Chapter 8 Covalent Bonds
Covalent Bonds are bonds that share their electrons
- different from ionic bonds because electrons are not donated or received
Covalent bonds form molecules
- molecular compounds- are compounds composed of molecules
ex. CO2 , H20, NO2 all molecular compounds
- diatomic molecules - are elements that must exist in pairs if found in their
elemental form in nature. (H, Br, O, N, Cl, I, F)
- ionic compounds cannot form molecules. There is no such thing as a NaCl
or MgCl2 molecule!
- Molecular formula- chemical formula for a molecule.
- shows how many atoms of an element the molecule contains.
- C2H8 contains 2 carbon atoms and 8 hydrogen atoms.
- molecular formulas do not have to be the lowest whole number ratios.
You must right exactly how many atoms are in the molecule.
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Molecular formulas do not indicate the structure of the atom.
How do Covalent Bonds form?
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Covalent bonds are usually between nonmetals or anions. The exception is
hydrogen, because it only has 1 electron it covalently bond.
Elements still want a noble gas configuration, so they must find electrons to
share to fill their valences.
F
F
O
H
H
- Hydrogen will always form single bonds. Other elements have the ability to
form multiple bonds.
Multiple Bonds include
- Single bonds- sharing 1 electron pair
- Double bonds – sharing 2 electron pairs
- Triple bonds- sharing 3 electron pairs
We can determine the number of bonds from the following equation:
1. Write out all of the atoms in a compound.
2. Determine how many valence electrons each has and how many would
be in a full octet for each element. ( remember Hydrogen only wants 2
valence electrons)
3. Subtract what the valence electrons you have from the value of
electrons if they were all stable.
4. Divide by 2.
5. That will give you the number of bonds formed.
Ex.
NH3 N has 5 wants 8
H has 1 wants 2
H has 1 wants 2
H has 1 wants 2
Total has 8 wants 14.
14-8=6/2=3bonds
Nitrogen forms a single bond with each of the hydrogen.
Naming molecular formulas
When naming molecular formulas,. you must use the prefixes that represent the
number of atoms and element has in the molecule.
1= mono
2= di
3= tri
4= tetra
5= penta
6= hexa
7= hepta
8= octa
9= nona
10= deca
CO2 = carbon dioxide
Common molecules that do not follow this pattern
Diatomics- they are called by their elemental name( O2= oxygen)
H2O = water
NH3 = ammonia
CH4 = methane
Coordinate covalent bonds
-When one atom donates both electrons in the covalent bond.
Ex. Carbon monoxide – CO
C
O
Carbon will share 2 electrons with oxygen.
This gives carbon 6 and oxygen 8 valence electrons.
C
O
Oxygen will donate 2 electrons to carbon so both
octets are full.
Polyatomic ions are formed this way. They must have electrons from cations to fill the
valences of covalent bonds.
Examples page 225.
Bond dissociation energy.
The amount of energy required to break a covalent bond.
The higher the bond dissociation energy, the stronger the covalent bond.
Ex. C-C single bond has a dissociation energy of 346 kJ/mol
C=C double bond has a dissociation energy of 657 kJ/mol
Resonance
- when there is more than one valid electron dot diagram
- ex. Ozone O3
can be written 2 ways:
O=O-O or O-O=O
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Resonance structures are hybrids of the same molecule. Shown by a double
headed arrow
O=O-O  O-O=O
Exceptions to the octet rule:
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odd number of valence electrons (NO2)
has 17 valence electrons, impossible to satisfy the octet rule, but it does
form a stable molecule.
Phosphorus and Sulfur
Phosphorus can share either 3 or 5 electrons
Sulfur can share 2 or 6 electrons
PCl5 or PCl3
SF2 or SF6
Bonding Theories
When atoms combine, they form molecular orbitals instead of atomic orbitals.
-2 electrons are still needed to fill a molecular orbitals. If this is accomplished by
a covalent bond, then the orbital is called a bonding orbital.
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two types of bonds
o Sigma bonds- when two orbitals overlap on the axis and form a bond.
 Can occur with s or p orbitals.
o Pi bonds- bond that occurs either below or above the bonding axis.
 weaker bonds than sigma bonds.
VSEPR Theory
Stands for Valence shell electron pair repulsion.
Helps to explain the shapes of molecules.
- all electrons want to be as far away from one another as possible.
- VSEPR theory cause molecular shapes to adjust to accommodate the repulsion
- CH4- forms a tetrahedral structure. All bonds form at a 109.5 degree angle
- NH3 the non bonding pair pushes the other electrons away to form a
pyramidal structure with bond angles of 107 degrees.
- CO2 has a bond angle of 180 or linear.
Orbital hybridization
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Provides information about both molecular bonding and shape.
Double and triple bonds create more pi bonds
Polar Bonds
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nonpolar covalent bonds
o diatomics are nonpolar covalent bonds
o any molecule that the electrons are shared equally is nonpolar.
polar bonds
o electrons are shared unequally.
o Depends on electronegativity.
o The more electronegative atom has a slight negative charge, forcing
the less electronegative atom to have a slight positive charge.
o H- Cl , Hydrogen has a lower electronegativity than chlorine, so the
electrons move closer to chlorine giving it a slightly negative charge.
Electronegativity and bond types.
o 0.0 – 0.4 nonpolar covalent
o 0.4 – 1.0 Moderately polar covalent
o 1.0 – 2.0 Very polar covalent
o > or = 2.0 Ionic
Problems 8.3 on pg 239
Polar molecules are called dipoles, meaning that the polar molecule has 2 poles.
Intermolecular forces.
- Intermolecular forces are weaker than ionic or covalent bonds.
o Van der Waals Forces- weakest of the forces
 Dipole interaction- polar molecule attraction
 Similar to but weaker than ionic bonds
 dispersion forces- the weakest of all molecular interactions.
 Caused by the motion of electrons, the more electrons
the stronger the force.
o Hydrogen bonds- hydrogen bonds to a very electronegative atom but
also bond to an unshared pair of another electronegative atom.
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