A.P. Chapter Eight Outline
Covalent Bonding



In nature, atoms are rarely found uncombined, generally only the noble gases.
In an ionic compound, the transfer of one or more valence electron from an atom
of a metal to an atom of a nonmetal produces ions who’s opposite charges hold
the ions together in a crystal lattice.
In a molecular covalent bond, atoms are held together by bonds consisting of one
or more pairs of electrons shared between the bonded atoms.
I.
Covalent Bonds: shared electron pairs connecting atoms in molecular
(covalent) compounds
1. A single covalent bond is two electrons (one pair) being shared by two
atoms
2. A double covalent bond is four electrons (two pairs) being shared by tow
atoms
3. A triple covalent bond is six electrons (three pairs) being shared by two
atoms
II.
Lewis Structures
 Lewis structures are diagrams that represent how atoms in a
molecule are bonded together. A Lewis structure shows all
valence electrons in a molecule as dots or lines that represent
covalent bonds. A line connecting two atoms represents a shared
pair of electrons. A dot represents an electron, so a pair of
electrons is two dots.
 The number of electrons that the atom must share to achieve a
noble gas configuration determines the number of covalent bonds
an atom can form.
 Octet Rule: To form bonds, main group elements gain, lose, or
share electrons to achieve a stable electron configuration
characterized by eight valence electrons.
 The number of valence electrons an atom has is equal to its main
group number, which is also the number of bonds that atom can
form.
 The number of electrons that an atom of a main group element
must share to achieve an octet equals eight minus it’s a group
number.
A. Guidelines for Writing Lewis Structures
1)
Count the total number of valence electrons is the
molecule. For an atom this is the main group
number. For a negative ion, the total number of
electrons is the group number plus the charge
amount. (For example, O2- has six valence electrons
2)
3)
4)
5)
because it is in group 6A, plus two because the ion
charge is –2). For a positive ion, the total electron
count is the atoms group number minus the charge
amount (for example for NH4+ the total electron
count is 5 for nitrogen since it is in group 5A, plus 1
for each of the four hydrogen’s since hydrogen is in
group 1A, minus one because the charge is +1,
giving a final count of 8).
Use atomic Symbols to draw a skeleton structure
by joining the atom with shared pairs of electrons (a
single line). A skeleton structure indicates the
attachment of terminal atoms to a central atom. The
central atom is usually the one written first in the
formula, and is the one that can form the most
bonds.
Place lone pairs (unbounded pairs of electrons)
around each atom to satisfy the octet rule (except
hydrogen), starting with the terminal atom
Place any leftover electrons on the central atom,
even if it will give the central atom more than an
octet.
If the number of electrons around the central atom
is less than eight, change single bonds around the
central atom(s) to multiple bonds.
III.
Single Covalent Bonds in Hydrocarbons
A. Hydrocarbons contain only carbon and hydrogen atoms. Their general formula is
CnH2n+2. They are referred to as alkanes.
- Saturated hydrocarbons: each carbon atom is bonded to a maximum number of
carbon atoms.
- Cycloalkanes: saturated hydrocarbons consisting of carbon atoms joined in rings.
- Functional groups: a distinctive group of atoms in an organic compound that
imparts characteristic chemical properties to the molecule. For example, the –OH
group is characteristic of alcohols.
B. Alkenes have one or more carbon-carbon double bond. Their formula is CnH2n.
They are unsaturated.
C. Alkynes have one or more triple bonds. Their formula is CnH2n-2.
D. Double bonds and isomerism: Double bonds hold the molecule stiff and unable to
rotate along the double bond. When two atoms or groups of atoms are attached to
carbon atoms on the same side of the C=C double bond, they are said to be cis to
each other and the isomer is the cis isomer. When two atoms or groups of atoms
are on opposite sides, they are trans to each other and the compound is the trans
isomer.
E. Saturated Fats are triglycerides: 3 fatty acids + glycerol  triglyceride + H2O
IV.
Bond Properties: Bond Length and Bond Energy
A. Bond Length: the distance between nuclei of two bonded
atoms
1. Bonds form between atoms to reduce potential energy. Energy is at a
minimum where there is a balance of electrostatic attraction and repulsion
(between electrons and protons)
2. The size of the atoms determines the size of the bond. A bond between
two small atoms will be much smaller than the bond between two large
atoms.
3. Single bonds are shorter than double bonds, and double bonds are shorter
than triple bonds.
B. Bond Enthalpies = Bond Energy: the enthalpy change that
results when the bond between two atoms is the gaseous
phase is broken and the atoms are separated (at constant
pressure).
1. Bond breaking requires energy – always! Remember
that breaking bonds is an endothermic process, whereas
bond forming is exothermic (energy released).
2. It takes more energy to break a double bond than a
single bond, and more to break a triple bond than a
double bond.
3. Remember bond enthalpies are for breaking bonds in
the gaseous phase. If a reactant or product is not in the
gaseous state, you need to vaporize it before you can
break the bonds.
V.
Bond Properties: Bond Polarity and Electronegativity
 A covalent bond is a pair of electrons being shared by two atoms.
 Unless the two atoms are identical, the electrons are not shared
equally – the more electronegative atom will have a greater
“pull” and will attract the shared electrons closer. This results is
3 types of bonds:
1) Nonpolar covalent bonds: The electron pair is
being shared equally between the two atoms.
2) Polar covalent bonds: the electrons are NOT shared
equally
3) Ionic bonds: the electrons aren’t shared at all but
are transferred from one atom to another.
 How to determine type of bond: Use tables of electronegativity. If the difference in
electronegativity between the two atoms is less than 0.3, than the bond is nonpolar.
 If the difference in electronegativities between the two atoms is between 0.3 and 1.7,
than the bond is polar covalent.
 If the difference is more than 1.8 than the bond is ionic.
*** Remember the periodic trend for electronegativity is across a period and up a group;
overall towards fluorine.

The more electronegative atom will have a partial negative charge, signified with
the symbol -, and the less electronegative atoms will have a partial positive
charge signified by +.
VI.
Formal Charge
 There are times when more than one Lewis structure can be drawn. To determine
which is correct, determine the formal charge of the individual atoms of the
molecule.
 The formal charge of a bonded atom is the charge it would have if its bonded
atoms were shared equally. To calculate formal charge:
1) All of the lone pair electrons are assigned to the
atom on which they are found.
2) Half of the bonding electrons are assigned to each
atom in the bond.
3) The sum of the formal charges must equal the actual
charge; zero for neutral molecules and the ionic
charge for an ion.
Formal charge for an atom = (number of valence electrons in an atom –
[(number of lone pairs of electrons) + (1/2 number of bonding electrons)]
4) Smaller formal charges are more stable (therefore
more favorable) than larger ones.
5) Negative formal charges should reside on the more
electronegative atoms.
6) Like charges should not be on adjacent atoms.
VII.

Lewis Structures and Resonance
Many molecules, such as ozone (O3) cannot be drawn as a single Lewis
structure.
There is a third possible structure, but it is not stable and has never been shown to exist.
Each of the oxygen’s is single-bonded to the other two oxygen’s, forming a triangle.
Notice how I used an equilibrium-type arrow between the two structures. Some people
misinterpret this to mean the actual molecule shifts back and forth between the two
structures.
IT DOES NOT!!!!!!!


VIII.
Each of the Lewis structures, called resonance structures, is thought of
contributing to the true structure that cannot be written. The actual structure is
a resonance hybrid.
For resonance structures, remember the following:
1) Lewis structures differ only in the assignment of
electron pair positions; the positions of the atoms
don’t change.
2) Resonance hybrids represent a hybrid structure
representing a composite of the Lewis structures,
not different structures vibrating back and forth.
Exceptions to the Octet Rule
1. Molecules or ions with central atoms having fewer than eight electrons:
Small atoms such as Boron do not form an octet. Boron has 3 valence
electrons and can only form 3 bonds. Giving it a full octet makes it very
reactive.
2. Molecules or ions with an odd number of valence electrons:
Atoms or molecules that have an unpaired electron are known as free
radicals. Free radical molecules are very reactive.
3. Molecules or ions with central atoms having more than an octet of
electrons:
Central atoms can have “expanded octets”, or more than 8 electrons.
Compounds with expanded octets occur only with elements in the 3rd
period and beyond.
IX.


Aromatic Compounds
Aromatic compounds are organic compounds characterized by the presence of
one or more benzene or benzene-like rings.
Constitutional isomers of aromatic compounds
- Atoms or functional groups can be substituted for hydrogen’s on the
benzene ring.
-
X.
Ortho – a prefix designating the substituents are adjacent to each other
Meta – a prefix designating the substituents are separated by one
carbon atom on the ring
Para – a prefix indicating two carbon atoms separate the substituents.
Molecular Orbital Theory
The Molecular Orbital Theory does a good job of predicting electronic spectra and
paramagnetism, when VSEPR and the V-B Theories don't. The MO theory does not need
resonance structures to describe molecules, as well as being able to predict bond length
and energy. The major draw back is that we are limited to talking about diatomic
molecules (molecules that have only two atoms bonded together), or the theory gets very
complex.
The MO theory treats molecular bonds as a sharing of electrons between nuclei. Unlike
the V-B theory, which treats the electrons as localized balloons of electron density, the
MO theory says that the electrons are delocalized. That means that they are spread out
over the entire molecule.
Now, when two atoms come together, their two atomic orbitals react to form two possible
molecular orbitals. One of the molecular orbitals is lower in energy. It is called the
bonding orbital and stabilizes the molecule. The other orbital is called an anti-bonding
orbital. It is higher in energy than the original atomic orbital and destabilizes the
molecule.
Below is a picture of the molecular orbitals of two hydrogen atoms come together to form
a hydrogen molecule:
The MO Theory has five basic rules:
1. The number of molecular orbitals = the number of atomic orbitals combined
2. Of the two MO's, one is a bonding orbital (lower energy) and one is an antibonding orbital (higher energy)
3. Electrons enter the lowest orbital available
4. The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle)
5. Electrons spread out before pairing up (Hund's Rule)
Below is a molecular orbital energy diagram for the hydrogen molecule. Notice that the
two AO's or atomic orbitals combine to form 2 MO's - the bonding and the anti-bonding
molecular orbitals. Also, notice that the five rules have been followed, the electrons
having been placed in the lowest energy orbital (rule 3) and have paired up (rule 4) and
there are only two electrons in the orbitals (rule 5).
.
If you notice at the very bottom of the above picture, "bond order" is mentioned. If a
molecule is to be stable, it must have a bond order greater that 0. Bond order is calculated
as: 1/2 (# of electrons in bonding orbitals - # of electrons in anti-bonding orbitals). If the
bond order is 0, the molecule is unstable and won't form. If the bond order is 1 a single
bond is formed. If the BO (bond order) is 2 or 3 a double or triple bond will be formed
respectively.
When the 2nd period atoms are bonded to one another, you have both 2s and three 2p
orbitals to contend with. When this happens, you have twice as many MO's! Below is a
diagram of the new molecular orbitals:
.
Finally, we can put the Molecular Orbital Theory to use!! Would you predict that
dilithium or diberylium is more likely to form, based on the diagram below?
.
The answer is dilithium because it has a bond order of 1, which is stable, and diberylium
has a BO of 0, which is unstable and therefore will not form.
Look at the following MO diagrams for some of the period two elements. Can you tell
which molecules are paramagnetic? Which molecules have the highest bond energy,
which has the lowest? Rank single, double, and triple bonds in order of bond energy and
bond length. (Hint a BO of 1 is a single bond, 2 a double...)
.