Unit 7
Covalent Bonding
Bonding
• A metal & a nonmetal transfer electrons
– An ionic bond
• Two metals mix
– An alloy (Metallic bond)
• What do two nonmetals do?
– Neither one will give away an electron
– So they share their valence electrons
– This is a covalent bond
Covalent Bonding
• Nonmetals hold on to their valence
electrons
• They can’t give away electrons to bond
• Still want to be stable!
– Need noble gas configuration (octet rule)
• Get it by sharing valence electrons with
each other.
• By sharing, both atoms get to count the
electrons toward noble gas
configuration.
Covalent Bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons…
 Both end with full orbitals

F
F
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons…
 Both end with full orbitals

F F
8 Valence
electrons
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

8 Valence
electrons
F F
Ways to Illustrate Covalent Bonds
• Molecular formula: shows the number
of atoms of each element in a molecule.
– Ex. PF3
• Lewis Structures: uses dots to represent
bonding between molecular compounds
• Structural Formulas: shows the
arrangement of atoms and bonds
– Shared electron dots are replaced with
a dash
• Models: ball and stick (3-D versions)
Single Covalent Bond
• Occurs between nonmetals or a
nonmetal & hydrogen
• Sharing of two valence electrons (1 pair)
• Different from an ionic bond – electrons
are SHARED not transferred
An example with dots…
•It’s like a jigsaw puzzle
•You will be given the formula
•You put the pieces together to
make everyone stable or happy 
– Most atoms need an octet
– H & He need a duo
– Carbon is often the center
Water
H
O
Each hydrogen has 1 valence
electron
and wants 1 more
The oxygen has 6 valence
electrons
and wants 2 more
They share to make each other
“happy”
Water
• Put the pieces together
• The first hydrogen is happy
• The oxygen still wants one more
HO
Water
• The second hydrogen attaches
• Every atom has full energy levels
HO
H
Structural formula…
•Replace shared dots with a dash
H O
H
Practice – Dots & Structures
•
•CH3I
•H2S
•CH2Cl2
•NH3
•C2H6
•SCl2
•AsF3
•SiH4
•CHF3
Multiple Bonds
•Sometimes atoms share more
than one pair of valence electrons.
•A double bond is when atoms
share two pair (4) of electrons.
•A triple bond is when atoms
share three pair (6) of electrons.
Carbon dioxide
C
O
• CO2 - Carbon is central
atom
• Carbon has 4 valence
electrons
• Wants 4 more
• Oxygen has 6 valence
electrons
• Wants 2 more
Carbon dioxide
• Attaching 1 oxygen leaves the oxygen
1 short and the carbon 3 short
CO
Carbon dioxide
 Attaching
the second oxygen
leaves both oxygen 1 short and the
carbon 2 short
OC O
Carbon dioxide

The only solution is to share more
O CO
Carbon dioxide

The only solution is to share more
O CO
Carbon dioxide

The only solution is to share more
O CO
Carbon dioxide

The only solution is to share more
O C O
Carbon dioxide

The only solution is to share more
O C O
Carbon dioxide

The only solution is to share more
O C O
Carbon dioxide
The only solution is to share more
 CO2 requires two double bonds
 Each atom gets to count all the atoms in
the bond

O C O
Carbon dioxide
The only solution is to share more
 CO2 requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
Carbon dioxide
The only solution is to share more
 CO2 requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
Carbon dioxide
The only solution is to share more
 CO2 requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
Carbon Dioxide
•Replace the shared pairs with
dashes
O C O
Practice
• O2
• CS2
• CH2O
• N2F2
• NO2
• HCN (triple)
• C2H2 (triple)
Exceptions to the Octet
Rule/Patterns of Bonding
1. Some elements with odd number
of valence electrons
–BF3
–PCl5
2. Coordinate covalent bonding
Coordinate Covalent Bond
• When one atom donates both electrons
in a covalent bond
• Carbon monoxide (CO)
CO
Coordinate Covalent Bond
 When
one atom donates both electrons
in a covalent bond.
 Carbon monoxide (CO)
C O
Coordinate Covalent Bond
 When
one atom donates both electrons
in a covalent bond.
 Carbon monoxide (CO)
C O
C
O
Summary of Covalent Bonding
• Covalent bonds occur by SHARING electrons
• Occurs between NONMETALS
• End product is called a MOLECULE
1. Molecular compound - formed with
different elements
2. Diatomic molecules - 2 of the same atom
– There are 7 elements that always form
diatomic molecules
– H2 , N2 , O2 , F2 , Cl2 , Br2 , and I2
Diatomic Molecules
Naming Molecular Compounds
• Easier than ionic compounds
• No balancing charges
•1 mono•2 di•3 tri•4 tetra•5 penta-
•6 hexa –
•7 hepta –
•8 octa –
•9 nona –
•10 deca –
Naming Molecular Compounds
• 1st element – add the prefix that matches
the subscript
– Exception – do not add “mono-” if there is
1 atom
– No aa, oo, or ao double vowels
• 2nd element – add the prefix that
matches the subscript
– Still ends in -ide
Naming
• CO2
– 1st element = Carbon; subscript 1
•Remember exception
•“Carbon”
– 2nd element = oxygen; subscript 2
•Prefix for 2 is di•“Dioxide”
– Full name = carbon dioxide
Practice Naming
• S2Cl2
– Disulfur dichloride
• CS2
– Carbon disulfide
• SO3
– Sulfur trioxide
• P4O10
– Tetraphosphorus decoxide
Name  Formula
• Just look at the prefixes!
• Carbon tetrachloride
– 1 Carbon, 4 Chlorine atoms
– CCl4
• Iodine heptaflouride
• Dinitrogen monoxide
• Sulfur dioxide
Common Names
• H2O – dihydrogen monoxide
– Water 
• NH3 – Nitrogen trihydride
– Ammonia
• CH4 – carbon tetrahydride
• Methane
• HCl – Hydrogen monochloride
– Hydrochloric acid
Names to know!
•NH3 - Ammonia
•H2O - Water
•CO – Carbon monoxide
•CO2 – Carbon dioxide
•SO2 – Sulfur dioxide
•CFl4 – Carbon Tetraflouride
Molecular Shapes
• Lewis diagrams & structural formulas
are 2-dimensional
• Real molecules are 3-D
• If there are 2 atoms, the molecule has a
LINEAR shape (no other options!)
– Carbon monoxide (CO)
• If it has more than 2, how do we figure
out the shape?
VSEPR Theory
• Valence Shell Electron Pair Repulsion
Theory
• Used to predict shape of a molecule
• Negative electrons repel each other and
pairs want to be as far apart as possible
Linear
• Linear: 2 atoms around central atom, no
unshared pairs on central atom
• With three atoms the farthest the two
outer molecules can get apart is 180º.
• Will require 2 double bonds or one triple
bond
180º
O C O
Trigonal Planar
• Trigonal planar: 3 atoms around central
atom, no unshared pair on central atom.
• Angle = 120°
Tetrahedral
H
H C H
H
• 4 molecules around
a central atom
• All single bonds
• Must think in 3-D!
Tetrahedral
• Tetrahedral: 4 atoms
around central atom,
no unshared pair on
central atom
• A pyramid with a
triangular base.
H
H
109.5º
C
H
H
So far…
SHAPE
# SHARED
PAIRS FROM
THE CENTRAL
ATOM
# UNSHARED
PAIRS ON THE
CENTRAL
ATOM
LINEAR
2
0
TRIGONAL
PLANAR
3
0
TETRAHEDRAL
4
0
Molecular Shapes
•But what if there are unshared
pairs on the central atom?
•They still repel each other…
Bent
• Bent: 2 atoms around central atom, 1 or 2
unshared pair(s) on central atom.
H O
H
O
H
H
<109.5º
Bent
• Ball and stick model does not show
unshared electron pairs
Pyramidal
• Pyramidal: 3 atoms around central atom, 1
unshared pair on central atom
H N H H
H
N
H
H
<109.5º
Pyrimidial
• Ball and stick model does not show
unshared electron pairs
Trigonal Bipyramidal
• Trigonal bipyramidal: 5 atoms around
central atom, no unshared pair on central
atom
• Angles = 90° and 120°
Adding to the chart…
SHAPE
# SHARED
PAIRS FROM
THE CENTRAL
ATOM
# UNSHARED
PAIRS ON THE
CENTRAL
ATOM
BENT
2
1 or 2
PYRAMIDIAL
3
1
TRIGONAL
BIPYRAMIDAL
5
0
So how do I determine the
shape of a given molecule?
1. Draw the Lewis diagram
2. Count the shared and unshared
pairs
3. Use the VSEPR Theory to
determine the shape
Which type of bond is it?
• Look at which elements are involved
– Metal & nonmetal = ionic bond
– 2 nonmetals = covalent bond
• Electronegativity – measure of a
tendency of an atom to attract a pair of
electrons
– Influenced by amount of positive
charge in the nucleus & electron
shielding
Differences in Electronegativity.
• Big difference between values (greater
than 1.70)
– One atom REALLY wants the
electrons and the other…not so much
– Ionic bonding  ionic compound
• Smaller difference between values (less
than 1.70)
– Both have “equal” attraction for the e– Covalent bonding  molecule
Differences in Electronegativity
•Medium difference
– Still a bit of a tug of war over e– Unequal sharing of electrons
– Results in a POLAR covalent bond
– Positive and negative poles
– Dipole – partially negative on one
side, partially positive on the other
Polar Covalent Bond
Differences in Electronegativity
• Very small difference
– Share electrons equally
–NONPOLAR covalent bond
– No positive and negative poles
Nonpolar Covalent Bond
Polar vs. nonpolar molecules
• Look at polarity of each bond
– All nonpolar bonds = nonpolar molecule
(O2)
• Look at the overall shape
– Symmetrical polar bonds cancel each other
out so molecule = nonpolar (CO2, CCl4)
– Nonsymmetrical polar bonds = polar
molecule (H2O)
Dipole-dipole attraction
• Attraction between + part of one dipole
and - part of another dipole
• Hydrogen bond - between an
electronegative atom and a hydrogen
atom bonded to another electronegative
atom
– Often involves F, N, or O
– Strongest of the intermolecular forces
Hydrogen Bonding
d+ dH O
+
Hd
Hydrogen Bonding
Van der Waals – London
dispersion force
•Weak intermolecular force caused
by negative electrons on one side
of a cloud being attracted to a
nearby positive nucleus
•Constantly changing
Properties of Molecular Compounds
• Poor conductors of heat & electricity
• Often found as liquids or gases
• Weaker attraction between atoms
• Low melting & boiling points
IONIC vs
COVALENT
Carbon (Organic) Chemistry
• Carbon plays a dominant role in the
chemistry of living things
• Bonding stability
– 4 valence electrons
– Very unlikely to form ionic bonds
– Can form covalent bonds with LOTS of
different elements (especially H & O)
• Small molecules link together resulting in the
formation of a large variety of structures
often with repeating subunits
Examples of carbon-based
compounds
•Simple hydrocarbons
•Small carbon molecules with
functional groups
•Complex polymers
•Biological molecules
Simple Hydrocarbons
• Petrochemicals – Propane, Butane, Octane
Functional groups
• Specific groups of atoms
that are responsible for
chemical characteristics
of a compound
• ALWAYS a close
relationship between
properties & structure
(aspirin, vitamins,
insulin)
Complex Polymers & Biological
Molecules
• Natural polymers
– Proteins, nucleic acids
• Synthetic polymers
– Polythene , Polystyrene
– Kevlar
– Nylon
Common organic molecules
• CH4
• C2H6
• C2H4
• C2H2
• CH3CH2OH
• CH2O
• C6H6
• CH3COOH