Molecules with hydrogen atoms bonded to electronegative atoms such as O, N, and F (and to a much lesser extent Cl and S) tend to exhibit unusually strong intermolecular interactions. Even the noble gases can be liquefied or solidified at low temperatures, high pressures, or both (Table 11.3). The LibreTexts libraries are Powered by MindTouch® and are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. Compounds with higher molar masses and that are polar will have the highest boiling points. Missed the LibreFest? For more information contact us at firstname.lastname@example.org or check out our status page at https://status.libretexts.org. In this section, we explicitly consider three kinds of intermolecular interactions:There are two additional types of electrostatic interaction that you are already familiar with: the ion–ion interactions that are responsible for ionic bonding and the ion–dipole interactions that occur when ionic substances dissolve in a polar substance such as water. Why? As a result, it is relatively easy to temporarily deform the electron distribution to generate an instantaneous or induced dipole. It can be measured experimentally and the values compared to give us an idea of the forces involved in the liquid state Is a similar consideration required for a bottle containing pure ethanol? Vigorous boiling requires a higher energy input than does gentle simmering. As the atomic mass of the halogens increases, so does the number of electrons and the average distance of those electrons from the nucleus. Explain your reasoning. Compounds such as HF can form only two hydrogen bonds at a time as can, on average, pure liquid NH3. In larger atoms such as Xe, however, the outer electrons are much less strongly attracted to the nucleus because of filled intervening shells. This effect, illustrated for two H2 molecules in part (b) in Figure 11.5, tends to become more pronounced as atomic and molecular masses increase (Table 11.3). Of the two butane isomers, 2-methylpropane is more compact, and n-butane has the more extended shape. London dispersion forces are due to the formation of instantaneous dipole moments in polar or nonpolar molecules as a result of short-lived fluctuations of electron charge distribution, which in turn cause the temporary formation of an induced dipole in adjacent molecules. The effect is most dramatic for water: if we extend the straight line connecting the points for H2Te and H2Se to the line for period 2, we obtain an estimated boiling point of −130°C for water! Water has two polar O–H bonds with H atoms that can act as hydrogen bond donors, plus two lone pairs of electrons that can act as hydrogen bond acceptors, giving a net of four hydrogen bonds per H2O molecule. For example, part (b) in Figure 11.6 shows 2,2-dimethylpropane (neopentane) and n-pentane, both of which have the empirical formula C5H12. Larger atoms with more electrons are more easily polarized than smaller atoms, and the increase in polarizability with atomic number increases the strength of London dispersion forces. Ethyl methyl ether has a structure similar to H2O; it contains two polar C–O single bonds oriented at about a 109° angle to each other, in addition to relatively nonpolar C–H bonds. There are three different types of intermolecular forces, London Dispersion forces, Dipole-Dipole forces, and Hydrogen Bonding forces. Explain any trends in the data, as well as any deviations from that trend. Describe the effect of polarity, molecular mass, and hydrogen bonding on the melting point and boiling point of a substance.
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