Chemical Bonding
A chemical bond is a lasting attraction between atoms, ions or molecules that enables the formation of chemical compounds. The bond may result from the electrostatic force of attraction between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds. The strength of chemical bonds varies considerably; there are strong bonds or primary bonds such as covalent, ionic and metallic bonds, and weak bonds or secondary bonds such as dipole-dipole interactions, the London dispersion force and hydrogen bonding.
Covalent bonds form between non-metal atoms, with high electronegativities. Because the electronegativities are not widely different, there is no single atom that pulls the electrons to itself, therefore, covalent bonds involve sharing a pair of electrons, each atom ends up with a complete outer shell because of the shared electrons.
An example of a covalent bond with electrons shared between Carbon and Hydrogen (resulting in Methane).(Source: Wikipedia-CC BY-SA 3.0)
A single bond between two atoms corresponds to the sharing of one pair of electrons. A double bond has two shared pairs of electrons. A triple bond consists of three shared electron pairs. Quadruple and higher bonds are very rare and occur only between certain transition metal atoms.
An illustration of double and triple bonds in carbon dioxide and carbon monoxide respectively. (Source: Wikipedia-CC BY-SA 3.0)
Non-polar covalent bonds occur when the electronegativity difference between the bonded atoms is small, typically 0 to 0.3. Non-polar covalent bonds are often immiscible in water or other polar solvents, but much more soluble in non-polar solvents such as hexane.
A polar covalent bond occur when the electronegativity difference between the two atoms is higher, typically 0.3 to 1.7. This difference creates an imbalance of charge as the electrons spend more time in time in the highly electronegative atom. Such bonds have significant ionic character. When the electronegativities are only moderately different (0.3 to 0.9) these bonds give rise to dipole–dipole interactions.
Ionic bonds occur between a metal and a non metal, both significantly different electronegativity (typically >1.7). Ionic bonding leads to separate positive ion (cation) and negative ions (anion). Electrons are transferred from metal (low electronegativity) to non-metal (high electronegativity) atoms. A typical feature of ionic bonds is the formation of a 3-dimensional crystal lattice structure. Ionic bonds are held strongly together by the attraction between the negative and positive ions.
An illustration of how ionic bonds are formed, one atom losing its electron(s) and the other gaining the electron(s). (Source: Wikipedia-CC BY-SA 3.0)
Metal atoms generally have lower electronegativity. Pure metals do not consist of metal atoms, but of closely packed cations suspended in a “sea” of free electrons which can move between the cations (+ ions). The free movement or delocalization of bonding electrons leads to classical metallic properties such as luster. (Lustre or luster is the way light interacts with the surface of a crystal, rock, or mineral, generally denotes radiance, gloss, or brilliance). The electrons act like a “glue” that holds the positive nuclei together. Free electrons transmit kinetic energy rapidly, thus metals are excellent conductors of heat. The free electrons can also move rapidly in response to an electric field, making metals excellent conductors of electricity. Individual atoms are held loosely to other atoms, so atoms slip easily past one another, so metals are ductile (Ductility is a measure of a material's ability to undergo significant plastic deformation before rupture or breaking). The atoms in metals have a strong attractive force between them. Much energy is required to overcome it. Therefore, metals often have high boiling points. Metals are insoluble in water or organic solvents unless they undergo a reaction with them.
As the term suggests, intermolecular forces (aka Van der Waals forces) are forces of attraction or repulsion that occur between molecules. These should be differentiated from intramolecular forces, which are forces that occur between atoms within the same molecule. These forces tend to be weak, as compared to intramolecular forces.
In polar molecules (such as water), the negative poles are attracted to positive poles of adjacent polar molecules. In a liquid, molecules can orient themselves to maximize attraction and minimize repulsion.
Named after Fritz London (1900 - 1954), London forces are caused by the attraction of electrons to the nuclei of adjacent molecules. It is caused by a brief shift in electrons from one side of an atom/molecule to another, called a momentary dipole. The greater the number of electrons in the atoms of a molecule, the 'stronger' the London forces. Stronger London forces are responsible for increasing the boiling points of molecules, where the higher the number of electrons the higher the boiling point. Among molecules that have the same number of electrons, other factors may also influence their boiling points. For example, both Br2 and ICl have a total of 70 electrons, however ICl is a polar molecule, so in addition to London forces, dipole-dipole forces make ICl have a higher boiling point (97oC) compared to Br2 (59oC).
A hydrogen bond occurs when a hydrogen ion (a proton) is shared among electrons in neighboring molecules. For this to happen, the hydrogen atom must be bonded to a highly electronegative atom such as Nitrogen, Oxygen or Fluorine. There must be at least one lone pair of electrons on the atom bonded to the hydrogen. Hydrogen bonding is in fact a special type of dipole-dipole attraction that only occurs between hydrogen ions and either F, N or O atoms. In water, hydrogen bonding is responsible for the high melting and boiling points when compared to other substances of similar characteristics such as number of electrons. Hydrogen bonding in water is also responsible for other characteristics such as surface tension, ice floating on the surface of lakes/rivers, capillary action and the shape of the meniscus.
Liquids with small intermolecular forces (such as gasoline, ethanol etc.) are volatile i.e. they tend to evaporate quickly.
Examples of weak intermolecular forces.
The intermolecular forces (such as those within the water molecules (cohesive forces)) create a film on the surface of water that resist disruption by external forces. This makes the surface of water act as if it was an elastic membrane. Insects are able to stand on the surface of water without sinking, making it easier for them to catch their prey. Surface tension is also responsible for the shape of water droplets making them look more spherical, an indication that the forces attracting the water molecules to each other are stronger than the forces between the water molecules and the environment.
An illustration of surface tension. (Source: Wikipedia-CC BY-SA 3.0)
Menisci can be classified as either concave or convex. The shape of the menisci is influenced by the intermolecular forces between the liquid molecules, and between the liquid molecules and the molecules of wall of the container.
Two types of menisci; A: Concave as observed in water, B: Convex as observed in Mercury. The line indicates the volume reading. (Source: Wikipedia-CC BY-SA 3.0)
Capillary action occurs when a liquid, such as water, is pulled up a narrow tube by the intermolecular forces between the water molecules and the tube. Capillary action is an important biological process in both plants and animals.
Download a pdf copy of Unit 2: Chemical Bonding for offline use.
C$0.99