The magnitude of the attraction of one particle for another is important in determining whether the substance containing those particles is a solid, a liquid, or a gas under normal conditions (20%deg;C, 1 atm). These attractive forces also represent a part of the potential or stored energy of a sample. We know that when the kinetic energy of a sample changes, its temperature changes. When the potential energy of a sample changes, the temperature does not change. Instead the energy that is added or subtracted breaks or forms bonds. For example, when 6.02 kJ of energy are added to 1 mol (18 g) of ice at 0°C, there is no change in temperature, only a change in state from solid to liquid. The added energy counteracts the forces that held the water molecules in the rigid ice structure. It does not break the bonds within molecules, only those between molecules.

H₂O(s) + energy H₂O(l)

The bonds between the atoms, ions, or molecules in a substance range from very strong to comparatively weak. The melting and boiling points of substances of similar molecular weight are a measure of the relative strength of their intermolecular, interatomic, or interionic bonds. The higher-melting substance contains stronger bonds between its particles; the lower-melting substance has comparatively weaker interparticle bonds.

A. Intermolecular Forces in Liquids
Low-melting solids and compounds that are liquids at room temperature are usually covalent. Ionic compounds almost always melt well above room temperature. (The melting point of sodium chloride, a typical ionic solid, is 801°C.) In samples of covalent compounds that are either low-melting or liquid, three types of intermolecular forces are possible. Two of these have already been discussed: dispersion forces Section 9.8 and dipole-dipole interactions Sections 7.4 and 9.8. The third type of intermolecular force is known as hydrogen bonding.

FIGURE 10.1 When a drop of a dye is placed on the surface of some water, the molecules of dye are buffeted by the moving molecules of water and dispersed throughout.

 

Hydrogen bonding is a special kind of dipole-dipole interaction. It represents the attraction between a small partially negative atom (usually nitrogen, oxygen, or fluorine) and a partially positive hydrogen (usually bonded to another very electronegative atom like nitrogen, oxygen, or fluorine) in another molecule. We can illustrate the existence of hydrogen bonding by comparing the boiling points of the hydrides of several elements. (A hydride of an element is the compond that element forms with hydrogen; for example, ammonia is the hydride of nitrogen.) We know that the boiling point of a substance is a measure of its intermolecular forces. Figure 10.2 plots the boiling points of the hydrides of the elements in Groups IV, V, VI, and VII of the periodic table.

FIGURE 10.2 Molecular weight versus boiling point of the hydrides of Groups IV, V, VI, and VII. The solid line shows the actual boiling points; the dashed line, the predicted boiling point of the hydride of the lightest member of each group if there were no hydrogen bonding.

 

Elements of Group IV show the regular increase in boiling point that would be expected for a regular increase in molecular weight. For Groups V, VI, and VII, the hydrides of the three heaviest members of each group show a regular increase in boiling point, but the boiling point of the hydride of the lightest member of each group is much higher than would be predicted. Their expected boiling points are shown by the dashed continuation of the line plotting boiling points against molecular weight. For Group VI, the graph shows that the boiling point of water is approximately 200°C higher than would be predicted on the basis of molecular weight. Ammonia and hydrogen fluoride also boil at a higher temperature than expected. We postulate that the abnormally high boiling points of ammonia, water, and hydrogen fluoride are due to strong interactions between the molecules, so strong that these compounds behave more as aggregates of molecules than as single molecules. These aggregates are held together by hydrogen bonds. Within this group, the hydrogen bond strength is greatest in H-F, less in H-O, and even less in H-N bonds. Notice that this order corresponds to the order of decreasing electronegativity of F, O, and N.

Let us consider the nature of hydrogen bonding in water at the molecular level. We know that the O-H bond is polar covalent. The oxygen atom bears a partial positive charge.

When two water molecules come close to each other, a partially positive hydrogen atom of one water molecule interacts with the partially negative oxygen atom of the other to form a hydrogen bond. The hydrogen atom is apparently bonded to oxygen atoms in both molecules. We say that the two molecules of water are held together by a hydrogen bond. Figure 10.3 shows several water molecules. The hydrogen bonds between neighboring molecules are shown as dashed lines. The overall effect is a network of bonds with no separate molecules.

In closing this discussion of hydrogen bonding, we must emphasize that, even though hydrogen bonds are stronger than dipole-dipole interactions or dispersion forces, they are still much weaker than intermolecular or interionic bonds.

Hydrogen bonds play an important role in chemistry, particularly in the chemistry of biochemical molecules such as proteins and nucleic acids.

FIGURE 10.3 Hydrogen bonding in ice. Each water molecule is bonded to four others. Its two hydrogen atoms are attracted to the oxygen atoms in two other water molecules, and its oxygen atom attracts hydrogens in two more water molecules.

B. Interparticle Forces in Solids
High-melting solids are metals, ionic compounds, or network covalent compounds. This latter group includes diamonds, quartz, and silica, the hardest materials known. The interparticle forces in high-melting solids also fall into three categories:

1. Network covalent solids
In Chapter 7 we discussed covalent bonds in small molecules. Most of these are low-melting, but the interatomic bonds in these small molecules are very strong. Network covalent solids (Figure 10.4a) can be thought of as huge covalent molecules of enormous molecular weight. In these molecules every atom is covalently bonded to several other atoms so as to form a network of covalent bonds (hence the name). Such compounds have very high melting points. Diamonds (mp 3550°C) and quartz (SiO₂; mp 1610°C) are typical network covalent solids. When these compounds melt, a regular covalent bond of shared electrons is broken, the same kind of bond that is broken when water is decomposed to hydrogen and oxygen.

2. Ionic bonds
In Chapter 7 we discussed ionic compounds, pointing out that the particles in an ionic compound are not atoms or molecules but ions. In the solid state, the ions are held in a regular, rigid structure (Figure 10.4b) by the electrostatic forces of attraction between oppositely charged particles, as noted in Section 7.1A ionic solids usually have very high melting points, indicating the great strength of the electrostatic forces between ions.

FIGURE 10.4 Bonding in solids: (a) a network covalent solid (notice how closely the structure resembles the structure of hydrogen bonding in ice); (b) an ionic solid; (c) metal.

3. Metallic bonding
The bonding in metals differs from that in ionic or covalent solids. We know that metals typically have one, two, or three valence electrons. One picture of a metallic solid shows these valence electrons as a fluid within which float the nuclei surrounded by their inner electrons (Figure 10.4c). The conductivity of metals is due to the movement of this "sea" of electrons. Hammering metal into a thin sheet spreads out the fluid; similarly, drawing a metal into thin wire is a rearrangement of the sea of electrons containing the nuclei into a thin stream. Metal fatigue, a problem in the aircraft industry and in nuclear power plants, is associated with metals losing their "fluid" nature and assuming a rigid structure. The melting points of metals vary over a wide range. Mercury is the lowest-melting (-39°C); tungsten is one of the highest-melting (3410°C).

Example: Two sets of compounds are given. Arrange each in order of increasing strength of interparticle force. a. CH3CH2OH   CH3CH2CH2CH3   Ni b. PH3   MgO CH3CH3 Solution a. The correct order is CH3CH2CH2CH3, CH3CH2OH, Ni. The first compound, CH3CH2CH2CH3, is a hydrocarbon. It contains no polar bonds. Between its molecules are only weak dispersion forces. Teh second compound, CH3CH2OH, contains hydrogen bonded to the very electronegative oxygen; its molecules would be held together by comparatively strong hydrogen bonds. The last substance, nickel, is a metal. Interatomic bonds in metals are strong. b. The order is CH3CH3, PH3, MgO. CH3CH3 is a hydrocarbon containing no polar bonds. Its intermolecular forces are very weak. In PH3, the phosphorus is not sufficiently electronegative to allow the formaiton of hydrogen bonds but the phosphorus-hydrogen bond is still polar; there would be weak dipole-dipole interaction between the molecules. MgO is essentially ionic; the difference in electronegativity between magnesium and oxygen is appreciable. There will be strong interionic bonds between Mg2+ and O2-.