Study Guide: Chapter 14 Liquids and Solids
(pg 1 of 2)
Why is it that, when we constantly apply heat to ice, it eventually melts to form liquid water? And why does this
water, when heated, eventually form steam?
Also, how can it be that different chemicals are different phases at the same temperature? For example, sodium
chloride is a solid, water is a liquid, and nitrogen is a gas, all at normal room conditions.
These are the types of questions that you should be able to answer by studying Chapter 14 in your text.
The bonds that we discussed in Chapter 12 were intramolecular bonds. That is, these were forces that hold atoms
together in a molecule. We now want to consider intermolecular bonds, or bonds that occur among molecules to
form a solid or a liquid. Figure 14.1 provides a microscopic representation of a solid, a liquid and a gas.
Recall from our discussion of the VSEPR model (Chapter 12) that water is a bent or V-shaped molecule, and is
therefore polar. In other words, there are regions of partial positive and negative charges in a water molecule, and
these opposite charges provide the force that holds these molecules together (see Fig. 14.4, 14.5 & 14.6 in your
book). This attraction is termed dipole-dipole attraction. For molecules such as water, which has a hydrogen
bonded to oxygen that is very electronegative, these dipole-dipole attractions are particularly strong and termed
hydrogen bonding. Hydrogen bonding accounts for some of the quite unusual properties of water. For example,
because these intermolecular bonds are rather strong, water is a liquid at room temperature conditions (water is a
rather small and “lightweight” molecule, and its boiling point is quite high for its molar mass and size).
So why does ice turn to water which turns to steam as we heat our sample? We know that temperature is a
measure of kinetic energy (see Chapter 10), and an increase in kinetic energy will cause the molecules to move more
quickly. The more movement, the more disorder. Look again at Figure 14.1, and note that solids are more ordered
that liquids, which are more ordered that gases.
So this provides us a way of thinking about phase changes. But what about the fact that some chemicals are solids,
some are liquids, and some are gases at the same room conditions? This too can be answered with our model of
intermolecular forces.
Consider, for example, nitrogen gas (N2). Draw the Lewis Structure for N2 and convince yourself that it is a
nonpolar molecule. Because of this there should be no permanent dipole-dipole interactions. Because there seems
to be no strong force holding these molecules together, it should make sense that N 2 is a gas. In fact, it is not
obvious that one N2 molecule would have any attraction for another N2 molecule.
However, we know that N2 can be a liquid if we get the temperature low enough (you may have seen or heard of
liquid nitrogen before). Therefore there must be intermolecular attractions of some sort among all molecules.
These forces are termed London dispersion forces and come about from a random uneven distribution of electrons
around the nucleus (see Fig 14.8). As the temperature is lowered, the molecules slow down. If they are slowed
down enough (which is why liquid nitrogen is very cold) the molecules have enough attraction to come together to
form a liquid.
Look at Figure 14.8, and think about this question, “Will a molecule with more electrons have a greater or smaller
chance of having an uneven distribution of electrons?” Think about this before reading on.
The more electrons there are, the greater the chance that there will be an uneven distribution of electrons. The
more electrons, the greater number of ways there are to be unevenly distribution of electrons. The most
electrons, the greater number of ways there are to be unevenly distributed. As a simple example, consider the
following. If there were only two electrons in a diatomic molecule, the electrons could be evenly distributed or
both could be on one side of the molecule. But with ten electrons there are many more ways for even distribution
(six on one side, four on the other; seven and three, etc.) and many combinations of these. This should led us to
Study Guide: Chapter 14 Liquids and Solids
(pg 2 of 2)
believe that the more electrons a molecule has, the stronger the London dispersion forces will be (the greater the
uneven distribution, the stronger the attraction).
Consider, for example, I2, Br2, and Cl2. Each of these is nonpolar and therefore relies on only London dispersion
forces for intermolecular attraction. Thus the type of attraction for each of these is the same. But I 2 has the
most electrons, and Cl2 has the least. We would predict, then that I2 molecules would exhibit the most attraction
for one another, and Cl2 would exhibit the least. As it turns out, I2 is a solid, Br2 is a liquid, and Cl2 is a gas, all at
room conditions. While our simple model could not predict that I 2 should be a solid, it is pleasing to note that we
can predict the order of attraction.
After reading and studying Chapter 14, you should be able to answer the following questions:
1. Why is sodium chloride a solid with such a high melting point?
2. Why is water a liquid but methane (CH4) a gas at room temperature?
Learning Review:
1. Why does ice float on the surface of liquid water?
2. Which of the following statements about water are true, and which are false?
a. At 0°C water can be either a solid or a liquid.
b. The normal boiling point of water is 97.5°C at 1 atm pressure.
c. Forces among water molecules are called intramolecular forces.
3. Why does it take more energy to convert liquid water to steam than it does to convert ice to liquid water?
4. How much energy is required when 15.0g of liquid water at 75°C is heated to 100.°C and then converted to
steam at 100.°C? The molar heat of vaporization for liquid water is 40.6kJ/mol, and the specific heat of
capacity of liquid water is 4.18J/g°C.
5. How much energy is required to convert 6.0g liquid water to steam at 100°C?
6. Which bond is stronger, a covalent bond or a hydrogen bond?
7. Draw a diagram showing how three NH3 molecules can form hydrogen bonds.
8. Why does water, H2O, boil at a higher temperature than H2S?
9. Explain how London dispersion forces form between two molecules of argon.
10. Match the following terms with the correct definition.
Terms
Definition
a. condensation
the pressure exerted by a liquid at equilibrium
b. vapor pressure
a balance between two opposite processes
c. vaporization
vapor molecules form a liquid
d. equilibrium
a liquid becomes a gas
11. Which of the molecules in each pair would have the greater vapor pressure? To help you determine which
of these molecules is polar, review Section 12.2 and 12.3 of your textbook.
a. CH4, CH3OH
b. H2O, H2S
12. Name the three types of crystalline solids, and give us an example of each.
13. What intermolecular force is responsible for holding each of the following compounds together in the solid
state?
a. KCl
b. HF
c. SiCl4
14. Mixtures of metallic elements form alloys. What are the two kinds of alloys, and how do they differ from
each other?