Basic Concepts of Chemical Bonding

Drawing Lewis Structures

The general procedure...

1. Sum the valence electrons from all atoms

• Use the periodic table for reference
• Add an electron for each indicated negative charge, subtract an electron for each indicated positive charge

2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond

• You may need some additional evidence to decide bonding interactions
• If a central atom has various groups bonded to it, it is usually listed first: CO₃^2-, SF₄
• Often atoms are written in the order of their connections: HCN

3. Complete the octets of the atoms bonded to the central atom (H only has two)

4. Place any leftover electrons on the central atom (even if it results in more than an octet)

5. If there are not enough electrons to give the central atom an octet, try multiple bonds (use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds)

Draw the Lewis structure of phosphorous trichloride (PCl₃)

This is an example of a central atom, P, surrounded by chlorine atoms

1. We will have 5(P) plus 21 (3*7, for Cl), or 26 total valence electrons

2. The general symbol, starting with only single bonds, would be:

3. Completing the octets of the Cl atoms bonded to the central P atom:

4. This gives us a total of (18 electrons) plus the 6 in the three single bonds, or 24 electrons total. Thus we have 2 extra valence electrons which are not accounted for. We will place them on the central element:

5. The central atom now has an octect, and there is no need to invoke any double or triple bonds to achieve an octet for the central atom. We are finished.

Draw the Lewis structure for the NO+ ion

1. We will have 5 (N) plus 6 (O) minus 1 (1+ ion), or 10 valence electrons

2. The general structure starting only with single bonds would be:

3. Completing the octet of the O bonded to N:

4. This gives us a total of 6 plus 2 for the single bond, or 8 electrons. There are 2 unaccounted for electrons and we will place them on the N:

5. There are only 4 atoms on the N atom, not enough for an octet, so lets try a double bond between the N and O:

The oxygen still has an octet, but the N only has 6 valence electrons, so lets try a triple bond:

Each atom now has a valence octet. We are finished.

In some cases we can draw several different Lewis structures which fulfill the octet rule for a compound. Which one is the most reasonable?

One method is to tabulate the valence electrons around each atom in a Lewis structure to determine the formal charge. The formal charge is the charge that an atom in a molecule would have if we considered each atom to have the same electronegativity in a compound.

To calculate formal charge, assign electrons to their respective atoms as follows:

1. All of the unshared electrons are assigned to the atom on which they are found
2. The bonding electrons are divided up equally between each atom involved in the bond
3. The number of valence electrons expected in the isolated atom is compared to the actual number of electrons assigned from the Lewis structure:

Example: Carbon Dioxide (CO₂)

Carbon has 4 valence electrons

Each oxygen has 6 valence electrons, therefore our Lewis structure of CO₂ will have 16 electrons:

One way we could draw the Lewis structure is:

Another way we could draw the Lewis structure is:

Both structures fulfill the octet rule. But what are the formal charges?

Which structure is correct? In general, when several Lewis structures can be drawn the most stable structure is the one in which:

• The formal charges are the smallest
• Any negative charge is found on the most electronegative atom

In the above case, the second structure is the one with the smallest formal charges (i.e. 0 on all the atoms).

• Furthermore, in the first possible Lewis structure the carbon has a formal charge of 0 and one of the oxygens it is bonded to has a formal charge of +1.
• Oxygen is more electronegative than Carbon, so this situation would seem unlikely.

It is important to remember that formal charges do not represent the actual charges on the atoms. Actual charges are determined by the electronegativity of the atoms involved.

1996 Michael Blaber