Drawing Lewis structures

G N Lewis recognized that the most stable covalently-bonded species have an even number of electrons. He made a major contribution to the understanding of bonding by proposing that
 
Although more sophisticated bonding theories have developed, Lewis structures are still used by practicing chemists to show the distribution of bonding and nonbonding electrons when predicting structure or reactivity. The Lewis structure of H₂O shows that this molecule has two bonds (four bonding electrons) and four non-bonding electrons.
  
To construct the Lewis structure for a given molecular formula:

Step 1:	Count the total number of  valence electrons available on this group of atoms. For H2O:  2(2H) + 6(1O from Group 16) = 8 For PCl3: 5(P from Group 15) + 3 × 7 (3Cl from Group 17) = 26 Subtract electrons if overall positive; Add electrons if overall negative.
Step 2:	Join atoms by single bonds.  Subtract the total number of electrons used from the Step 1 total.   For H2O:  Subtract 4 because each of two bonds use two electrons (this leaves 4). For PCl3: Subtract 6 because each of three bonds uses two electrons (20 remain)
Step 3:	Complete octets at outer atoms other than H.  Subtract electrons used from the Step 2 total. An octet is complete if there are eight electrons around the atom. These may be bonding and non-bonding. Each single bond to an atom counts as two electrons.                Cl                 | For: Cl – P – Cl Six additional electrons are required at each chlorine. These will be non-bonding. (Subtract 18 from 20 to give 2 remaining) For H-O-H:  No change because the outer atoms are hydrogen (4 electrons remain).
Step 4:	Place the remaining valence electrons at the central atom.   For H2O 4 electrons remaining; place two nonbonding pairs at O as above. For PCl3 2 electrons remaining: place one non-bonding pair at P   Central atoms from the second row of the periodic table have a maximum of eight electrons. Central atoms from other rows have a minimum of eight electrons.