SAT Subject Test Chemistry

PART 2

REVIEW OF MAJOR TOPICS

CHAPTER 3

Bonding

INTERMOLECULAR FORCES OF ATTRACTION

The term intermolecular forces refers to attractions between molecules. Although it is proper to refer to all intermolecular forces as van der Waals forces, named after Johannes van der Waals (Netherlands), this concept should be expanded for clarity.

Dipole-Dipole Attraction

One type of van der Waals forces is dipole-dipole attraction. It was shown in the discussion of polar covalent bonding that the unsymmetrical distribution of electronic charges leads to positive and negative charges in the molecules, which are referred to as dipoles. In polar molecular substances, the dipoles line up so that the positive pole of one molecule attracts the negative pole of another. This is much like the lineup of small bar magnets. The force of attraction between polar molecules is called dipole-dipole attraction. These attractive forces are less than the full charges carried by ions in ionic crystals.

London Dispersion Forces

Another type of van der Waals forces is called London dispersion forces. Found in both polar and nonpolar molecules, it can be attributed to the fact that a molecule/atom that usually is nonpolar sometimes becomes polar because the constant motion of its electrons may cause uneven charge distribution at any one instant. When this occurs, the molecule/atom has a temporary dipole. This dipole can then cause a second, adjacent atom to be distorted and to have its nucleus attracted to the negative end of the first atom. London dispersion forces are about one-tenth the force of most dipole interactions and are the weakest of all the electrical forces that act between atoms or molecules. These forces help to explain why nonpolar substances such as noble gases and the halogens condense into liquids and then freeze into solids when the temperature is lowered sufficiently. In general, they also explain why liquids composed of discrete molecules with no permanent dipole attraction have low boiling points relative to their molecular masses. It is also true that compounds in the solid state that are bound mainly by this type of attraction have rather soft crystals, are easily deformed, and vaporize easily. Because of the low intermolecular forces, the melting points are low and evaporation takes place so easily that it may occur at room temperature. Examples of such solids are iodine crystals and moth balls (paradichlorobenzene and naphthalene).

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Weakest of all, the London dispersion forces are one-tenth the force of most dipole attractions.

Hydrogen Bonds

A proton or hydrogen nucleus has a high concentration of positive charge. When a hydrogen atom is bonded to a highly electronegative atom, its positive charge will have an attraction for neighboring electron pairs. This special kind of dipole-dipole attraction is called a hydrogen bond. The more strongly polar the molecule is, the more effective the hydrogen bonding is in binding the molecules into a larger unit. As a result the boiling points of such molecules are higher than those of similar nonpolar molecules. Good examples are water and hydrogen fluoride.

Studying Figure 14 shows that in the series of compounds consisting of H2O, H2S, H2Se, and H2Te an unusual rise in the boiling point of H2O occurs that is not in keeping with the typical slow increase of boiling point as molecular mass increases. Instead of the expected slope of the line between H2O and H2S, which is shown in Figure 14 as a dashed line, the actual boiling point of H2O is quite a bit higher—100°C. The explanation is that hydrogen bonding occurs in H2O but not to any significant degree in the other compounds.

Figure 14. Boiling Points of Hydrogen Compounds with Similar Electron Dot Structures

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Notice how hydrogen bonding elevates the boiling point of H2O above the expected slope.

This same phenomenon occurs with the hydrogen halides (HF, HCl, HBr, and HI). Note in Figure 14 that hydrogen fluoride, HF, which has strong hydrogen bonding, shows an unexpectedly high boiling point.

Hydrogen bonding also explains why some substances have unexpectedly low vapor pressures, high heats of vaporization, and high melting points. In order for vaporization or melting to take place, molecules must be separated. Energy must be expended to break hydrogen bonds and thus break down the larger clusters of molecules into separate molecules. As with the boiling point, the melting point of H2O is abnormally high when compared with the melting points of the hydrogen compounds of the other elements having six valence electrons, which are chemically similar but which have no apparent hydrogen bonding.