{"id":766,"date":"2018-03-20T15:50:38","date_gmt":"2018-03-20T15:50:38","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-orgbiochemistry\/?post_type=chapter&p=766"},"modified":"2018-09-19T14:40:22","modified_gmt":"2018-09-19T14:40:22","slug":"8-1-intermolecular-interactions","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-monroecc-orgbiochemistry\/chapter\/8-1-intermolecular-interactions\/","title":{"raw":"8.1 Intermolecular Interactions","rendered":"8.1 Intermolecular Interactions"},"content":{"raw":"
<\/div>\r\n
\r\n
\r\n
\r\n
\r\n

Learning Objectives<\/h3>\r\n
    \r\n \t
  1. Define phase<\/em>.<\/li>\r\n \t
  2. Identify the types of interactions between molecules.<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n

    A phase<\/span><\/strong><\/span>\u00a0is a certain form of matter that includes a specific set of physical properties. That is, the atoms, the molecules, or the ions that make up the phase do so in a consistent manner throughout the phase. Three of the phases are: the solid phase<\/em>, in which individual particles can be thought of as in contact and fixed in place; the liquid phase<\/em>, in which individual particles are in contact but moving past each other; and the gas phase<\/em>, in which individual particles are separated from each other by relatively large distances. Not all substances will readily exhibit all phases. For example, carbon dioxide does not exhibit a liquid phase unless the pressure is greater than about six times normal atmospheric pressure. Other substances, especially complex organic molecules, may decompose at higher temperatures, rather than becoming a liquid or a gas.\u00a0 Several other phases exist but will not be considered here.<\/p>\r\n

    Which phase a substance adopts depends on the pressure and the temperature it experiences. Of these two conditions, temperature variations are more obviously related to the phase of a substance. When it is very cold, H2<\/sub>O exists in the solid form as ice. When it is warmer, the liquid phase of H2<\/sub>O is present. At even higher temperatures, H2<\/sub>O boils and becomes steam.<\/p>\r\n

    Pressure changes can also affect the presence of a particular phase, as indicated for carbon dioxide, but its effects are usually less obvious. We will focus on the temperature effects on phases. Chemical substances follow the same pattern of phases when going from a low temperature to a high temperature: the solid phase, then the liquid phase, and then the gas phase, although some compounds decompose without entering the more mobile phases. However, the temperatures at which these phases are present differ for all substances and can be rather extreme. Table 8.1 \"Temperature Ranges for the Three Phases of Various Substances\"<\/a> shows the temperature ranges for solid, liquid, and gas phases for three substances. As you can see, there is extreme variability in the temperature ranges.<\/p>\r\n\r\n

    \r\n
    Table 8.1<\/span> Temperature Ranges for the Three Phases of Various Substances<\/strong><\/h5>\r\n\r\n\r\n\r\n\r\n\r\n\r\n\r\n\r\n\r\n
    Substance<\/th>\r\nSolid Phase Below<\/th>\r\nLiquid Phase Above<\/th>\r\nGas Phase Above<\/th>\r\n<\/tr>\r\n<\/thead>\r\n
    hydrogen (H2<\/sub>)<\/td>\r\n\u2212259\u00b0C<\/td>\r\n\u2212259\u00b0C<\/td>\r\n\u2212253\u00b0C<\/td>\r\n<\/tr>\r\n
    water (H2<\/sub>O)<\/td>\r\n0\u00b0C<\/td>\r\n0\u00b0C<\/td>\r\n100\u00b0C<\/td>\r\n<\/tr>\r\n
    sodium chloride (NaCl)<\/td>\r\n801\u00b0C<\/td>\r\n801\u00b0C<\/td>\r\n1413\u00b0C<\/td>\r\n<\/tr>\r\n<\/tbody>\r\n
    The melting point<\/em> of a substance is the temperature that separates a solid and a liquid. The boiling point<\/em> of a substance is the temperature that separates a liquid and a gas.<\/th>\r\n<\/tr>\r\n<\/tfoot>\r\n<\/table>\r\n<\/div>\r\n

    What accounts for this variability? Why do some substances become liquids at very low temperatures, while others require very high temperatures before they become liquids? It depends on the strength of the intermolecular interactions<\/strong><\/span><\/span>\u00a0between the particles of substances. (Although ionic compounds are not composed of discrete molecules, we will still use the term intermolecular<\/em> to include interactions between the ions in such compounds.) Substances that experience strong intermolecular interactions require higher temperatures to become liquids and, finally, gases because it takes more energy to overcome the intermolecular attractions and mobilize the particles to enter the liquid or gas phase. \u00a0 Substances that experience weak intermolecular interactions do not need much energy to become liquids and gases and thus will exhibit these phases at lower temperatures.\u00a0 The following paragraphs will consider substances with weakest to strongest intermolecular attraction, thus generally lowest to highest melting\/boiling points.<\/p>\r\nAll molecules\u00a0 experience temporary charge separation caused by chance uneven distribution of electrons in the molecule, which also induces charge separation in adjacent molecules by attracting or repelling that molecule's electrons.\u00a0 While these charge separations are tiny and temporary, they never-the-less cause attractions between neighboring molecules.\u00a0 These very weak intermolecular interactions are called London <\/strong>dispersion forces.\u00a0 <\/span><\/strong><\/span>Molecules that experience no other type of intermolecular interaction will at least experience dispersion forces. Substances that experience only dispersion forces are typically soft in the solid phase and have relatively low melting points. Because dispersion forces are caused by the instantaneous distribution of electrons in a molecule, larger molecules with a large number of electrons can experience substantial dispersion forces. Examples include waxes<\/em>, which are long hydrocarbon chains that are solids at room temperature because the molecules have so many electrons.\r\n

    Weak London dispersion forces are the only attractive forces between nonpolar covalent molecules, resulting in generally low melting and boiling points.\u00a0 Molecule may be nonpolar by having only nonpolar bonds or by having polar bonds that cancel each other.\u00a0\u00a0 Carbon dioxide (CO2<\/sub>) and carbon tetrachloride (CCl4<\/sub>) are examples of nonpolar molecules having polar bonds that cancel each other.\u00a0 Under normal atmospheric pressure, carbon dioxide sublimes rather than melting and boiling.\u00a0 Carbon tetrachloride melts at -22.9\u00a0o<\/sup>C and boils at 76.7\u00a0o<\/sup>C.\u00a0 (Figure 8.1 \"Nonpolar Molecules\"<\/a>).<\/p>\r\n\r\n

    \r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"1500\"]\"image\" Figure 8.1 Nonpolar Molecules. <\/em>Although the individual bonds in both CO2<\/sub> and CCl4<\/sub> are polar, their effects cancel out because of the spatial orientation of the bonds in each molecule. As a result, weak London dispersion forces are the only attractions between molecules and melting and boiling points are generally low.[\/caption]\r\n\r\n<\/div>\r\nIn addition to London dispersion forces, polar molecules experience\u00a0dipole-dipole interactions.<\/strong>\u00a0 As discussed in Sections 4.4 and 4.5, molecules are polar if they have polar bonds that do not cancel each other, resulting in distinct areas of the molecule having \u03b4+ and \u03b4- charges.\u00a0 Figure 8.3 shows the dipole-dipole attractions between molecules of acetone.\u00a0 Acetone melts at -95.4\u00a0o<\/sup>C and boils at 56.5 o<\/sup>C.\u00a0 Despite acetone's stronger dipole-dipole attractions, its melting and boiling points are lower than those for nonpolar carbon tetrachloride.\u00a0 However, carbon tetrachloride is almost 3 times more massive than acetone, so its London dispersion forces\u00a0 add up to a good deal of attraction between molecules.\r\n\r\n\"\"\r\n
    Figure 8.2 Dipole-dipole interaction between molecules of propanone, a polar covalent compound.\u00a0\u00a0\u00a0 https:\/\/chem.libretexts.org\/@api\/deki\/files\/4654\/image089.png?revision=1<\/div>\r\n

    When a polar molecular compound has a hydrogen atom bonded to an atom of one of most electronegative elements, fluorine, oxygen, or nitrogen, the \u03b4+ on the H atom and the \u03b4- on O, N or F atom form especially strong dipole-dipole attractions known as hydrogen bonds<\/strong> between molecules, again in addition to the weaker London forces and any other dipole-dipole attractions. A hydrogen bond is about 10% as strong as a covalent bond.\u00a0 The physical properties of water, which has two O\u2013H bonds, are strongly affected by the presence of hydrogen bonding between water molecules. Figure 8.3 \"Hydrogen Bonding between Water Molecules\" shows how molecules experiencing hydrogen bonding can interact.\u00a0 The melting point of water is 0 o<\/sup>C, and its boiling point is 100 o<\/sup>C.\u00a0 Note that water is a much smaller molecule than previous examples carbon tetrachloride and acetone, yet has higher melting and boiling point due to the relatively strong intermolecular hydrogen bonding experienced by water molecules.<\/p>\r\n\"\"\r\n

    Figure 8.3 Hydrogen Bonding between Water Molecule.\u00a0 Label 1 indicates hydrogen bonds as they are typically depicted using dotted lines. \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 By User Qwerter at Czech wikipedia: Qwerter. Transferred from cs.wikipedia to Commons by sevela.p. Translated to english by by Michal Ma\u0148as (User:snek01). Vectorized by Magasjukur2 - File:3D model hydrogen bonds in water.jpg, CC BY-SA 3.0, https:\/\/commons.wikimedia.org\/w\/index.php?curid=14929959<\/div>\r\n
    <\/div>\r\n
    \u00a0Melting and boiling points for ionic compounds are generally much higher than for molecular compounds because the strongest force between any two particles is the ionic bond, in which two ions of opposing charge are attracted to each other.\u00a0 These i<\/strong>onic interactions<\/span><\/strong><\/span>\u00a0between particles are another type of intermolecular interaction.\u00a0 In the crystal lattice of an ionic compound, each cation is attracted to the anions surrounding it in all directions, and each anion is attracted to all of its surrounding cations.\u00a0 It takes a large amount of energy to overcome these numerous strong attractions simultaneously, so ionic substances typically have high melting and boiling points. Sodium chloride (Figure 8.4 \"Sodium Chloride\") is an example of a substance whose particles experience ionic interactions, with a melting point of 801 o<\/sup>C and a boiling point of 1413\u00a0o<\/sup>C.<\/div>\r\n\"\"\r\n
    Figure 8.4 Ionic crystal lattice in sodium chloride.\u00a0 Image is Public Domain.<\/div>\r\n

    Substances with the highest melting and boiling points have covalent network bonding<\/span><\/strong><\/span>.\u00a0\u00a0 In these substances, all the atoms in a sample are covalently bonded to other atoms; the entire sample is essentially one large molecule, which cannot truly melt or boil, but can decompose.\u00a0 Many of these substances are solid over a large temperature range because it takes so much energy to disrupt all of the covalent bonds at once. One example of a substance that shows covalent network bonding is diamond, shown in Figure 8.5,\u00a0 which is a form of pure carbon. At temperatures over 3,500\u00b0C, diamond finally vaporizes into gas-phase atoms.<\/p>\r\n\"\"\r\n

    Figure 8.5 Covalent network bonding in diamond, Image from https:\/\/en.wikipedia.org\/w\/index.php?title=Allotropes_of_carbon&oldid=854366718, accessed 8\/15\/18<\/div>\r\n

    The phase that a substance adopts at a given temperature and pressure depends on the type and magnitude of the intermolecular interactions the particles of a substance experience. If the intermolecular interactions are relatively strong, then a large amount of energy\u2014in terms of temperature\u2014is necessary for a substance to change phases. If the intermolecular interactions are weak, a low temperature is all that is necessary to move a substance out of the solid phase.\u00a0 The strengths of the individual attractions fall in this order: ionic interaction >>hydrogen bonding > dipole-dipole interaction>>London dispersion forces.\u00a0 In future chapters, this concept will be used to predict and explain relative melting and boiling points for organic compounds.<\/p>\r\n\r\n

    \r\n

    Example 1<\/h3>\r\n

    What intermolecular forces besides dispersion forces, if any, exist in each substance? Are any of these substances solids at room temperature?<\/p>\r\n\r\n

      \r\n \t
    1. potassium chloride (KCl)<\/li>\r\n \t
    2. ethanol (C2<\/sub>H5<\/sub>OH)<\/li>\r\n \t
    3. bromine (Br2<\/sub>)<\/li>\r\n<\/ol>\r\n

      Solution<\/p>\r\n\r\n

        \r\n \t
      1. Potassium chloride is composed of ions, so the intermolecular interaction in potassium chloride is ionic forces. Because ionic interactions are strong, it might be expected that potassium chloride is a solid at room temperature.<\/li>\r\n \t
      2. Ethanol has a hydrogen atom attached to an oxygen atom, so it would experience hydrogen bonding. If the hydrogen bonding is strong enough, ethanol might be a solid at room temperature, but it is difficult to know for certain. (Ethanol is actually a liquid at room temperature.)<\/li>\r\n \t
      3. Elemental bromine has two bromine atoms covalently bonded to each other. Because the atoms on either side of the covalent bond are the same, the electrons in the covalent bond are shared equally, and the bond is a nonpolar covalent bond. Thus, diatomic bromine does not have any intermolecular forces other than dispersion forces. It is unlikely to be a solid at room temperature unless the dispersion forces are strong enough. Bromine is a liquid at room temperature.<\/li>\r\n<\/ol>\r\n<\/div>\r\n
        \r\n
        \r\n

        Skill-Building Exercise<\/h3>\r\n

        What intermolecular forces besides dispersion forces, if any, exist in each substance? Are any of these substances solids at room temperature?<\/p>\r\n\r\n

          \r\n \t
        1. \r\n
          \r\n

          methylamine (CH3<\/sub>NH2<\/sub>)<\/p>\r\n\r\n<\/div><\/li>\r\n \t

        2. \r\n
          \r\n

          calcium sulfate (CaSO4<\/sub>)<\/p>\r\n\r\n<\/div><\/li>\r\n \t

        3. \r\n
          \r\n

          carbon monoxide (CO)<\/p>\r\n\r\n<\/div><\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n

          \r\n
          \r\n
          \r\n

          Concept Review Exercise<\/h3>\r\n
            \r\n \t
          1. \r\n
            \r\n

            What types of intermolecular interactions can exist in compounds?<\/p>\r\n\r\n<\/div><\/li>\r\n<\/ol>\r\n<\/div>\r\n

            \r\n

            Answer<\/h3>\r\n[reveal-answer q=\"390815\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"390815\"]1. polar and nonpolar covalent bonding, ionic bonding, dispersion forces, dipole-dipole interactions, and hydrogen bonding[\/hidden-answer]\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
            \r\n
            \r\n

            Key Takeaways<\/h3>\r\n