Covalent Bonding and
Naming
I. Types of Covalent Bonds
• l. Nonpolar covalent bond-a covalent
bond in which the bonding electrons are
shared equally
• 2. Polar covalent bond-a bond is which
the bonded atoms do NOT share the
bonding electrons equally.
Types of Covalent Bonds
• If the difference in electronegativity
between two bonded atoms is from 0.3 to
1.7, a polar bond will exist
• If the difference in EN is less than 0.3 then
the bond is nonpolar covalent.
Practice
a. H and Cl
b. F and Br
c. S and I
d. O and H
The Octet Rule
• chemical compounds tend to form so that
each atom, by gaining, losing, or
sharing electrons, has an octet of
electrons in its highest energy level.
Multiple Bonding
• single bond -two atoms share one pair of
electrons
• double bond -two atoms share two pair of
electrons
• triple bond -two atoms share three pair of
electrons
Drawing Lewis Dot Structures
Add up the TOTAL number of valence
electrons in the substance
Decide what is the central atom. The central
atom is the one that is least represented.
(or the least electronegative)
Arrange the dots so that each atom has an
octet
Practice
1. Br2
2. CCl4
3. H2O
4. NH3
5. O2
6. N2
Exceptions to the Octet Rule
• Some atoms have less than an octet.
Example: Hydrogen only needs 2
electrons surrounding it and boron only
needs 6.
• H2
• BF3
Exceptions to the Octet Rule
Some atoms have more than an octet
(One reason is because of bonding d
orbitals as well as s and p orbitals.)
Example: Sulfur can have up to 12
electrons surrounding it.
SF6
Resonance
• – a concept in which two or more Lewis
structures for the same arrangement of
atoms (resonance structures) are used to
describe the bonding in a molecule or ion.
To show resonance, a double-headed
arrow is placed between a molecule’s
resonance structures. Example: O3
Molecular Shapes
The Valence Shell Electron Pair Repulsion
Theory (VSEPR) is used to predict the
shape of molecules.
The VSEPR theory states that the bonding
and nonbonding pairs of electrons will
arrange themselves so that the repulsive
forces between them are at a minimum.
Molecules will adjust their shapes so that
the nonbonding electrons are as far apart
as possible.
Molecular Shapes
• The number of shared and unshared
pairs of electrons determines the shape of
the molecule.
Linear
•
•
•
•
The bond angle is 180º .
2 atoms bonded to the central atom
0 lone pairs
Ex: CO2
Bent
The bond angle is 90º
2 atoms bonded to
the central atom
2 lone pairs
Write the Lewis
structure for SF2
Trigonal Planar
The bond angle is
120°
3 atoms bonded to
the central atom
0 lone pairs
Write the Lewis
structure for BCl3
Trigonal Pyramidal
The bond angle is
107°
3 atoms bonded
to the central
atom
1 lone pairs
Write the Lewis
structure for NH3
Tetrahedral
The bond angle is
109.5°
4 atoms bonded to the
central atom
0 lone pairs
Write the Lewis
structure for CH4
Trigonal bipyramidal
The bond angle is
120°
5 atoms bonded to
the central atom
0 lone pairs
Write the Lewis
structure for PCl5
Octahedral
The bond angle is
120°
6 atoms bonded to
the central atom
0 lone pairs
Write the Lewis
structure for SF6
Polarity
A. nonpolar covalent bonds – EN
difference is less than 0.3
There is equal sharing of electrons
B. polar covalent bonds – EN is between
0.3 and 1.7
There is unequal sharing of electrons
Practice
• Determine whether each of the following is
a polar covalent bond, nonpolar covalent,
or ionic bond.
•
1. N-O bond
2. Cl-Cl bond
•
3. C-S bond
4. S-O bond
•
5. P-O bond
6. Na-Cl bond
Determining Polarity
In a polar molecule there is an uneven
distribution of charge. One end of the
molecule is more negative or positive than
the other
In a nonpolar molecule there is an even
distribution of charge.
Determining Polarity
Molecules with nonbonding pairs of
electrons on the central atom are polar.
H2O
NH3
Determining Polarity
• The polarity of molecules with no
nonbonding pairs depends on the atoms
bound to the central atom.
• If the surrounding atoms are identical, the
molecule is nonpolar. This is because the
bond dipoles cancel.
• Example: CH4
• When the atoms surrounding the central
atom are different, the molecule is polar.
• Example: BBrI2
Practice
SO2
H2S
CO2
Intermolecular Forces of
Attraction (IMFs)
Van der Waals forces
The attractive forces between molecules are
collectively referred to as van der Waals
forces.
1. There are very strong intermolecular
forces of attraction between polar
molecules and weak intermolecular forces
of attraction between nonpolar molecules.
2. Many physical properties of a
substance, such as freezing point and
melting point, depend on the strength of
the intermolecular forces of attraction.
Dipole-Dipole Forces
•
1. Dipole-dipole forces exist between
polar molecules. The negative dipole of
one molecule attracts the positive dipole of
another molecule.
• Example: HCl
• 2. Dipole-dipole forces are weaker than
chemical bonds.
Hydrogen Bonding
• 1. Hydrogen bonding is a special type of
dipole-dipole force. Since no electrons are
shared or transferred, hydrogen bonding is
not a chemical bond.
• 2. Hydrogen bonding always involves
molecules containing hydrogen that is
chemically bonded to a highly
electronegative atom of small atomic size,
specifically nitrogen, oxygen or fluorine
Example: H2O
London Dispersion Forces
• 1. London dispersion forces are most noticeable
in nonpolar molecules
• 2. London dispersion forces arise from the
motion of electron clouds . From the
probability distributions of orbitals, it is
concluded that the electrons are evenly
distributed around the nucleus. However, at any
one instant, the electron cloud may become
distorted as the electrons shift to an unequal
distribution. It is during this instant that a
molecule develops a temporary dipole.This
temporary dipole introduces a similar response
in neighboring molecules, thus producing a
short-lived attraction between molecules
Strength of Forces of Attraction
from Weakest to Strongest
• 1. Chemical Bonds
– a. Nonpolar Covalent
– b. Polar Covalent
– c. Ionic
• 2. Van der Waals Forces
– a. London Dispersion Forces
– b. Dipole-Dipole Forces
– c. Hydrogen Bonding
Comparing Boiling Points
• 1. When molar masses are similar, substances
with stronger intermolecular forces of attraction
have higher boiling and freezing points.
• 2. In order to compare boiling and freezing
points of a substances with similar molar
masses you must:
• a. Identify the structural formula
• b. Determine the polarity of the molecule.
• c. Determine the dominant type of Van der
Waals forces
• 3. Which of the following molecules has the
higher boiling point: CH4 or NH3? Explain your
answer.
Naming and Writing Covalent
Compounds
• Writing Formulas for Binary Molecular
Compounds those containing 2
nonmetals. Use the prefix naming system
- know theses prefixes
Naming and Writing Covalent
Compounds
•
•
•
•
•
•
•
•
•
•
mono – one
di – two
tri – three
tetra – four
penta - five
hexa – six
hepta – seven
octa - eight
nona - nine
deca – ten
Practice
•
•
•
•
•
•
•
•
nitrogen tetrasulfide
carbon dioxide
oxygen monofluoride
sulfur hexachloride
trioxygen decanitride
tetrafluorine monophosphide
hexafluorine nonasulfide
heptabromine octanitride
Naming Binary Molecular
Compounds
• The less electronegative element is given
first.
• The second element is named by
combining a prefix indicating the number
of atoms contributed by the element to the
root of the name of the second element
and then adding –ide to the end.
Practice
• CCl4
• NF3
• PBr5
SF6
• SO3
• PCl5
• N2O
PF6
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Naming Acids
Binary or hydrohalic acids
Name
“hydro____ic acid”
Ex: HF
____________________
HCl
____________________
HBr
____________________
HI
____________________
H 2S
____________________
Oxyacids – contain a polyatomic ion
If the polyatomic ion ends “ate” the acid will
end in “-ic”
HNO3
____________________
H3PO4
____________________
H2SO4
____________________
H2CO3
____________________
HC2H3O2
____________________
Oxyacids – contain a polyatomic ion
If the polyatomic ion ends “ite” the acid will
end in “-ous”
HNO2
H3PO3
H2SO3
____________________
____________________
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Metallic Bonds
• Each metal donates its valence
electron(s) to form an electron cloud
• This leaves positive particles which are
"cemented" together with the negative
electron cloud, often called a “sea of
electrons.”