Chemical bonds between two or more atoms in a molecule are referred to as intramolecular forces. They are generally categorized as ionic, covalent, or metallic. These are the bonds that are broken and reformed during a chemical reaction. Significantly lower in strength are intermolecular forces. Intermolecular forces are generally attractive forces that exist between and among molecules of all shapes, sizes, and masses. They exist in the gaseous, liquid, and solid phases, as well as in the solution phase, where water is acting as the solvent to hydrate the solute.

They are important to understand because they are directly related to macroscopic properties such as melting point, boiling point, vapor pressure and volatility, and the energy needed to overcome forces of attraction between molecules in changes of state. Just as important, they also help determine the solubility of gases, liquids, and solids in various solvents (i.e., whether two substances are soluble or miscible in each other) and in part explain or reflect the “like dissolves like” principle. They are also critical in determining the structure of biologically active molecules such as DNA and proteins.

The categories of intermolecular, attractive forces for neutral or uncharged molecules in order of increasing strength are:

  • • London dispersion forces—exhibited by all molecules
  • • Polar forces or interactions—exhibited by all asymmetric molecules
  • • Hydrogen bonding—exhibited by molecules containing O-H, N-H, or F-H bonds

More generally, however, these forces of attraction, including charged and uncharged species, can be usefully organized and expanded, again in order of increasing strength, as follows:

  • • London dispersion forces—important in nonpolar substances (substances having no permanent dipoles).
  • • Ion-dipole forces—important in salts (consisting of ions) dissolved in water (aqueous solutions), also referred to as “hydration forces”
  • • Dipole-dipole forces—important in polar, covalent substances, which have permanent dipoles
  • • Hydrogen bonding, important in O-H, N-H, and F-H interactions
  • • Ion-Ion forces—important for ions only, usually present in aqueous solutions

The term “van der Waals forces” refers to dipole-induced dipole interactions and are a special case of the more general term “London dispersion forces.” Note that the distance between the oxygen and hydrogen atom within the water molecule is about 100 pm (where pm represents picometer, which is 1 x 10-12 meters). This is an intramolecular distance. In contrast, the distance between the oxygen atom of one water molecule and the hydrogen atom of a neighboring water molecule is about 180 pm. This is the intermolecular distance, typical in a hydrogen bond, and is almost double the intramolecular distance.

For example, polar substances such as acetone (C3H6O) will dissolve in other polar substances such as methyl chloride (CH3Q) but not in carbon tetrachloride (CCU), a nonpolar substance. Some polar substances, such as ethanol (C2H3OH), will dissolve in water in all proportions, that is, are completely miscible, because of hydrogen bonding between the hydrogen atom of the ethanol molecule and the oxygen atom of the water molecule. The same is true of ammonia and hydrofluoric acid, where hydrogen bonding occurs. Note that dimethyl ether (C2OH6), an isomer of ethanol, does not undergo hydrogen bonding with water because of its different molecular structure, that is, an ether vs. an alcohol (see Chapter 3, Table 3.6—ethanol vs. dimethyl ether). Hexane (C6H14) will dissolve in octane (C6H18) because both substances are characterized by nonpolar bonds. Naturally, size (length) and shape (branching of the molecule) also play a role in solubility considerations, and it is often difficult to determine which factor is more important in predicting solubility between two substances.

The greater the forces of attractions between molecules in a liquid, the greater the energy that must be supplied to separate them. Hydrogen bonding is a key reason why low molecular weight alcohols have much higher-than-expected boiling points, that is, heats of vaporization, in comparison to nonpolar hydrocarbons like hexane.

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