C OVALENT C OMPOUNDS
T WO T YPES
OF
B ONDS

Ionic: Electrons are transferred

Covalent: Electrons are shared

Non-polar covalent: equally shared

Polar Covalent: unevenly shared
B OND P OLARITY
R EVIEW:
W HAT IS ELECTRONEGATIVITY ?
ability of an atom to attract electrons
Which element is the most electronegative?
Fluorine
- Has 7 valence e- and wants 8
H
F
P OLAR
BOND
:
covalent bond with greater electron density
around one of the two atoms
electron poor
region
H
electron rich
region
F
e- poor
H
+
d
e- rich
F
d
1
18
2
13
3
4
5
6
7
8
9
10
11
12
14
15
16
17
W HAT TYPE OF B OND IS IT ?
Electronegativity
Difference
Bond Type
0 to 0.3
Nonpolar Covalent
0.4 to 1.6
Polar Covalent
 1.7
Ionic
Increasing difference in electronegativity
Nonpolar
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Classify the following bonds as ionic, polar
covalent,or covalent:
Cs to Cl
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H to S
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4 Polar Covalent
N – 3.0
3.0 – 3.0 = 0 Nonpolar Covalent
Cl to N
Cl – 3.0
D O YOU NOTICE A PATTERN FOR
THE COMBO OF ELEMENTS THAT
ARE IONIC VS COVALENT ?

Ionic bonds form between:

Covalent bonds form between:

Identify the following as ionic, covalent,
or both:
CaCl2
BaSO4
CO2
AlPO4
SO3
H2O
P ROPERTIES OF C OVALENT
C OMPOUNDS

Usually soft and squishy

Not soluble in water

Does not conduct electricity

Soluble in organic solvents

Low melting points

Low boiling points
P ROPERTIES OF I ONIC
C OMPOUNDS

Combination of ions (cation/anion)

Tightly packed solids in a crystal lattice

Hard and Brittle

Usually soluble in water

Conducts electricity when dissolved

High melting points

High boiling points
N AMING C OVALENT C OMPOUNDS
NAMING COMPOUNDS
Nonmetal – Nonmetal
USE PREFIXES!
1.
Change the ending of the second
word to -ide
2.
No mono on the first word
3.
Drop any double vowels
C OVALENT P REFIXES
Number of Atoms
1
Prefix
Mono-
2
3
4
DiTriTetra-
5
6
7
PentaHexaHepta-
8
9
OctaNona-
10
Deca-
E XAMPLES
1.
CO
1. Carbon Monoxide
2.
CO2
2. Carbon Dioxide
3.
SO2
3. Sulfur Dioxide
4.
SO3
4. Sulfur Trioxide
5.
N2H4
5. Dinitrogen Tetrahydride
6.
N2O3
6. Dinitrogen Trioxide
E XAMPLES
1.
disilicon hexafluoride
1. Si2F6
2.
tricarbon octachloride
2. C3Cl8
3.
phosphorus pentabromide
3. PBr5
4.
nitrogen monoxide
4. NO
5.
selenium difluoride
5. SeF2
6.
dihydrogen monoxide
6. H2O
EMPIRICAL AND MOLECULAR
FORMULAS
Define Empirical Formula:
A chemical formula that gives the
simplest whole-number ratio of the
elements in the formula.
Which of the following is an empirical formula?
CO2
C2O4
N2H4
NH2
Define Molecular Formula:
A chemical formula that gives the
actual number of the elements in the
molecular compound.
For the following molecular formulas, write the
empirical formula:
Molecular:
C2H4
C6H12O6
C9H21O6N3
Empirical:
L EWIS S TRUCTURES
O CTET R ULE

Eight electrons in the valence shell (filling s
and p orbitals) make an atom STABLE
This is called the octet rule

Bond formation follows 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 valence energy level.
L EWIS D OT D IAGRAMS
•
an electron-configuration notation with
only the valence electrons of an element
are shown, indicated by dots placed
around the element’s symbol.
•
tracks the number of valence electrons
•
the inner core electrons are not shown
L EWIS D OT P RACTICE
Li
Be
N
O
F
Ne
L EWIS S TRUCTURES FOR
C OMPOUNDS

The pair of dots between two symbols
represents a shared pair.


How many shared pairs does each fluorine have
below?
An unshared pair, also called a lone pair, is
a pair of electrons that is not involved in
bonding and that belongs exclusively to one
atom.
F F
L EWIS S TRUCTURES

The shared pair of electrons is often
replaced by a long dash.
F
F
Each dash represents TWO electrons
W HY SHOULD TWO ATOMS
SHARE ELECTRONS ?
To get a valence of 8 electrons!
+
F
7e-
F
F
7e-
8e-
F
8e-
Lewis structure of F2
single covalent bond
lone pairs
F
F
lone pairs
F
F
single covalent bond
lone pairs
lone pairs
M ULTIPLE C OVALENT B ONDS

double bond:
covalent bond in which two pairs of
electrons are shared between two atoms

shown by two side-by-side pairs of dots
or by two parallel dashes
H
H
C C
H
H
M ULTIPLE C OVALENT B ONDS

triple bond:
covalent bond in which three pairs of
electrons are shared between two atoms

shown by three side-by-side pairs of dots
or by three parallel dashes
L ENGTHS OF C OVALENT B ONDS
Bond
Type
Bond
Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
Bond Lengths
CN
138
Triple bond < Double Bond < Single Bond
CN
116
B OND L ENGTH AND B OND E NERGY

As atomic size increases, bond
length increases, and as a result
bond energy decreases

As you increase the number of
bonds between two atoms, energy
increases, while bond length
decreases.
B OND L ENGTH AND B OND
E NERGY E XAMPLES
1.
Which bond is greater in length: Br2 or F2?
2.
The HF bond is 570 pm, the H2 bond is
436 pm, which bond requires more
energy to break?
3.
Which bond would require more energy
to break C-C single bond or C=C double
bond?
Which bond is longer?
W RITING L EWIS
S TRUCTURES
1.
Draw skeletal structure of compound showing
what atoms are bonded to each other. Put least
electronegative element in the center.
2.
Count total number of valence e-. Add 1 for each
negative charge. Subtract 1 for each positive
charge.
3.
Complete an octet for all atoms except hydrogen
4.
If structure contains too many electrons, form
double and triple bonds on central atom as
needed.
W RITE THE L EWIS STRUCTURE OF
NITROGEN TRIFLUORIDE (NF 3 ).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
F
N
F
F
W RITE THE L EWIS STRUCTURE OF
THE CARBONATE ION (CO 3 2- ).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
O
C
O
O
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24
RESONANCE STRUCTURE :
When there are two or more Lewis structures for a single
molecule
What are the resonance structures of the
carbonate (CO32-) ion?
-
O
C
O
O
-
O
C
O
O
-
-
-
O
C
O
O
-
S OME ELEMENTS DO NOT
FOLLOW THE OCTET RULE
H
Be
H
F
B
F
F
There can also be expanded octets!
M OLECULAR G EOMETRY
VSEPR THEORY

Lewis Dot Diagrams are 2D but we
live in a 3D world.

How are molecules actually
arranged??

Follows the Valance Shell Electron
Pair Repulsion Theory or VSEPR
AB 2 – L INEAR
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
2
0
180˚
Cl
Be
Cl
AB 3 – T RIGONAL P LANAR
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
3
0
120˚
AB 2 E 1 – B ENT
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
2
1
<120˚
AB 4 – T ETRAHEDRAL
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
4
0
109.5˚
AB 3 E 1 – T RIGONAL
P YRAMIDAL
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
3
1
107˚
AB 2 E 2 – B ENT
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
2
2
104.5˚
P REDICTING M OLECULAR
G EOMETRY
1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and
number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
O
S
AB2E
bent
F
O
F
C
F
AB4
F
tetrahedral
I NTERMOLECULAR F ORCES
Intermolecular forces:
attractive forces between molecules.
Intramolecular forces:
attractive forces within a molecule (the
bonds)
Intermolecular
Forces
Intramolecular
Forces
Intramolecular
Forces
intramolecular forces are much stronger than
intermolecular forces
D IPOLES

What is a dipole?

A polar molecule

Uneven sharing of electrons so
there is a separation of charge
D IPOLE -D IPOLE F ORCES

Attraction between two polar molecules
—
+
—
+
H YDROGEN B ONDING

Special type of Dipole – Dipole

Attraction between:
Hydrogen and Nitrogen/Oxygen/Fluorine
D IPOLE – I NDUCED D IPOLE

Electrons shift
toward
positive end
of dipole
Attraction between one polar and one
nonpolar molecule
—
—
+
+
—
+
L ONDON D ISPERSION F ORCES

Attraction between two nonpolar molecules
Electrons
become
uneven and
form a dipole
—
+
—
+
S TRENGTH
OF
IMF

Hydrogen Bond

Dipole – Dipole

Dipole – Induced Dipole

London Dispersion Forces
strongest
weakest
Which of the following molecules is polar?
H2O, CO2, SO2, and CH4
O
S
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
What type(s) of intermolecular forces exist
between each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are
also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces.
S
SO2
SO2 is a polar molecule: dipole-dipole forces. There are
also dispersion forces between SO2 molecules.
W HAT
DOES
IMF

Viscosity

Surface Tension

Cohesion/Adhesion

Boiling Point
EFFECT ?
V ISCOSITY

Measures a fluid’s resistance to flow
Stronger IMF  Higher Viscosity
S URFACE T ENSION

result of an imbalance of forces at the
surface of a liquid.
Stronger IMF  Higher Surface Tension
B OILING P OINT

Point at which liquid particles escape the
surface of the liquid into the gas phase
Stronger IMF  Higher Boiling Point
A DHESION

AND
C OHESION
Cohesion:
intermolecular attraction between like molecules

Adhesion:
intermolecular attraction between unlike molecules
Adhesion
Cohesion