Notes
2.5 Lone Pairs: Octet Rule
Sections
Lone Pairs: Octet Rule
Exceptions to the octet rule
- second-row elements can have fewer electrons than an octet (NOT more because they only have s and p orbitals)
- third-row elements can have more bonds than allowed by the octet rule because they have d orbitals
Carbene
Neutral molecule in which a carbon is surrounded by only six valence electrons
Sulfur
Neutral bonding configuration with two bonds and two lone pairs OR with six bonds
Phosphorus
Neutral bonding configuration with three bonds and one lone pair OR with five bonds
Determining the number of lone pairs on atoms:
For a neutral atom, subtract number of bonds from number of valence electrons
For a charged atom, a positive charge is from a lone pair that has replaced a bond and a negative charge is from a bond that has replaced a lone pair
Full Text
Here, we will learn how to assign lone pairs to atoms that do not form an octet.
Write that exceptions to the octet rule come in two forms.
Second-row elements can have fewer electrons than an octet (but not more).
Third-row elements can have more bonds than are allowed by the octet rule because they have access to d orbitals, whereas the second-row elements only hold electrons in s and p orbitals.
For the first drill…
Write that a carbene is a neutral molecule in which a carbon is surrounded by only six valence electrons.
So how many lone pairs are on each atom in dichlorocarbene, formula CCl2?
Show that dichlorocarbene has carbon as a central atom, which is bonded to two chlorine atoms.
Each of the chlorine atoms has three lone pairs for a neutral configuration.
Because the two chlorine atoms only form one bond, the carbon atom cannot have four bonds. We must figure out how to use lone pairs to keep the carbon atom neutral.
We know, from the periodic table, that carbon has 4 valence electrons.
Subtract from this, the number of bonds: 2.
From this, we get 2 nonbonding electrons.
So, draw one lone pair on the carbon atom.
For the next drill…
When depicting the line structures of compounds containing sulfur, the sulfur atom can sometimes be shown with varying numbers of bonds.
Draw the line structure for sulfuric acid, H2SO4, in a way that minimizes all formal charges. How many lone pairs are on the sulfur atom?
Since sulfur is below oxygen in the periodic table, we can predict that like oxygen, it would be neutral with two bonds and two lone pairs; however, this is only applicable when sulfur is bonded to two atoms. In sulfuric acid, sulfur is bonded to four oxygen atoms…
Show sulfur bonded to 4 oxygen atoms.
On the two oxygens to the right and left, add hydrogens to make hydroxyl groups.
Looking at the other two oxygen atoms, we see that they each need another bond for a neutral configuration.
Therefore, turn the S-O single bonds into S-O double bonds for these two oxygen atoms.
Now, we finish the neutral electron configuration for all four oxygen atoms with two lone pairs on each.
Let's count the number of electrons around the sulfur:
6 bonding electrons
- 0 nonbonding electons.
Equals 6 valence electrons.
We know from the periodic table that sulfur has 6 valence electrons, so the atom is neutral. Thus, we have successfully created a line structure with no formal charges.
Next, redraw the Lewis structure for sulfuric acid—but here, only use single bonds.
The first thing that we should do is redraw the atoms in the same arrangement: sulfur surrounded by oxygens, with two peripheral hydrogens.
Next, draw in the bonds. The double bonds in our first structure will be replaced with single bonds in this structure, but otherwise the bonds will be the same.
In terms of electrons, we now need to assign the two electrons that had formed each of the double bonds to an atom.
What is the more electronegative element, sulfur or oxygen?
Oxygen (it sits higher on the periodic table).
Oxygen has the higher affinity for electrons, so put an additional lone pair on each of the terminal oxygens.
Write that each oxygen has one less bond than its neutral bonding configuration.
Indicate that these terminal oxygens each have a -1 charge.
What is the formal charge on the sulfur atom in this case?
On the sulfur, there are four single bonds.
Write that it has two fewer bonds than the neutral bonding arrangement.
And it has no unshared electrons (because oxygen is more electronegative), so indicate that sulfur has a charge of +2.
Why bother to show the sulfur-oxygen bonds as single bonds?
Because the p orbitals that form the double bonds have inefficient overlap due to their sizes:
The 3p orbital from sulfur and 2p orbital from oxygen overlap inefficiently.
For the last drill…
Let's draw the lone pairs on the nitrogen, oxygen, and sulfur atoms in the compound esomeprazole, trade name Nexium, which reduces the formation of stomach acid, and is, therefore, used to manage gastroesophageal reflux disease.
Although the structure contains carbon and hydrogen atoms (in addition to nitrogen, oxygen, and sulfur), we do not need to assign any lone electron pairs to carbon or hydrogen atoms since they are only surrounded by bonds in their bonded, neutral configurations.
First, consider the neutral bonding pattern of nitrogen:
Three bonds + 2 nonbonding electrons.
How many bonds are around each nitrogen atom?
Three.
Therefore, we add one unshared electron pair to each of the three nitrogen atoms.
How about the neutral bonding pattern of oxygen?
2 bonds + 4 nonbonding electrons.
The oxygen atoms each have 2 bonds.
So add two lone pairs onto each atom to complete the bonded configuration.
How about sulfur?
It's in the same group as oxygen, so it has 6 valence electrons.
How many bonds and nonbonding electrons?
Take the number of valence electrons: 6.
And subtract the number of bonds: 4.
Equals: 2 nonbonding electrons – add them (1 lone pair) to the sulfur atom.