Covalent Bonding and
Naming
Chemistry 11
Borrowed from Mrs. Kay
Read Pages 168-171, 185-196
Pure Covalent Bonding



Equal sharing of electron between two of the
same non-metals
The electronegativity between the two atoms
must ZERO!
Ex: hydrogen gas (H2)
HOBrFINCl


These elements naturally form as diatomic
molecules (2 atoms bonded covalently)
Hydrogen (H2), Oxygen (O2), Bromine (Br2),
Fluorine (F2), Iodine (I2), Nitrogen (N2), and
Chlorine (Cl2)
Single Bonds




Sharing of 2 electrons, one pair
A single line represents the 2 electrons
Longest and the weakest of the covalent bonds
(easiest to break apart)
Ex: Fluorine, F2
Double Bonds

Sharing 4 electrons



2 pairs of electrons
Shorter and stronger
than a single bond,
takes more energy to
break it apart.
Drawn with two lines,
each line represents 2
electrons
Triple Bonds
Sharing 6 electrons, 3 pairs of electrons
 Shortest and strongest of the covalent
bonds
 Ex: nitrogen, N2

Naming simple molecules
 If its diatomic (HOBrFINCl)
you simply name the nonmetal its made of Must
memorize the prefixes
 RULES: if there is only one
of the first atom than don’t
use a prefix, otherwise use
a prefix.
 Ex: CO = carbon monoxide
 Ex: P2O4 = diphosphorous
tetroxide
Prefix
Number
Mono
1
Di
2
Tri
3
Tetra
4
Penta
5
Hexa
6
Hepta
7
Octa
8
Nona
9
Deca
10
Practice:
1.
2.
3.
4.
5.
6.
CO2
NH3
BF3
NO2
N2H4
N2F2
1.Carbon dioxide
2.Nitrogen trihydride
3.Boron trifluoride
4.Nitrogen dioxide
5.Dinitrogen tetrahydride
6.Dinitrogen difluoride
Polar Covalent Bonding




Electrons are shared unequally.
They are not ionic, because the electron is not totally
removed because there was not enough attraction to totally
remove the electron.
Based on difference in electronegativity
Ex: HCl
Electronegativity
The degree to which an atom attracts
electrons to itself
 It is not a measurement, but a scale.
 Periodic trend: generally increases from
left to right across a period and from
bottom to top in a group.
 What is the most electronegative element?
 Fluorine (F)

The greater the electronegative difference,
the more polar the bond because the more
electronegative atom will attract the
“shared” electron pair closer to itself.
Ex: H-N = more polar covalent than H-C,
because electronegativity of H (2.20) and
N (3.04), while C (2.55)
H-N: ΔEN = 3.04-2.20 = 0.84 (more polar)
H-C: ΔEN = 2.55 – 2.20 = 0.35

Bonding Continuum
Idea that bonds can behave as mostly
ionic or mostly covalent
 Use the difference in electronegativity to
label the bond type

0 = pure covalent
 0.4 to 1.7 = mostly covalent
 1.7 or greater = mostly ionic

Lewis Structures for molecules


Need to show the structure of a molecule.
Will use Lewis structures (electron dot diagrams)
to show where there are lone pairs (filled
orbitals) and bonding pairs (places where bonds
most likely occur)
Lewis Structures
1.
2.
3.
Look at valence electrons of all atoms
Pick a central atom (least electronegative
usually, has most bonding sites)
Align all atoms so that each have their
ideal amount of valence electrons
achieved through sharing.

Usually 8 (stable octet), but can be 2 (H, He)
and 6 (B)
Carbon tetrachloride
Carbon is the central atom.
 It has 4 bonding pairs.
 Chlorine wants to share
one bonding site each.
 Need 4 chlorines for every
one carbon
(Cl has 3 lone pairs and 1
bonding pair)

Some examples
Practice drawing and naming Lewis
Structures


H2O
CH2O
What about ions?
Count up all valence electrons that you
are allowed to place.
 Still pick the central atom.
 Still have the correct number of electrons
around each atom (usually 8, except for H
and He)
 Add extra electrons if an anion and take
away electrons if a cation


TRY [CO3]2- and [NH4+]
Metallic Bonding
What is a Metallic Bond?
- A metallic bond occurs in
metals. A metal consists of
positive ions surrounded by a
“sea” of mobile electrons.
This
shows
what a
metallic
bond
might
look
like.
Name 4 Characteristics of a
Metallic Bond.
1. Good conductors of heat
and electricity
2. Great strength
3. Malleable and Ductile
4. Luster