1. Coulomb Forces: A Simplified View of Bonding
Bonds are made by simultaneous coulombic attraction and electron exchange.
When two atoms approach, the electrons of one are attracted by the protons of the other and vice-versa.
Energy is released as the two atoms approach each other. When the atoms get too close together the energy begins to rise again due to repulsions between the two nuclei and the two sets of electrons.

2. This minimum energy is called the bond strength, and the distance between the two nuclei at this point is called the bond length.

3. Ionic and Covalent Bonds: The Octet Rule
Covalent Bonds are based on the sharing of electrons.
If the electrons are not shared equally, a polar covalent (partially ionic) bond is formed, otherwise a pure covalent bond is formed.
Ionic Bonds are based on the transfer of one or more electrons from one atom to another. The resulting cation and anion are electrostatically attracted to each other.

5. Ionic and Covalent Bonds: The Octet Rule
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Ionic and Covalent Bonds: The Octet Rule
The periodic table underlies the octet rule.
Electrons in atoms occupy levels or shells of fixed capacity.
The first has room for 2, the second 8, and the third 16.
Noble gases have 8 valence electrons (Helium 2) and are particularly stable.
Other elements lack octets in their outer electron shells and tend to form molecules in such a way as to create a stable octet arrangement.

6. In pure ionic bonds, electron octets are formed by transfer of electrons.
Alkali metals react with halogens by the transfer of one electron from the alkali metal to the halogen.
Both atoms achieve a noble gas configuration: the alkali metal that of the preceding inert gas, the halogen that of the following inert gas.
IPNa = +119 kcal mol-1
EACl = kcal mol-1
Electostatic Attraction = -120 kcal mol-1
E = kcal mol-1

7. Valence electrons are conveniently indicated by placing dots around the symbol for an element. The letters represent the nucleus and the core electrons, and the dots represent the valence electrons:
Hydrogen can either lose an electron to form an H+ ion, or gain an electron to form a H-, or hydride, ion:

8. In covalent bonds, electrons are shared to achieve octet configurations
Ionic bonds between identical atoms of the same element do not form.
The high ionization potential of hydrogen prevents it from forming ionic bonds with halogens and other non-metallic elements.
Ionic bonds are also unfeasible for carbon since it would require the loss of 4 electrons to achieve the octet of the preceding inert gas, or the gain of 4 electrons to achieve the octet of the following inert gas.

9. In these and similar cases, covalent bonding occurs
In these and similar cases, covalent bonding occurs. Atoms share electrons to achieve a noble gas configuration.
In certain cases, one atoms supplies both of the electrons in the bond:
Often 4 electron (double) and 6 electron (triple) bonds are formed:

10. In most organic bonds, the electrons are not shared equally: polar covalent bonds.
Pure covalent bonds (perfect sharing of electrons) and ionic bonds (complete transfer of electrons) are two extreme types of bonding.
Most bonds lie somewhere between these extremes and are called polar covalent bonds.
Each element can be assigned an electronegativity value which represents its electron accepting ability when participating in a chemical bond.
The larger the difference in electronegativety between two atoms participating in a chemical bond, the more ionic is the bond.
Bonds between atoms of different electronegativity are said to be polar bonds. A partial negative charge is found on the atom of higher electronegativity and an equal but positive charge on the other atom.

11. As a rule of thumb, electronegativity differences less than 0
As a rule of thumb, electronegativity differences less than 0.3 represent pure covalent bonds, from 0.3 to 2.0 polar covalent bonds, and greater than 2.0 ionic bonds.
The separation of opposite charges in polar covalent molecules results in the formation of dipoles:

12. In symmetrical molecules such as CO2 and CCl4, the individual dipoles will cancel and the molecule is left with a zero dipole moment.

13. Electron repulsion controls the shapes of molecules.
The shapes of molecules can be predicted using the VSEPR method.
Bonding and non-bonding electron pairs on the same atom will arrange themselves in three-dimensions to be as far apart as possible.
In the case of 2 electron pairs, as in BeCl2, a linear arrangement results. For 3 electrons pairs, as is in BCl3, a trigonal arrangement results, and in the case of 4 electron pairs, a tetrahedral arrangement occurs: