Lecture Notes:

Covalent Bonding

• Molecular Compounds
  • Compounds that contain only covalent bonds are called molecular compounds
  • Molecular compounds have a different structure and very different properties from ionic compounds
  • Molecular compounds have lower melting points, lower boiling points, and are very poor conductors of electricity compared with ionic compounds
  • They exist as individual molecules rather than as part of a crystal lattice: each formula unit represents a single entity that is not strongly bonded to other molecules of the same type
  • Ionic compounds by contrast do not exist alone as their formula units might seem to imply: they exist only as large groups in a crystalline arrangement
  • Molecular compounds have lower melting and boiling points precisely because they consist of separate molecular units; the atoms are still bound strongly together but the molecules are not bound to one another
• The Nature of Covalent Bonding
  • Covalent bonding is characterized as a sharing of electrons to achieve stability and complete octets; contrast this with ionic compounds which are said to exchange electrons
  • Ionically-bonded atom are held together by electromagnetic forces; covalently bonded atoms are held together by the same forces in a slightly different arrangement
  • Ionic bonds form between ions of opposite charge: the whole atom (or group of atoms in the case of polyatomic ions) have a charge
  • Covalent bonds form because a region of higher electron density can be found between the bonding atoms
  • Electrons in both atoms are attracted to both nuclei; the repulsive forces between particles of like charge is overcome by the build-up of a region of negative charge between the bonding atoms
  • The two positive nuclei are both attracted to this region of negative charge; in this way electrons can be said to be shared
  • One more reason this type of bonding occurs is that electrons from separate atoms can become spin-pairs when the atoms are bound together; that is, one will be spin up and the other spin down; as you know, this combination is favorable for electrons in the same orbital
• Properties of Bonds (No 007 jokes, please)
  • Bond Length
    • The distance between two bonded atoms is called the bond length
    • This is actually an average length because the atoms are constantly in vibrational motion
    • Bond length is inversely proportional to the energy of the bond; that is the shorter the average bond length, the more energy is required to break that bond
    • The higher the bond energy, the more energy is released when a bond forms
    • Bond energy is a good indicator of the force of attraction between two atoms: the shorter the bond, the stronger it is
    • It is easier to break the bonds between two large atoms than between two small atoms, simply because the small atoms can be closer together
  • Electronegativity and Bonding
    • So far we have been speaking as if there were exactly two kinds of bonds between atoms; this is not strictly true
    • All atomic bonds actually fall on a spectrum between the extremes of ionic and covalent bonds
    • Electronegativity allows you to decide the polarity of a bond
    • If a bond is very polar, like the bond between Na and Cl in salt, it is said to be wholly ionic
    • If a bond is very nonpolar, like the bond between two O atoms in O₂, then it is said to be wholly covalent
    • Many bonds fall between these extremes and belong to molecules which can, as a whole, be polar or nonpolar
    • The reason behind this is that atoms with a high electronegativity value attract electron density to themselves more than atoms with a low value
    • Water is an excellent example: the oxygen atom has an electronegativity of 3.5 on the Pauling scale, the two hydrogen atoms have an electronegativity of 2.1; this difference is not great enough to make the bonds between these atoms ionic but it is great enough to cause a separation of charge
    • Separation of charge means that there is a partial positive charge on the atom from which electron density has been ‘stolen’ and a partial negative charge on the atom which ‘stole’ the electron density
    • Generally, there are three broad categories of atomic bonding:
      1. nonpolar covalent: this is a bond with no polar characteristics; an example is the bond between H atoms in H₂; the electronegativity difference between the atoms in a nonpolar covalent bond lies between 0.0 and 0.4
      2. polar covalent: this is a bond with a definite separation of charge; a partial positive charge can be found on the less electronegative atom and a partial negative charge can be found on the more electronegative atom; the electronegativity difference between the atoms in a polar covalent bond lies between 0.4 and 2.1
      3. ionic: this is a bond between atoms or groups of atoms that each have their own charge; the electronegativity difference between the atoms in an ionic bond lies between 2.1 and 4.0
  • Intermolecular Bonding
    • Weak bonds form between molecules, especially between molecules that have polar character
    • This happens because of the attractive forces caused by the separation of charge in these molecules; they are said to have a dipole
    • One important kind of intermolecular bonding is called hydrogen bonding; this occurs between molecules that contain hydrogen bonded to an electronegative atom such as oxygen, nitrogen, or fluorine
    • This works because H is such a small atom that it represents a high charge density and it can come closer to atoms in other molecules than larger atoms can