Chemical Bonding

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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:

Delocalizing electrons over two atoms instead of one lowers the energy of the system.
Atoms with less than filled shell electron configurations can share valence electrons to fill their highest or valence subshell and thus gain quantum mechanical stability. G.N. Lewis counted the valence electrons with dots to show how the tendency to create a filled s and p subshell (8 electrons) influenced molecular stochiometery and structure. He called the tendency to have or share 8 electrons the 'Octet Rule'.

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). Replace all bonding pairs with a single line (non-bonded pairs of electrons or lone pairs are left as two dots). 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. Try to pair all the electrons in the structure (this is not possible if the number of valence electrons is odd).

Triumphs of Lewis dot structure:

Predicts multiple bonds
Explains stoichiometry of covalent molecules

Example Lewis structures:

Another short description of Lewis Dots?

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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Chemical Bonding