By now you should already know how elements in groups 1, 2 and 13 through 18 behave when it comes time to form ions. For example, Fluorine tends to gain one electron in order to have the electron configuration of Neon (and have a -1 charge).

The formation of ions is only one way in which atoms can form

I borrowed this image from HyperPhysics.

bonds. Ionic compounds form between atoms or groups of atoms that carry positive and negative charges. The positively charged particles are attracted to the negatively charged particles and in much the same way that a balloon rubbed on your hair will stick to the wall, ions form compounds. Molecular compounds form due to a completely different kind of interaction. Instead of losing or gaining electrons to form ions the atoms in molecular compounds share their electrons with each other.

Atoms become ions in a process of energy minimization: they are more chemically stable when they have the same number of valence electrons as a noble gas (group 18). Another way for atoms to minimize their energy and have full valence shells is for them to share electrons with each other. For example, in H₂O the bond between a Hydrogen atom and the Oxygen atom contains two electrons and Hydrogen needs two electrons to fill its valence shell. Oxygen needs two electrons to fill out its valence shell and it borrows one from each Hydrogen atom. This type of bond is called a covalent bond.

Given a molecular formula, there are some steps to follow in drawing a Lewis Dot Diagram for the molecule. Conventionally, electrons are considered in pairs (an exception is a molecular radical with an odd number of electrons). Electron pairs in a bond are represented as a line between the atomic symbols:
H—O—H. Non-bonding pairs (also called lone pairs) are represented by a pair of dots near the atomic symbol they belong to (for example, :N—F₃; the three single bonds to three F atoms are represented as one line for the sake of simplicity in this text). Lone pair electrons are valence shell electrons that do not participate in bonding.

1. Count the number of valence electrons for each of the atoms in the molecule and calculate the total.
2. If the molecule has a negative charge add enough electrons to account for it. If the molecule has a positive charge, subtract the number of electrons needed to create that charge.
3. Organize the atoms around a central atom. Hydrogen is never the central atom. Atoms other than the central atom are referred to as ligands.
4. Draw a single line between each ligand atom and the central atom to represent bonds between those atoms. Each bond requires one electron pair.
5. Add lone pairs of electrons to each atom as required to make an octet, giving each atom a noble gas electron configuration. Remember not to exceed the number of electrons you calculated earlier!
6. Redistribute lone pairs as double- or triple-bonds if any atoms do not have an octet.
7. Formal charge will determine if your electron distribution is correct. Calculate it by comparing the number of electrons that belong to an atom to the number of valence electrons for that atom. Do this as follows:
   1. Count the total number of bonding electrons attached to the atom. Call this total b.
   2. Count the total number of lone pair electrons. Call this total n.
   3. Formal charge (F) is defined as follows: F = V – (n + b/2).
   Formal charge should nearly always be zero for every atom in a molecule. The exception to this rule is when a molecule has an overall charge. Total formal charge for a molecule should equal the overall charge of the molecule. Minimize formal charges so that they are all as small as possible and so that as few atoms as possible have formal charges.