Foundations of Chemistry. Philippa B. CranwellЧитать онлайн книгу.
dipole showing charge distribution; (c) neighbouring chlorine molecule with an induced dipole, showing charge distribution.
This is because the shift of electrons generating a δ− charge forces the electrons in a nearby bond to be repelled, so a dipole is formed in the nearby molecule, as shown in Figure 2.30c. London dispersion forces are reasonably weak because they are short‐lived, but they are important nevertheless.
Instantaneous dipoles (or dispersion forces) increase with increasing polarisability of the molecule. The more readily polarisable the molecule, the larger the instantaneous dipole and induced dipole. Polarisability is a measure of how easily the charge distribution in an atom or molecule can be distorted by the application of an external electrical field or charge. The greater the number of electrons, the more readily polarisable the molecule is, so instantaneous dipoles increase with increasing molecular mass. In addition, the larger the surface area of a molecule or the larger the area of possible contact between two molecules, the stronger the intermolecular forces.
Instantaneous dipoles are responsible for the very weak intermolecular forces formed between noble gas atoms. The noble gases consist of monatomic molecules: single atoms of neon, argon, etc. Clearly, there are no intramolecular forces here, as there are no bonds, and the intermolecular forces are extremely weak. Instantaneous dipoles are formed by the random movement of electron density from one side of the atom to another, which then induces a dipole in a neighbouring atom, as shown in Figure 2.31. Helium has only two electrons, so the size of the dipole is very small. The heavier noble gases possess more electrons and so have larger dipoles.
Figure 2.31 (a) Helium atom (Z = 2) showing even distribution of electrons; (b) helium atom with instantaneous dipole due to temporary movement of charge; (c) induced dipole in a neighbouring helium atom.
A monatomic molecule is composed of just one atom. A diatomic molecule such as Cl2 has two atoms. A triatomic molecule such as H2O has three atoms. A polyatomic atom has several atoms.
2.4.3 Hydrogen bonding
The final type of interaction is hydrogen bonding, which is the strongest type of intermolecular force and has about one‐fifth the strength of a typical covalent bond. Hydrogen bonding is actually a special type of dipole–dipole interaction. Hydrogen bonds are formed between molecules that contain a hydrogen atom bonded to a small, strongly electronegative element such as nitrogen, oxygen, or fluorine.
Hydrogen bonding can occur when a hydrogen atom in a molecule is bonded to a strongly electronegative element such as nitrogen, oxygen, or fluorine. A hydrogen bond is formed between the H atom of one molecule and the N, O, or F atom of a neighbouring molecule.
The hydrogen atom is very small and has only one electron in its 1s orbital. When bonded to an electronegative element such as nitrogen, N, oxygen, O, or fluorine, F, a dipole exists, and the hydrogen atom adopts a partial positive charge (δ+) while the other atom becomes slightly negatively charged (δ−). The partial positive charge is therefore concentrated over a very small volume and makes the hydrogen atom strongly polarising. This allows the hydrogen atom from one molecule to attract electron density from a small electronegative atom (such as nitrogen, oxygen, or fluorine) in a neighbouring molecule, and the hydrogen atom becomes sandwiched between the two more electronegative atoms. This can be seen in Figure 2.32, which depicts the formation of a hydrogen bond between two molecules of water. The highly polarising hydrogen atom in one water molecule (δ+) is able to attract electron density in the form of a lone pair of electrons from a neighbouring oxygen atom (δ−). The lone pair is now ‘shared’ between the hydrogen atom and the oxygen atom of a neighbouring molecule, and a hydrogen bond is formed. Because hydrogen bonding requires both the bonded and the neighbouring atom to be highly electronegative and small, hydrogen bonding generally only takes place between hydrogen and nitrogen, oxygen, and fluorine.
Figure 2.32 (a) Formation of a hydrogen bond between two molecules of water where molecules are in constant motion; (b) formation of hydrogen bonds in ice, where molecules are in a fixed position.
In liquid water, hydrogen bonds are constantly forming and breaking, as in Figure 2.32a, whereas in ice (Figure 2.32b), the molecules are held together in fixed positions by hydrogen bonds. Each oxygen atom can form two hydrogen bonds to neighbouring molecules as it possesses two lone pairs of electrons.
Hydrogen bonds are also formed between different types of molecules. The group of organic molecules known as alcohols contains an —OH bond. Ethanol, C2H5OH, is a well‐known alcohol. Hydrogen bonding can occur between molecules of ethanol (Figure 2.33a). Hydrogen bonding can also occur between molecules of ethanol and molecules of water, as shown in Figure 2.33b.
Figure 2.33 (a) Hydrogen bonding between two ethanol, C2H5OH, molecules; (b) hydrogen bonding between ethanol, C2H5OH, and water molecules.
2.4.4 Summary of strengths of intermolecular forces
Intermolecular forces are the weak forces of attraction formed between simple covalently bonded molecules. There are three types of intermolecular forces:
London dispersion forces: instantaneous dipole–induced dipole forces
Permanent dipole–permanent dipole forces
Hydrogen bonding
The order of strengths of these forces is:
instantaneous–induced dipole forces < permanent dipole–permanent dipole forces < hydrogen bonding. Table 2.2 gives approximate values for the strengths or energies of different types of intermolecular forces and bond types.
Table 2.2 Approximate strengths of different types of bonds and intermolecular forces.
| Bond type/intermolecular force type | Strength (approx.)/kJ mol−1 |
|---|---|
| Ionic bond (e.g. NaCl) | 771 |
| Single covalent bond (e.g. Cl—Cl) | 242 |