Covalent Bond
Def - Bond due to the sharing of electrons between two nonmetals, or a nonmetal and Hydrogen.
Example: F2 – Fluorine
In using Lewis diagrams to model Covalent bonding the atoms are placed next to each other with
the single, unpaired electrons between the atoms. These are the atoms that will be shared to form
the covalent bond.
The bond length is the distance between the nuclei of adjacent atoms. This is the distance at
which the potential energy will be lowest.
Bond Energy – Energy required to separate the atoms (break the bond).
The bond length has a direct affect on the bond energy; the shorter the bond length the higher the
bond energy.
Table of Bond Lengths and Energies:
Bond
H-F
H-O
H-Cl
H-Br
C-O
C-N
H-I
(Slide 1)
Bond Length (pm)
92
96
127
141
143
147
161
Bond Energy (KJ/mol)
569
459
432
366
358
305
299
Table of Bond Lengths and Energies for single vs. multiple bonds.
(Slide 2)
Bond
C-C
C=C
C≡C
Bond Length (pm)
154
134
120
Bond Energy (KJ/mol)
346
612
835
C-N
C=N
C≡N
147
132
116
305
615
887
Most bonds are not purely ionic or covalent but fall somewhere in between. We compare the
electronegativities of the atoms involved to determine:
1) Bond type
2) which atom will have the electrons orbiting around them more frequently
What is important is the electronegativity difference between the elements.
Example 1:
Sodium + Chlorine
Cl = 3.16
Na = 0.93
Difference = 2.23
This is a large difference and so the electrons will spend their time around the chlorine atom
producing an ionic bond.
Example 2:
Chlorine + Bromine
Cl = 3.16
Br = 2.96
Difference = 0.20
While there is a slight difference between these two atoms it is not enough to make a difference
and so the electrons will spend an equal amount of time on both sides of the bond. The electrons
are being shared equally by these two atoms in this covalent bond.
Example 3:
Sulfur + Oxygen
O = 3.44
S = 2.58
Difference = 0.86
This is not a large enough difference for the electrons to be transferred fully, but too large for the
electrons to be shared evenly. In this situation the oxygen, which is stronger, will have the
electrons orbiting about them more often than the Bromine. This results in an uneven
distribution of charges giving bromine a slight positive charge, and the oxygen a slight negative
charge.
δ – symbol for a partial charge
Covalent bonds are classified as either polar or non-polar based on the distribution of charge.
Example 2 is a nonpolar bond, while example 3 is a polar bond.
Polar Bond – A covalent bond in which electrons spend more time on one side giving it a partial
charge.
Nonpolar Bond – A covalent bond in which electrons spend an equal amount of time on both
sided of the bond, giving both sides the same charge.
< 0.3 Nonpolar bond
0.3 ≤ X ≤ 1.7 Polar bond
> 1.7 Ionic Bond
(Slide 3)
Drawing Lewis Structures for Covalent Bonds
Example:
1)
H2O
Determine the valence of each element in the compound. For negative valences the
number indicates the number of spaces to be filled. The atom with the greatest number of
spaces to be filled is usually the central atom.
(Hydrogen and Halogens only have 1 space available, so they will always be at the end of
a series of atoms.)
Hydrogen is -1. Oxygen is -2. Therefore, Oxygen will be the central atoms.
2)
Arrange the atoms in a skeletal structure around the central atom with the atoms
connected by forming electron pairs.
3)
Add any unshared electron pairs around the central atom.
4)
Check to see if the octet rule has been filled.
5)
Determine the electronegativity difference between the elements and indicate the partial
charges for the compound where appropriate.
Oxygen = 3.44 Hydrogen = 2.20 Difference = 1.24 this is a polar bond.
Example 2:
CH4
Carbon has a -4 valence. Hydrogen has a -1 valence. Carbon will be the central atom.
Hydrogen has no unshared pairs of electrons.
Carbon = 2.55
Hydrogen = 2.20
Difference = 0.35
this will be a polar bond.
Example 3:
CH3Cl
Carbon has a -4 valence. Hydrogen has a -1 valence. Chlorine has a -1 valence. Carbon will be
the central atom.
Chlorine = 3.16
Carbon = 2.55
Hydrogen = 2.20
Chlorine/Carbon difference = 0.61 Polar bond with Chlorine being more negative
Carbon/Hydrogen difference = 0.35 Polar bond with Hydrogen more positive
One important exception to the octet rule are the elements in Column 13 which can be stable
with only 6 valence electrons (3 electron pairs)
Example:
BCl3
Boron has a +3 valence. Chlorine has a -1 valence.
Chlorine = 3.16
Boron = 2.04
Difference = 1.12
this will be a polar bond
Resonance
Ozone has the formula O3. Draw a Lewis structure for an Ozone molecule.
It appears that this molecule contains 1 double bond and one single bond. However, tests done
on the molecule indicate that the 2 bonds are identical with a bond length halfway between the
O-O and O=O bond lengths.
This molecule would best be represented by:
The double arrow indicates that both forms are present at the same time. This condition is called
resonance.
Resonance – condition where the molecule cannot be adequately represented by a single model.
Benzene
C6H6
(Kekule 1865)