Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, and solubilities (miscibilities) of substances. Molar mass, molecular shape, and polarity affect the strength of different intermolecular forces, which influence the magnitude of physical properties across a family of molecules.
Temporary attractive forces like dispersion are present in all molecules, whether they are polar or nonpolar. They cause gases to condense (liquefy) and liquids to freeze (solidify) under very low temperature (or high pressure) conditions. Dispersion forces arise from temporary dipoles caused by the asymmetrical distribution of electrons around the atom's nucleus. Atoms (or molecules) with a greater number of electrons (higher molar mass) display stronger dispersion forces than lighter atoms (or molecules). The melting point and boiling point trend of halogens demonstrate this effect. Moving down the group, from fluorine to iodine, melting points and boiling points increase with increasing atomic size (or mass). This increase may be rationalized by considering how the strength of dispersion forces is affected by the electronic structure of the atoms or molecules in the substance. In a larger atom, the valence electrons are, on average, farther from the nuclei than in a smaller atom. Thus, they are less tightly held and can more easily form the temporary dipoles that produce the attraction. The measure of how easy or difficult it is for another electrostatic charge (for example, a nearby ion or polar molecule) to distort a molecule’s charge distribution (its electron cloud) is known as polarizability.
A molecule that has a charge cloud that is easily distorted is said to be very polarizable and will have large dispersion forces; one with a charge cloud that is difficult to distort is not very polarizable and will have small dispersion forces. The shapes of molecules also affect the magnitudes of the dispersion forces between them. For example, boiling points for the isomers n-pentane, isopentane, and neopentane are 36 °C, 27 °C, and 9.5 °C, respectively. Even though these compounds are composed of molecules with the same chemical formula, C5H12, the difference in boiling points suggests that dispersion forces in the liquid phase are different, being greatest for n-pentane and least for neopentane. The elongated shape of n-pentane provides a greater surface area available for contact between molecules, resulting in correspondingly stronger dispersion forces. The more compact shape of isopentane offers a smaller surface area available for intermolecular contact and, therefore, weaker dispersion forces. Neopentane molecules are the most compact of the three, offering the least available surface area for intermolecular contact and, hence, the weakest dispersion forces.
Polar substances exhibit dipole–dipole attractions. The effect of this attraction is apparent when comparing the properties of polar HCl molecules to nonpolar F2 molecules. Both HCl and F2 consist of the same number of atoms and have approximately the same molecular mass. At a temperature of 150 K, molecules of both substances would have the same average KE. However, the dipole–dipole attractions between HCl molecules are sufficient to cause them to “stick together” to form a liquid, whereas the relatively weaker dispersion forces between nonpolar F2 molecules are not, and so this substance is gaseous at this temperature. The higher normal boiling point of HCl (188 K) compared to F2 (85 K) is a reflection of the greater strength of dipole–dipole attractions between HCl molecules, compared to the attractions between nonpolar F2 molecules.
A special type of dipole–dipole force—hydrogen bonds—have a pronounced effect on the properties of condensed phases (liquids and solids). For example, consider the trends in boiling points for the binary hydrides of group 15 (NH3, PH3, AsH3, and SbH3), group 16 hydrides (H2O, H2S, H2Se, and H2Te), and group 17 hydrides (HF, HCl, HBr, and HI). On progressing down the groups, the polarities of the molecules decrease slightly, whereas the sizes of the molecules increase substantially. The effect of increasingly stronger dispersion forces dominates that of increasingly weaker dipole–dipole attractions, and the boiling points are observed to increase steadily. Using this trend, the predicted boiling points for the lightest hydride for each group would be about −120 °C (for NH3), −80 °C (for H2O), and −110 °C (for HF). However, the measured boiling points for these compounds are about −33.34 °C (for NH3), 100 °C (for H2O), and 19.5 °C (for HF), all of which are dramatically higher than the predicted trends. The stark contrast between our naïve predictions and reality provides compelling evidence for the strength of hydrogen bonding.
Effect of Polarity on Miscibility
Liquids that can be homogeneously mixed in any proportion are said to be miscible. Miscible liquids have similar polarities. Consider, for example, methanol (CH3OH) and water (H2O), two liquids that are polar and capable of hydrogen bonding. On mixing, methanol and water will interact through intermolecular hydrogen bonds and mix; thus, they are miscible. Likewise, nonpolar liquids like hexane (C6H14) and bromine (Br2) are miscible with each other through dispersion forces. The chemical axiom “like dissolves like” is useful to predict the miscibility of compounds. Two liquids that do not mix to an appreciable extent are called immiscible. For example, nonpolar hexane is immiscible in polar water. Relatively weak attractive forces between the hexane and water do not adequately overcome the stronger hydrogen bonding forces between water molecules.
This text is adapted from Openstax, Chemistry 2e, Section 10.1: Intermolecular forces. and Section 11.3: Solubility.
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