10.1 OBJECTIVES
Draw Lewis structures of simple molecules and ions.
Assign formal charges correctly.
Draw several valid resonance structures when appropriate.
10.2 BACKGROUND
Being able to draw a clear picture of what a molecule looks like can be very helpful in
several ways. First, it can help you visualize the molecule in your mind, so that it is not
just a formula of symbols and subscripts. Second, it helps us keep track of all the
electrons and bonds in the molecule. Finally, we will use it later to help predict shapes.
Drawing Lewis Structures
There are four basic steps to draw a Lewis Structure:
1. Add up all the valence electrons in the molecule. If you are working with an ion,
add an electron for each overall negative charge, or subtract an electron for each
positive charge. Write the total number of electrons down.
2. Correctly connect all the atoms with single bonds. As a place to begin, put the
least electronegative atom in the center and arrange the atoms as symmetrically as
possible. This step is the trickiest, but there are often clues in the way the formula is
written that will help you.
3. Give every atom an octet, using exactly the remaining valence electrons, and
knowing that each single bond requires two electrons. In the case of H, only 2
electrons (a duet) are needed.
4. Check that (a) you have used the correct number of electrons and (b) you have
followed the octet rule for each atom. If you meet both these criteria, then you have
a valid Lewis structure.
Example: Draw the Lewis structure of acetate, CH3CO2–
1. Count the electrons.
2C
2 x 4 e–
3H
3 x 1 e–
2O
2 x 6 e–
–1 charge
1 e–
total
24 e–
2. Correctly connect the atoms with single bonds.
Notice that the two least electronegative elements (C) are in the
middle of the molecule, and that 3 H atoms are on one, and 2 O
O
C
H
C
H
H
O
atoms are on the other. These clues can be inferred from the way the formula was
written.
Also: Do not use dots for bonds. This will cause you and your instructor to go crosseyed.
3. Use the remaining electrons to give every atom an octet.
O
The H atoms and the left-hand C are satisfied, and adding
electrons around the oxygens gives them octets as well.
C
H
C
H
H
O
So far we have used a total of 24 electrons. There are no more to
use. Yet the carbon on the right is still sharing only six electrons.
We can fix this by moving two electrons from one of the oxygens to be shared between
the C and the O.
O
O
4. Check the structure.
H
H
C
C
H
C
C
O
O
H
H
H
All atoms have an octet (or
duet), and we
–
have used 24 e , the correct total number of electrons. We can also add a formal
negative charge to one of the oxygens, which we will discuss next.
Assigning Formal Charges to Atoms
Up to this point in your chemistry career, we have treated the electrons in a Lewis
structure like the human bodies in Body Worlds (the traveling plastination exhibit). We
do not ask where they came from. We just count them and move on. However, if we
want to be more careful, we can ask ourselves how many electrons a particular atom
brings to the molecule, and whether it currently has that many electrons (shared or
unshared).
Each atom in a molecule has a formal charge. To determine the formal charge, we
need to know two things:
1. How many valence electrons that atom has when it is neutral
2. How many electrons are assigned to that atom in your structure
The formal charge is the difference between the number of valence electrons and the
number of assigned electrons.
FCatomX = # valence electrons – # assigned electrons
Every first grader can find the number of valence electrons for a neutral atom, so we
only need to know how to find the number of assigned electrons. To find the assigned
electrons, consider how many electrons spend their time on the atom on average (this
will usually not be eight). If an electron is in a lone (nonbonding) pair, then it is always
on this atom, and it counts as one electron. But if it is forming a bond, half the time it is
on this atom, and half the time it is on another atom. So bonding electrons count for
half an electron. Since each bond consists of two electrons, we could just say that
overall, each bond counts for one electron.
# assigned e– = # nonbonding e– + ½ (# bonding e–)
Example: Assign the formal charges to both atoms in carbon monoxide.
C
O
C
O
e^{-}
This is how we show the formal charges for CO.
To check your formal charges, add them all up. They should add to the total charge on
the molecule or ion. In our example above, the sum of the formal charges is 0, which is
the same as the charge on carbon monoxide. Bingo!
For many atoms, the formal charge is zero, and we do not write it. But if an atom has
formal charge, you must write it. Always. Also, once you have practiced counting
formal charges a little, it should become a pretty quick habit for you. There is no need
to show calculations if you can do it correctly in your head.
One more thing you might ask: Do these formal charges mean that C really has a
negative charge, and O really has a positive charge in CO? Well, not really, unless you
are talking about an ion, where there really is a whole charge somewhere. In CO, the
atoms probably have partial – and + charges, respectively. But the formal charges
do give us an idea of where and what the delta charges will be. They also help us
choose between resonance structures, which we will consider next.
Resonance Structures
Many times, there is more than one valid Lewis structure for a particular molecule. In
the example of acetate, we could choose to make the double bond to the carbon from
either oxygen. (Note that formal charges have been assigned).
O
O
C
H
C
C
H
O
C
H
O
H
H
H
Both structures satisfy the octet rule and have the correct number of electrons. So
which one is right? Asked another way, which oxygen really has the negative charge?
Before answering this question, we should define resonance structures. A resonance
structure is a valid Lewis structure that is drawn from another structure by moving only
electrons. If any atoms are moved, then the two structures are not resonance
structures. Notice above that the first structure for acetate can be rearranged into the
second structure (as shown) by sharing two electrons from the lower oxygen with
carbon and then moving two electrons from the C=O bond onto the upper oxygen. No
atoms have been moved.
-1/2
O
C
H
C
O
-1/2
H
H
So, which is the correct resonance structure? Both are valid, and
the real structure is something in between, like the ion on the left.
Each oxygen has an average charge of -½, and the CO bonds are
really 1.5 bonds. These structures, though, are a little tedious to
draw, so we usually draw several resonance structures and
assume the reader can combine them mentally.
The simplest way to draw resonance structures is to move around the double bonds
(and any lone pairs as necessary). Once the double bonds have been everywhere they
can go, you should have most of the valid structures.
Example: Draw the possible resonance structures for SCO.
S
C
O
S
C
S
O
C
O
One of these structures has two double bonds, and the others have one single and one
triple bond, arranged in two possible ways. The resonance structures are separated by
double-headed “resonance arrows.” Arrows are important in chemistry, and this is
particular type of arrow is reserved for separating resonance structures. All three
structures follow all the rules for drawing Lewis structures. So which is closest to the
true structure? Do they all contribute equally, or are some more important than others?
To answer this, we should assign the formal charges.
S
C
O
A
S
C
B
O
S
C
C
O
Now, we have a way to compare the structures. Generally, in the best resonance
structures:
1. Each atom has an octet.
2. There is as little formal charge as possible.
3. Any negative formal charges are on the most electronegative atoms.
Using these criteria, we can say that structure B is the best because it has the least
amount of formal charge. After this, structure A is next best because it places the
negative formal charge on the oxygen, which is most electronegative. Finally, we can
say that C is a much less important resonance structure (although it still could contribute
something) because it places a positive charge on O, which is a more electronegative
atom that prefers to be negatively charged.
We can conclude that the “real” structure of SCO is probably something between A and
B, but closest to B.
Sharp-eyed Chem 1A students may notice that in the example of CO, we put a positive
formal charge on O. Isn’t that bad? Yes, it is. But that is the only structure of CO that
has octets, (see below) so we will accept the unusual formal charges rather than lose
an octet.
-2
C
+2
C
O
no octet on O
bad formal charges
O
better than others
C
O
no octet on C
Expanded Valence Shells
In certain situations, it is acceptable to assign atoms with an atomic number of 14 (Si) or
greater a so-called “expanded valence shell” where the central atom has more than 8
electrons. This should be avoided if possible, but there are certain structures where the
only way to eliminate large formal charges is to exceed the octet rule. One case is
when just connecting all the atoms with single bonds gives you more than 8 electrons
(as in AsF5). Another example is XeF6, where the molecule seems to have too many
electrons. (See figures). In such cases, assign the extra electrons to the central atom.
Remember that this is possible only for atoms heavier than aluminum Al.
F
F
F
F
F
As
F
F
As has 10 e–
Xe
F
F
F
F
Xe has 14 e–
Count the valence electrons in each molecule below, and draw Lewis structures for
them.
Part I. Simple structures. None of these molecules have formal charges or common
resonance structures.
PH3
#e– ______
CF4
#e– ______
CCl2F2
#e–
______
H2O
#e– ______
N2
#e– ______
F2
#e– ______
H2N2
#e– ______
HCCH
#e– ______
SF2
#e– ______
Part II. Structures with possible formal charges. Draw one valid Lewis structure,
and show all nonzero formal charges.
H3O+
#e– ______
NH3
NO2–
#e– ______
AsCl5
O3
#e– ______
N2O
SF6
#e– ______
CN–
#e– ______
#e– ______
SO2
#e– ______
XeF4
#e– ______
#e– ______
I 3–
#e– ______
#e– ______
NO3–
#e– ______
Part III. Free-for-all! In this section, draw the Lewis structure and any valid resonance
structures. Label each structure with a letter (A, B, C, etc.), and indicate which is the
best. If they are all equally good, then write equivalent.
SO3
#e– ______
best
structure:
________
CS2
#e– ______
best
structure:
________
N3-
#e– ______
best
structure:
________
CO32–
#e– ______
best
structure:
________
N2O
#e– ______
best
structure:
________
CNO–
#e– ______
best
structure:
________
10.4 POSTLAB QUESTIONS
1. Does adding more electrons to a molecule tend to increase or decrease the number
of bonds? Explain.
2. Do you think having more valid resonance structures would tend to make a molecule
more stable or less stable? Explain.
(Continued)
3. Which of the structures in the table are valid resonance structures of the acetic acid
structure shown? For each invalid structure, give a reason.
H
H
O
C
C
O
H
acetic acid
H
Possible Resonance
Structure
H
H
H
O
C
C
O
H
H
H
O
C
C
H
O
H
Valid?
Reason (if not valid)
H
H
H
O
C
C
O
H
H
H
H
H
O
C
C
H
O
C
C
H
O
O
H
H