Simple Procedure for writing Lewis Structures – Examples #2

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A simple procedure for writing Lewis structures is given in a previous article entitled “Lewis Structures and the Octet Rule”.

Examples for writing Lewis structures following the above procedure are given bellow:

Consider the case of the cyanate ion, NCO^-:

Step 1: The central atom will be the C atom since it is the less electronegative.  Connect the C atom with the N and O atoms with single bonds

           

Step 2: Calculate the # of electrons in π bonds (multiple bonds) using formula (1) in the article entitled “Lewis Structures and the Octet Rule”.

Where n in this case is 3 since NCO^-  consists of three atoms

Where V = (5 + 4 + 6) – (-1)  = 16  

Therefore, P = 6n + 2 – V = 6 * 3 + 2 – 16 = 4      \  there are 4 π electrons in NCO^-   \

2 double   bonds must be added to the structure of Step 1 or 1 triple bond.

Step 3 & 4: One double bond between C and O is added to the structure in step 1 and a second double bond between C and N. Alternatively, 1 triple bond is added either between C and O or between C and N. Unshared electron pairs are added so that there is an octet of electrons around each atom. All the equivalent resonance structures are drawn by delocalizing electron pairs. Therefore, the Lewis structures for NCO^-  are as follows:

Figure 1: Lewis structures for the cyanate ion NCO^-. Because of the extreme charge separation in the resonance form 2, it is considered that it does not contribute much to the ground state of the molecule. Resonance forms 1 and 3 are equivalent and contribute to the ground state of the molecule – there is less charge separation in them in comparison to resonance form 2. The above Lewis structures indicate that the cyanate ion is capable of bonding at either the nitrogen or oxygen to a metal atom, as indeed is known to occur.

Let us consider the case of CO₂:

Step 1: The central atom will be the C atom since it is the less electronegative.  Connect the C atom with the O atoms with single bonds

           

Step 2: Calculate the # of electrons in π bonds (multiple bonds) using  formula (1):

Where n in this case is 3 since CO₂ consists of three atoms.

Where V = (6 + 4 + 6) = 16  

Therefore, P = 6n + 2 – V = 6 * 3 + 2 – 16 = 4      \  there are 4 π electrons in CO₂   \

2 double bonds must be added to the structure of Step 1 or 1 triple bond.

Step 3 & 4: Two double bonds between C and O are added to the structure in step 1. Alternatively, 1 triple bond is added between C and O. Unshared electron pairs are added so that there is an octet of electrons around each atom. All the equivalent resonance structures are drawn by delocalizing electron pairs. Therefore, the Lewis structures for CO₂  are as follows:

Figure 2: Lewis structures for the CO₂ molecule. Resonance form 1 contributes significantly to the ground state of the molecule. Resonance form 2 is not significant – it is not energetically favored – because of the larger charge separation compared to 1.