1. Sum up the total number of valence electrons for the elements in the molecule. If the molecule is charged, subtract one electron for each positive charge or add one electron for each negative charge.
2. Write the structure for the molecule with a pair of electrons (or a dash) between each atom.
Groups of atoms will usually have the less electronegative atom surrounded by atoms having greater electronegativity. Never place a hydrogen atom in the center since it can only form one bond.
Think about the valence of each atom and make sure that these valences are not exceeded except in the case of ions. Ions always have more or less bonds than the normal uncharged atom.
The common valence or number of bonds formed for some common atoms are...
H = 1
O and S = 2
N = 3
C = 4
F, Cl, Br and I = 1
3. Place electrons around the outer atoms to fill their outer shells.
Most atoms require eight electrons ("octet rule") so they will resemble an inert gas. Elements in groups IA to IIIA do not follow the octet rule and have less than eight electrons in the final formula.
4. Subtract the number of electrons used so far from the total calculated in step 1 and place these remaining electrons on the central atom or atoms.
5. If the central atom ends up with less than 8 electrons, then it probably forms a multiple bond with an adjacent outer atom.
6. Finally you should calculate the formal charge on the atoms.