Intermolecular Forces (sections 11.1 – 11.3) Particles in gases are presumed to be so far apart that any forces they might exert on one another are considered negligible. In liquids and solids, however, particles are tightly packed, and have substantial attractions for one another. Were it not for these attractive forces, substances would not exist as solids or liquids. Intermolecular forces are a collective term for the attractions between particles. Intermolecular forces differ from chemical bonds, which hold atoms together to form molecules—intermolecular forces act between separate molecules. The attractions are generally much weaker than those of chemical bonds, although their strength is ultimately determined by the identity of the substances. Attractions between ions are stronger than attractions found among covalent particles, and are not classified as intermolecular forces (since ionic compounds do not exist as discrete molecules). The strength of ionic attractions is responsible for the properties generally associated with ionic substances—high melting points, high boiling points, hard, rigid crystalline structures—due to the great amount of energy required to separate ions from one another. While it is convenient at times to collectively describe the attractions between molecules, there are several types of intermolecular forces which must be considered when predicting or explaining the physical properties of molecular substances.

It should not be surprising that polar covalent molecules should attract one another; their polarity causes them to act at least somewhat as ionic particles. The greater the dipole moment, the greater the attractive forces between the molecules, as the more electronegative atom of one molecule attracts the less electronegative atom of another. [The more electronegative atom has a greater electron density, resulting in a “partial negative charge”, symbolized as –, with + representing the partial positive charge resulting from lesser electron density near the less electronegative atom.] +  –

–

+

+

–

–

+

–

+

–

+

Figure 1 Polar molecules in the solid state (and, to a lesser extent, in the liquid state) will tend to align in three dimensions as suggested by Figure 1 due to the dipole-dipole forces among them. [Polar covalent molecules will also be attracted by ions, in the form of ion-dipole forces. The strength of the attractions will again be affected by the dipole moment of the covalent molecule, and by the charge density of the ion. A familiar example of an ion-dipole attraction is the hydration of ions by

polar water molecules. Pure substances, however, cannot be both ionic and covalent and do not exhibit ion-dipole attractions. We will restrict further discussion to attractions among the particles of pure substances.] Even nonpolar molecules can attract one another if they become temporarily polarized. The random nature of electron motion in atoms will induce slight, highly transient dipoles even in completely-nonpolar substances. When random redistribution of electrons creates a dipole in one atom or molecule, dipoles will be induced in the neighboring atoms, and the effect will cascade through the entire sample, creating attractions between the particles. The separation of charge in the induced dipole is not as great as in a molecule with a permanent dipole, but is enough to allow slight attractions among nonpolar and nonpolar molecules. The more electrons in the nonpolar molecule, the more polarizable it is (because there are more electrons available to redistribute). The dipoles are highly transient, disappearing almost instantly, but the resulting van der Waals forces are sufficient to condense even the least polar of atoms and molecules. van der Waals forces occur in all molecules, and are the only intermolecular attractions among nonpolar molecules; they can be quite strong among species with many electrons, although in general these are the weakest of intermolecular attractions.

The types of intermolecular attractions determine the physical properties of substances. Weaker attractions allow particles to be easily separated from one another, while stronger attractions prevent such separation. One can determine the relative strength of attractions by observing the volatility of a substance (the extent to which it evaporates at a given temperature). Substances with greater volatility have lower melting and boiling points. It may be appropriate to list the types of attractions visited thus far, along with their associated properties. Attraction

Exhibited in

Properties

van der Waals

nonpolar covalent molecules

low melting points low boiling points

dipole – dipole

polar covalent molecules

low melting and boiling points (but higher than for nonpolar molecules with comparable molar mass)

ionic bonding

ionic compounds

high melting and boiling points

While most nonpolar covalent compounds exhibit only van der Waals forces, exceptional covalent compounds, such as diamond (crystalline carbon), quartz (a form of silicon dioxide, SiO2), and silicon carbide (SiC, known informally as “carborundum”) exists as covalent network solids. A crystal of a network solid is not a collection of separate molecules attracted to one another via dipole-dipole or dispersion forces. Instead, the entire crystal is a single molecule, in which the atoms are chemically bonded to one another. The resulting properties of network solids resemble those of ionic compounds: covalent network solids have very high melting and boiling points (typically higher than those found in ionic compounds), and very hard, brittle crystals. Indeed, diamond and silicon carbide are two of the hardest substances known. We can rank the relative strength of intermolecular attractions based on the bonding type exhibited by the substance. Molecular covalent compounds typically have the weakest attractions, ionic compounds exhibit much stronger attractions, and network covalent molecules exhibit the strongest attractions. The bonding triangle and other methods will indicate whether covalent or ionic bonding occurs; distinguishing between molecular covalent and network covalent bonding is more difficult. Fortunately, there are only four significant examples of network covalent bonding: diamond, silicon, silicon carbide (SiC), and silicon dioxide (SiO2).

Attraction

Exhibited in

Properties

van der Waals

nonpolar covalent molecules

low melting points low boiling points

dipole – dipole

polar covalent molecules

low melting and boiling points (but higher than for nonpolar molecules with comparable molar mass)

ionic bonding

ionic compounds

high melting and boiling points

covalent network

diamond, Si, SiC, SiO2

very high melting and boiling points