \mathrm{F}, 4.0 .\) (a) Which of the two molecules is more polar? Explain your answer. (b) For which substance are the dipole-dipole attractions between the molecules stronger? Explain your answer. (c) Which gas would, upon cooling, liquefy first? Explain your answer. Also, indicate which compound would have the higher boiling point based on your answer. (d) Do both molecules also experience London forces of attraction? If yes, which would have the greater London forces?

When discussing the intermolecular forces between methanol molecules, chemists usually ignore any London forces between them. Why are they justified in doing this?

Long-chain hydrocarbon molecules of the type \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{20} \mathrm{CH}_{3}\) are solids and are used for things like waxes. The CH bonds are essentially nonpolar. Why are waxes solid at room temperature?

Which of the molecules below would not form hydrogen bonds? (a) \(\mathrm{CH}_{3} \mathrm{OH}\) (b) \(\mathrm{CH}_{3} \mathrm{OCH}_{3}\) (c) \(\mathrm{CH}_{3} \mathrm{COOH}\) (d) \(\mathrm{NH}_{3}\) (e) None of these would form hydrogen bonds.

What is so special about hydrogen atoms that makes hydrogen bonding possible?

Two different compounds have the same elemental composition, \(\mathrm{C}_{3} \mathrm{H}_{8} \mathrm{O}\). One has a low boiling point and the other a much higher boiling point. What attractive force must be present in one of these compounds that is not present in the other?

What do we mean by induced dipole when discussing London forces?

Why do we use dotted lines rather than solid lines to represent hydrogen bonds?

Why would either covalent bonds or London forces be inappropriate for attaching the two strands of DNA to each other?

How is it possible for a nonpolar molecule to have a higher boiling point than a polar one?

What is the fundamental difference between a molecular substance and a nonmolecular substance?

Is it sometimes, always, or never possible to tell from a compound's formula whether it is a molecular or nonmolecular substance?

What do we mean by a "lattice of ions"?

In general, nonmolecular solids have much higher melting points than molecular solids. Why is this so?

What is a network covalent substance? Give an example

Predict which compound in each pair will have the higher melting point, and explain why for each pair. (a) HI or KI (b) \(\mathrm{Na}_{2} \mathrm{O}\) or \(\mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{CaF}_{2}\) or \(\mathrm{HF}\) (d) \(\mathrm{SiO}_{2}\) or \(\mathrm{CO}_{2}\)

All of the following are solids at room temperature. Classify them as molecular, ionic, network, or metallic. (a) Potassium \((\mathrm{K}), \mathrm{mp}=64{ }^{\circ} \mathrm{C}\) (b) Potassium chloride \((\mathrm{KCl}), \mathrm{mp}=770{ }^{\circ} \mathrm{C}\) (c) Red phosphorus \((\mathrm{P}), \mathrm{mp}=590{ }^{\circ} \mathrm{C}\) (d) Boron triiodide \(\left(\mathrm{BI}_{3}\right), \mathrm{mp}=50{ }^{\circ} \mathrm{C}\)

How does the attraction between two molecules depend on the distance between them?

Describe how the molecules in a liquid behave as the temperature of the liquid increases.

Solid ice is less dense than liquid water. What does this say about how many water molecules there are per unit volume of ice relative to how many water molecules there are per unit volume of liquid water?

Which of the following best describes a liquid: (a) The phase of matter in which particles are typically separated by the least distance (b) The phase of matter in which particles are in a fixed, rigid arrangement (c) The phase of matter in which particles completely fill the volume of their container (d) The phase of matter in which particles are in a loose, changeable arrangement but do not completely fill the volume of their container

At \(25^{\circ} \mathrm{C}\), fluorine, \(\mathrm{F}_{2}\), and chlorine, \(\mathrm{Cl}_{2}\), are gases but bromine, \(\mathrm{Br}_{2}\), is a liquid. What does this say about the intermolecular forces in bromine relative to those in fluorine and chlorine?

In which of the following are there no dipoledipole forces between molecules? Justify your answer. (a) \(\mathrm{AsH}_{3}\) (b) \(\mathrm{CO}_{2}\) (c) \(\mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{SeCl}_{2}\)

Which is more likely to be a gas at room temperature, \(\mathrm{CH}_{4}\) or \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) ? Justify your answer.

Which of the following does not form hydrogen bonds? Justify your choice. (a) Methyl alcohol, \(\mathrm{CH}_{3} \mathrm{OH}\) (b) Hydrofluoric acid, HF (c) Ammonia, \(\mathrm{NH}_{3}\) (d) Methane, \(\mathrm{CH}_{4}\)

Which of the following would you expect to have the highest boiling point? Justify your choice. (a) Propane, \(\mathrm{C}_{3} \mathrm{H}_{8}\) (b) Carbon dioxide, \(\mathrm{CO}_{2}\) (c) Ethyl alcohol, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) (d) Methyl fluoride, \(\mathrm{CH}_{3} \mathrm{~F}\)

Which of the following would most likely be a gas at room temperature? Justify your choice. (a) \(\mathrm{NaCl}\) (b) \(\mathrm{C}_{2} \mathrm{H}_{2}\) (c) Na metal (d) \(\mathrm{CH}_{3} \mathrm{~F}\)