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Jump to one of the following Lewis structures' parts:   Choose One Introduction to Lewis Structures Part 1- Basics Part 2- Multiple Bonding and Polyatomic Ions Part 3- Resonance Part 4- Hypervalency Part 5- Formal Charge I Part 6- Formal Charge II Part 7- Electron Deficiency Part 8- Oxoacids Part 9- Quizzes

Now, on to more advanced Lewis structures!

Pull up the Periodic Table if you need one.

How about CO? Let's tackle this just like the previous examples and hit it step by step until we hit a snag. You do realize that we'll hit a snag since we're in a new section, right?

Here's a current list of the rules.

CO
Your turn- enter your answer in the first box and hit "Verify" to see how you did.
1.	Total number of valence electrons.
2. Draw the skeleton structure. 3. Connect the atoms by making bonds. 4. Satisfy the octet rule for the surrounding atom.

Ah, hah! The next step is to satisfy the octet rule for the central atom, but the problem is that you have 1 pair of electrons left when you need 3 additional pair about the carbon atom. Sure, you've satisfied the octet rule for oxygen (remember, even though we only have two atoms, C is considered the central atom for building purposes since it is the least electronegative of the two) but what about carbon? It will only have 4 electrons once you drop that remaining pair of electrons about it and you need 8. Looks like we need a new rule.

Upon distribution of all of the valence electrons, if the central atom does not satisfy the Octet Rule, "borrow" lone pairs of electrons from one or more surrounding atoms and create multiple bonds until the Octet Rule is satisfied.

Two atoms sharing a pair of electrons is a single bond. Two atoms sharing 2 pairs of electrons is a double bond. And two atoms sharing 3 pairs of electrons is a triple bond.

Le't see it in action to make in clearer.

With CO, oxygen has a full octet but carbon is short. What happens if we take a lone pair from oxygen and share it with carbon? Using the structure we finished with in step 4 above-

Up until now, we've stuck with single bonds, but carbon and oxygen are sharing 4 electrons or 2 pairs of electrons so it's a double bond. Do we have a full octet around carbon? Nope. There are only 6 and so we are 2 short. So, grab another lone pair from oxygen and share it with carbon to make a triple bond.

How did I decide where to put the lone pairs of electrons about the atoms? Is there a rule that you need to know? Nope. I could have used any of following structures-

Another way to refer to bonds is "bond order". A single bond has a bond order equal to 1. The bond order of a double bond is 2. And, the bond order for a triple bond is 3. More formally, bond order is defined as the number of pairs of electrons shared among the two atoms in the bond.

A future revision will include some data comparing bond lengths and single, double, and triple bond. For now, consult your textbook.

Let's modify a rule to our growing list for drawing Lewis structures.

Let's modify our list of rules.

Now that you know the rule, how about another molecule? Say, O₂?

O2
Your turn- enter your answer in the first box and hit "Verify" to see how you did.
1.	Total number of valence electrons.
2. Draw the skeleton structure. 3. Connect the atoms by making bonds. 4. Satisfy the octet rule for the surrounding atom. 5. Try to satisfy the octet rule for the central atom. 6. Create a double bond to satisfy the octet rule for the central atom.

In a previous chemistry class or perhaps in the book you are using, you may have seen the Lewis structure for O₂ and other similar compounds (those with only 2 lone pairs of electrons about one or more atoms) drawn in slightly different ways-

or

The first one is the one that I've drawn. The second one is one that you may have seen before. The reason I don't like the second one (the one with the 2 lone pairs of electrons distributed symmetrically about the atoms) is because we are not talking about molecular shape. There's an implication in the second display which I don't think is appropriate at this stage. Shape can only be determined when Lewis structures are used in conjunction with another theory like VSEPR.

Just like the Lewis structure for CO, there are other similar ways to draw the one for O₂. I'll leave it up to you to extrapolate from the various CO Lewis structures.

Let's move on to polyatomic ions. An ion is created when one or more electrons is/are removed from a chemical species. It could be an atom which would then be a monatomic ion. Or, it could be a polyatomic molecule which would result in a polyatomic ion. Just to reiterate- we're talking about the addition or removal of electrons.

This causes problems right from the beginning for Lewis structures. How the heck do you calculate the number of valence electrons? Well, we modify our initial rule to account for ions. If the chemical species has a positive charge (resulting from the removal of electrons), the number of valence electrons is decreased by the amount of charge. If it's a negative charge (resulting from the addition of electrons), the number of valence electrons is increased by the amount of charge.

We need to modify the rule for calculating the total number of valence electrons for the molecule. If needed, you can see the current list.

We need to modify the rules for calculating the total number of electrons for the molecule. Here's the current list of rules.

Let's apply this to NO¹⁺:

NO1+
Your turn- enter your answer in the first box and hit "Verify" to see how you did.
1.	Total number of valence electrons.
2. Draw the skeleton structure. 3. Connect the atoms by making bonds. 4. Satisfy the octet rule for the surrounding atom. 5. Try to satisfy the octet rule for the central atom. 6. Create a triple bond to satisfy the octet rule for the central atom.

Did you notice the difference in the final Lewis structure for NO¹⁺ from that of all the other structures? When you draw one for a polyatomic ion, just like those for ions, you need to indicate the charge. All you need to do is draw brackets around the entire structure and then indicate by a superscript the total charge on the ion.

How about one more before moving on to even more advanced Lewis structures? Take a shot at the cyanide ion, CN^1-.

CN1-
Your turn (press "Reveal!" to check your answer):
1.	Determine the total number of valence electrons.
2. Draw the skeleton structure. 3. Connect the atoms by making bonds. 4. Satisfy the octet rule for the surrounding atom. 5. Try to satisfy the octet rule for the central atom. 6. Create a triple bond to satisfy the octet rule for the central atom.