Chapter 7 Lewis Formula
In the last chapter we studied ionic compounds. In this chapter we will study
covalent compounds.
In the last chapter the Lewis formula was a nice way to look at the electrons in
the outer shell of an ionic compound and it to help you to understand the
difference between core and valence electrons. In this chapter we use Lewis
formulas to really begin to understand the bonding between atoms in covalent
molecules
7-1 Covalent bonds
Consider the Cl2 molecule
If it was ionic, the reaction to form this compound would look like this:
:
:
:
:
:Cl@ + :Cl@ 6 [:Cl:]- [:Cl]+
:
:
:
:
and we could calculate the energy released to for the ionic gass just as we did in
the last chapter
RXN 1 Ionization E of Cl
+2.09 aJ
RXN 2 Electron Affinity of Cl
-.580 aJ
RXN 2 Coulomb attraction
(assuming Cl+ radius = 100 pm and Cl- radius = 181 pm)
-.82 aJ
NET
+.69 aJ
Positive energy - Not stable - will NOT form
Why? Because ionization E of Cl6 Cl+ too large
So we know Cl2 exists, but it is not ionic. What is it then?
It is a covalent compound, held together with a covalent bond
:::Cl:Cl:::
where a pair of electrons is shared between atoms so both atoms can obtain a
noble gas configuration
Key Concepts:
A covalent bond is formed when two atoms share electrons
In a Lewis Formula a covalent bond is indicated with a line
In a Lewis formula unpaired or lone pairs of electrons are shown with :
Unlike our ionic compounds there is no alternating charge holding the solid
together. Thus the Cl2 solid is much weaker with a MP of -101oC
2
Diatomic molecules do have a distinct distance between atoms
Key Concept:
Bond length - The distance between two atoms held together by a covalent bond
7-2 The Octet rule and Lewis Formulas
In most covalent compounds we will find that atoms try to have 8 electrons. We
call this the octet rule. (There will be exceptions that we will get into later.)
Key Concept:
Each element in a covalent compound forms bonds to that each atom has 8
electrons in its outer shell (Octet Rule)
So let’s start making Lewis formulas for different compounds
Here are some general rules to start with
1. Write the atoms on your piece of paper. Generally start with the different
atom in the middle (Odd man in the middle)
2. Count up all the valence electrons in all atoms and divid eby two to find the
number of bonds or lone pairs
3. Make bonds between the atoms with dashes. (Each bond is a pair of shared
electrons)
4. Add lone pairs of electrons so you use up all electrons, and all atoms have a
filled octet.
Let’s make some Lewis formulas
Let’s apply this to something complicated NF3
Total valence electrons? 5 + 3(7) = 26
Put pairs into bonds
F
:
F:N:F
F
|
or F-N-F
Add remaining electrons to satisfy octet rules
:
:
:F:
:F:
: : :
: | :
:F:N:F: or :F-N-F:
: : :
: : :
So we can see NF3 will have 3 pairs of bonding electrons shared between the N and the
F’s and one lone pair of electrons held by the N.
3
7-3 Hydrogen and Lewis Formulas
Hydrogen is our first exception to the octet rule
Since it only needs 1 additional electron to get to a noble gas, it will only need 1
pair of electrons around it. This is sometimes called the duet rule.
With one pair of electrons in hydrogen it is almost always on the outside of a
compound (terminal atom)
Key Concept:
Hydrogen should only have 1 pair of electrons
Let’s try a NH3
Total valence electrons? 5 + 3(1) = 8
Put pairs into bonds
H
:
H:N:H
Add remaining electrons to satisfy duet and octet rules
H
:
H:N:H
:
So we can see ammonia will have 3 pairs of bonding electrons shared between the N
and the H’s and one lone pair of electrons held by the N.
How about water H2O
H. H. :O:
:
:
H:O:H
:
4
Let’s add one more complication and go to an ion, NH4+
Total valence electrons? 5 + 4(1) = 9
But the + charge says we are missing an electron so
9-1=8
Put pairs into bonds
H
|
H-N-H
|
H
The proper way to finish an ion is the to put it in a quare bracket with a charge
as a superscript
[The above]+
Another complication. So far we have had an odd man in the middle, and that
odd guy was the central atom. What if we don’t have a single central atom?
Let’s try N2H4 (Hydrazine)
In this case try for a symmetric molecule. Also, knowing H are usually on the
outside we can start with :
H
H
N
N
H
H
Adding bonds
H
H
|
|
N----------------N
|
|
H
H
Adding lone pairs
H
H
|
|
:N---------------N:
|
|
H
H
5
This is how you would start more complicated molecules. If we get a chance we
will do organic chemistry, that is molecules based on carbon, this included things
like proteins and DNA where you can have thousand and millions of atoms all
linked together in one Lewis Formula!
7-4 Formal Charge
As we get to more complicated molecules, you will occasionally find that there
can be more than one arrangement of bonds and electron pairs that will satisfy
the octet rules for all the atoms. The way you decide which arrangement is best
is by looking at the formal charge on the atoms
In this process we will use a set of arbitrary rules to assign a charge for every
atom. Since these rules are arbitrary, the charges we assign may not reflect the
actual charge in the molecule. That is why we use the term Formal Charge.
Key Concept/Calculation:
Formal Charge on an atom = # of valence electron in atom - total number of lone
pair electrons - ½(total number of bonding pairs)
-orF.C. = # Valence Electron - # lone pair electrons -1 for each bond
Note that the sum of the formal charge in all atoms in a molecule or ion should
match the net charge on that molecule or ion.
Going back to our NH4+ ion
For the H
Each H has a single bond so
F.C. for a H = 1- ½ (1 bonding pair) = 0
Or F.C = 1-0 pairs -1 bond =0
For the Nitrogen
it has 4 single bonds so
F.C. = 5 - ½ (bonding pairs) = 5 - ½ (8) =+1
or 5-4(1) =+1
Note that before when we had an ion, we would put the ion in a square bracket
and put the charge of the ion as a superscript on that bracket to indicate that the
molecule as a whole had a charge. Once we assign a formal charge to an atom,
we can actually put a charge by the atom(s) we have assigned the charge(s) to.
H
|
H-N+-H
|
H
6
Now that you know how to assign form charge, how do you use that to tell which
structure is best?
Key Concept:
The best structure have:
1. The most atoms with 0 formal charge
2. Most negative formal charge is on most electronegative atom
(Section 7-9)
Example:
Which is the better structure for OF2:
F-O-F or F-F-O
: : :
:F-O-F:
: : :
12 3
: : :
:F-F-O:
: : :
45 6
FC atom 1 = 7-7-1 =0
FC atom 2 = 6-4-2 =0
FC atom 3 = 7-7-1 =0
FC atom 4 = 7-6-1=0
FC atom 5 = 7-4-2=+1
FC atom 6 = 6-6-1= -1
Bad structure, two atoms with FC …1
7-5 Multiple bonds
It is sometimes necessary to use more that one bond between atoms to get all
atoms satisfied. Take the example of C2H4
Total = 4(2) + 1(4) = 12
H:C:C:H
: :
H H
Only accounts for10 of 12 electrons. We need to put one more pair
Of electrons between the two C to get everybody happy
When you have multiple bonding pairs, these are written as dashes
H:C=C:H
: :
H H
You can have triple bonds as well. Let’s try O2 and N2
7
O2
6 electrons / atom for a total of 12 electrons
:O=O:
: :
N2 5 electrons/atom for a total of 10 electrons
:N/N:
Besides helping us to make nice Lewis Formulas, Double and triple bonds have
important physical properties that allow us to confirm their existence in the actual
molecule.
1.) Multiple bonds have shorter bond lengths
2.) Multiple bonds have higher bond energies
See Table 7.3
Note on multiple bonds. In general only use if at lease 1 atom in pair is C, N, or O.
These are usually the only elements that use multiple bonds
Key Concept:
You can use multiple bonds in Lewis Formulas, but generally only when one of
the atoms in the bond is N, C, or O.
7-6 Resonance Hybrids
For many compounds it is possible to write two or more equally good Lewis
formulas without moving the atoms involved in the structure
Take for example the NO3: 24 electrons
:O:
: :
:O:N=O: with a net -1 charge
:
:
Use [ ]-:[ ]-:[ ]- formalism
Notice that I had the double bond to the left and the single bonds up and to the
right. Any reason I can’t switch these?
So double bond could be in any one of three positions
Now another fact. How can you tell double bonds form single bond?
(Higher energy, shorter distance)
If look at NO3- find 3 bonds of equal length, intermediate between single and
double bonds in length, and all bonds have exactly the same energy.
8
The actual answer here is that none of these three structures is itself correct, but
that the true structure is an average or superposition of all three.
Key Concept:
Resonance occurs when more than one valid Lewis structure can be written for
a particular molecule. We usually write double headed arrows between these
structure. This is to indicate not that the true structure is flipping from one form
to the next, but that it is a simultaneous average of all the resonance state.
7-7 Free Radicals and Electron Deficient Compounds
Free Radicals
Free radicals are compounds (or atoms) with an odd number of electrons
It is impossible to come up with a good Lewis Formula for a free radical
compound because you can never satisfy all your Lewis rules
Free Radicals are usually highly reactive because they are either trying to
get rid of their odd electron by passing it off, or they are trying to grab an
electron from some other compound to get a pair.
Examples: NO, ClO2, NO2
Key Concept:
Free radicals are compounds with an odd number of electrons. They are highly
reactive. They do not yield good Lewis Formulas
Electron Deficient compounds
Be, B, and Al are another exception to the octet rule
Several Boron molecules exist where B hasn’t fully filled its octet. One of
these compounds is BF3. We can propose 2 structures for BF3
:
:F:
:F:
: ::
: || :
:F:B:F: And :F:B:F:
: :
:
:
The first doesn’t look like a good structure because the B needs another
pair of electrons. The second looks more reasonable from electron count,
but I just said to avoid double bonds if an atom isn’t N, C, or O, so which
is correct?
9
Experimentally BF3 reacts like it wants another pair of electrons. It reacts
strongly with compounds like NH3 and H2O that have pair of electrons to
donate. Thus, experimentally the evidence supports the 1st of the two
forms. Other highly reactive electron deficient boron compounds also
exist, so always watch out for boron, Be and Al
Key Concept:
In compounds Be, B, or Al may have less than an octet
7-8 Expanded Valence Shells
Here is yet another exception
Elements in the third row can have more than an octet of electrons. Can you
figure out why? (3s,3p & 3d orbitals)
Thus a molecule like SF6 exist. Here we have a total of 6 + 6(7) or 48 electrons.
The only way we can make this molecule work is if we have 6 shared pairs of
electrons with the S, not the 4 we would expect using the octet rule. We assume
that sulfur is used its 3d orbital to hold these extra electron pairs.
Another nasty example is I3This should have 7(3) + 1 or 22 electrons
The simplest model would be
:::
:I:I:I: which would only require 20 electrons, thus we have to put 1 more
:::
Electrons pair some place.
As a general rule, when it is necessary to exceed the octet for 3rd or higher row
elements, assume the extra electrons are placed on the central atoms (part 2 of
rule 7)
::: :
:I:I:I: which would only require 20 electrons, thus we have to put 1 more
:::
Electrons pair some place.
Let’s try 1 more. While I have said in the past that the noble gases are nonreactive, Radon (6th period) can react with Cl to form RnCl2
Electrons 8 +2(7) = 22
: : :
Simplest model :Cl:Rn:Cl: but this has only 20 electrons, so adding a pair
: : :
To the central Rn atom
10
Another example, SO42One Lewis structure we can come up with for this is:
:
:O:
: : :
:O:S:O: Here everybody is happy, all octets filled
: : :
:O:
:
If we assume however that since sulfur is in the third row, it might be happy with
more than an octet, we can get several other structures, one of which is:
:
:O:
: :
:O=S=O: where our sulfur has 12 electrons around it.
: :
:O:
Which structure is better? Let’s assign formal charge to each atom
: In the first structure each oxygen has 6 lone pairs electrons and 1 shared
Pair of electrons for a total of 7 valence electrons In an uncharged state it
would have 6 valence electrons to it has a net FORMAL charge of -1
For the Sulfur we have 8 shared pairs for 4 valence electrons when it was 6 in
the elemental state to this is +2 so our formal charges are
+2, -1, -1, -1, -1 (total of -2 for consistency check)
In the second model we have 1 oxygen at -1, 2 oxygens at 6-6 = 0
and the sulfur at 6 -6 =0 for a net of
0, -1, -1, 0, 0. So this model has more charges close to zero and is therefor
considered the better model. Can actually make several resonance forms of this
as well.
7-9 Electronegativity
Last chapter we talked about ionic bonds, where electrons are not shared. The
cation is missing and electron and has a positive charge, the anion has extra
electrons and has a negative charge, and it is the charge/charge interaction that
holds the compound together
This chapter we have talked about covalent bonds, where atoms share electrons
equally so there is no charge on either atom, and it is the share electrons that
hold things together
In reality bonds can be in between. There is sometimes covalent character in
ionic bonds, and there is often ionic character in covalent bonds.
11
When we assigned our formal charges we assumed that the 2 electrons in a
bonding pair were shared equally between the atoms. Ain’t necessarily so. In
reality, there are some atoms that are electron hogs and others that will give
them away to somebody else.
Key Concept:
Electronegativity is a measure of the tendency of an atom to pull electrons to
itself in a covalent bond
Eletronegativity is something that we can’t measure directly. Several people
have tired to come up with various measures of this tendency, and the most
successful was Linus Pauling.
Shown in Figure 7.13 is Pauling’s electronegativity scale
This scale runs form 4 for the most electronegative to 0 for the least
eletronegative
The noble gases are not given values in this table because they usually don’t
form covalent bonds. (If they don’t form a bond, then we can’t measure their
tendency to hold electrons in this bond.)
If you study this table you will see that electronegativity is another periodic
property. If get bigger as you go to the right and up, it get smaller as you go left
and down.
Because this is an arbitrary scale we will only worry about the difference between
two numbers on the scale, not about the absolute size of any single number on
the scale
7-10 Polar Bonds
The difference between the electronegativities of two atoms in a bond
determines how equally the atoms are shared.
If the difference is between 0 and < .4 we call this a ‘pure covalent bond,
and assume the electrons are equally shared
If the difference is between .4 and 2.0. then we call this a polar covalent
bond
A polar covalent bond is a covalent bond that has some ionic
character to it, to one atom is slightly +, and the other atom sis
slightly -, and this +/- interaction gives the bond added strength
If the difference is >2, then we have an ionic bond
12
Key concept/calculation:
To determine the character of a bond
1.) Look up the electronegativity of the atoms in a bond
2.) Subtract the larger number from the smaller number
If the result is <.4 you have a pure covalent bond
If the result is >2.1 you have a pure ionic bond
If the result is in between you have a polar covalent bond that has
some ionic properties
Sample calculations:
What kind of bond is a
CH bond (2.55-2.1=.45 slightly polar)
OH bond (3.44-2.1 =1.34 distinctly polar)
NaH bond (2.1-.93= distinctly polar)
NaCl bond (3.16-.93 = 2.23 ionic)
Clicker question: evaluate the polarity of some bonds
The partial ionic character of a bond is something that we can actually measure
because the more electronegative atom a little more negative charge than it
should, and the least electronegative atom has a slightly positive charge
because it has less electrons around it than it should
We indicate these partial charges with the symbols ä+ and ä-. The ä helps us to
remember that we don’t really have a full + or - charge, it is just a slight charge
Now that you know what eletronegativity is, let’s take a second look at one of our
formal charge problems
So let’s look at the 3 SCN- structures:
:::S-C=N: :::S=C=N:: : :S=C-N:::
1
2
3
e in element
S
6
C
4
N
5
e- in structure 1
F.C.
7
-1
4
0
5
0
2
6
0
4
0
6
-1
3
5
+1
4
0
7
-2
-
F.C.
F.C.
13
#3 is the worst because it has 2 atoms with non-zero formal charge
Since the electronegativities are: S 2.5, C 2.5, and N 3.0, structure 2 is better
than 1 because it has the negative formal charge on the most electronegative N
7-11 Dipole moments
When we have a polar bond we have two atoms, one with a partial + charge and
one with a partial negative charge. The existence of a ä+ and ä- in a molecule
can lead to what is called a dipole moment.
Dipole means two poles, a + pole and a negative pole
We represent a dipole with an arrow pointed from one pole to the other.
Here I will break from the text
The symbol I will use for a dipole is an arrow, but I put a +
sign on one end of the arrow to indicate where the + pole is,
and point the arrow toward the negative end.
The book is just the opposite. It doesn’t use a + sign at the
base (feather) end of the arrow, and in this book the
arrow points from the - end to the positive!
Whether the arrow points from positive to negative or negative to positive is not
as important at the fact that the arrow indicates a vector quantity.
Vector quantities have a magnitude and a direction. And if you add vector
quantities together you have to take both the magnitude and the direction into
account.
In fact, vector quantities can cancel each other out.
14
Dipole moments in Molecules
We now know how to determine if a bond is polar (had a dipole moment), but
now let’s go to the molecular level, how can we tell if the molecule as a whole is
polar and why is this important?
First the importance. I’m sure you have heard the phrase ‘like dissolves like’
What this means is that polar molecules will dissolve in polar solvents and
nonpolar molecules will dissolve in nonpolar solvents.
Polar molecules are molecules with a dipole, non polar molecules ones with very
little or no dipole moment. So the dipole of a molecules is the key to
understanding what solvents it will and won’t dissolve in. (Much more on this in
chapter 16)
The dipole of a molecule is found my summing up, in vector manner, all the
dipoles of all the individual bonds within the molecule. This may sound like lot’s
of nasty math, but don’t worry; we will do this graphically not mathematically.
Let’s start with CO2. Is CO2 polar or nonpolar? That is, does it have a net
dipole(polar) or no net dipole (non-polar)
In the next chapter you will learn how to predict that CO2 has a linear structure
::O=C=O::
O is more electronegative than C so we have two bond dipoles pointing in
opposite directions away from the C (Note the book is the opposite pointing in)
These two dipole cancel out so there is no net dipole in the molecule, and CO2 is
nonpolar.
Now look at water. Again in the next chapter you will find out how to determine
that H2O had a V shape: H-O::
|
H
Now we have two vector pointing from the H to the O, but they don’t cancel each
other out, so the molecule has a dipole, and is polar (Well DUH)