Chemical Bonding


We know the structure of the atom. We know that each atom is a compromise between electrostatic attraction between the electrons and the nucleus and electron-electron repulsion. Chemical bonds between atoms must have the same features.

The energy of interaction between the atoms changes with distance between the nucleii. There is an optimal distance for the chemical bond which is where this energy is at a minimum. The minimum energy with respect to the energy of dissociated fragments (r -> infinity) is called the bond energy.


The above picture is for the case of the Hydrogen molecule, but each particular chemical bond has its' own equilibrium distance and its' own bond dissociation energy. The bond lengths of the halogen molecules are used to determine an approximate radius for chemical (covalent) bonding of the halogen atoms. Perfect electron sharing is expected between atoms of the same type, so the bond in this case is perfectly covalent.

Why do atoms for covalent bonds? two major factors:

The 'Lewis Dot' symbology is simple. Draw atoms with their valence electrons only as dots, grouped in four possible pairs around the atom. Fill the four places around the atom as if they were four degenerate orbitals.

Now combine atoms together to form molecules by pairing electrons without changing the total number of electrons. Make an 'octet' around each atom in this way (except Hydrogen which can only support 2 valence electrons and heavy elements which can support 'super-octets' due to unfilled d- and f- orbitals). Each pair of electrons is either shared between two atoms, and is called a Bonding Pair or is entirely owned by a single atom, and called a Lone Pair Sometimes the bonding pairs will be replaced with a line for each pair of electrons. Usually lone pair electronss are left as dots, but they two may be replaced by lines. If more than one pair of electrons is shared between a given pair of atoms, a multiple bond has formed. Draw a solid line for each pair of bonding electrons in the multiple bond. A good Lewis structure has all the electrons paired (this is not possible if the number of valence electrons is odd) and an octet around each atom.
An additional feature of the Lewis structure is the so called Formal Charge, which is assigned to each atom in the molecule. The Formal Charge is defined as the number of valence electrons the neutral atom should have minus the number of lone pair electrons it has, minus one half the number of bonding electrons it shares. The sum of the formal charges on all the atoms in the structure is equal to the total charge of the species. A good Lewis structure has the minimum number and magnitude of formal charges, and avoids having formal charges of the same sign on adjacent atoms.

Triumph of Lewis dot structure:
  • Predicts multiple bonds
  • Explains stoichiometry of covalent molecules
    • Example Lewis structures:

    Try this interactive builder of     Lewis Dot Structures
    See this index of many webites related to Lewis Dot Structures.
    Covalent versus Ionic Bonding
    In our early discussion of chemical compounds, we said that if a non-metal and a metal bond, one or more electrons will be transferred from the metal to the nonmetal and the resulting ions stick by electrostatics. This is an extreme case of unequal sharing of electrons, but leads to the same kind of octet configuration of the atoms involved in the bond. Take LiF, for example:

    How evenly or not the bonding electrons are shared will determine the polarity of the bond and the nature of the interaction. What determines how the electrons are shared is the relative electronegativity (electron greed) of the bonding atoms. The degree of polarity or degree of ionic bonding of any given bond can vary continuosly zero to nearly 100%. We normally say that bonds between atoms with electronegativity difference ( DEN )greater than 1.7 are ionic, although this really means only more than about half ionic in character. Here are some examples of the bonds formed with different electronegativity differences, DEN,between the atoms:
    Bond Energetics
    Since we are treating the chemical bond as largely depending only upon the nature of the two atoms in contact through the bond, perhaps we can use this idea to determine the overall stability of a molecule by adding up its bond energies. This assumes that all chemical bonds between the same pair of atoms of the same type are approximately equal in properties. Namely, in this case, we will assume all C-H bonds take about the same amount of energy to break, regardless of the molecule they are in.

    The hypothetical state of a molecule after all its bonds are broken can be used as a 'reference', just like we used the standard states of the elements as a reference for the Enthalpies of Formation of molecules. Thus the energetics of a chemical transformation can be estimated from the bonds broken and formed in the reaction


    A specific example can be made from our old familiar combustion of methane reaction. We calculated the enthaly change during this transformation before from trditional thermochemcial methods. We can do this agian by using the average bond enthalpies of C-H, C=O, {O=O}, and O-H bonds

    So, the Heat of Formation of new molecule, or the Heat of Reactions of a given transformation can be estimated by using average bond energies and the above thermochemical analysis. This is not as accurate as using directly measured heats of formation (which is not an approximation!) but is sometimes very useful as a starting guess.


    Other average properties of bonds are also useful. For instance, the equilibrium bond length of a given type of bond is usually pretty constant from molecule to molecule. Therefore, average bond lengths can be used to predict parts of the structure of new and unknown molecules.
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