CHAPTER 6: COVALENT
COMPOUNDS
6.1 Covalent Compounds
What is a covalent bond?
• Sharing of electrons between atoms
• Electron orbitals overlap to form molecular orbitals
• Molecular orbitals are the region where the shared
electrons move
• Covalent bonds are also called molecular bonds.
How is a molecular bond different from an
ionic bond?
• Ionic Bond: Transfer of electrons
• Covalent or Molecular Bond: Sharing of electrons
Bond Length
• Average distance between nuclei of two atoms
• Not Fixed: atoms vibrate; bond acts like a spring
Bond Energy
• Energy needed to break a bond
• As bond energy goes up, bond length goes down (inverse
relationship)
H
Br
141.1 pm
Bond energy = 366 kJ/mol
kJ/mol
H
F
91.7 pm
Bond energy = 570
Bond
Type
Electron Bond
s
Length
Shared
Single
2
Long
Double
4
Triple
6
Short
Bond
Energy
Low
High
Electronegativity
• Difference in electronegativities can be used to determine
bond type
Predicting Bond Type from
Electronegativity Difference
• Nonpolar Covalent: Electronegativity Difference < 0.5
• Polar Covalent: Electronegativity Difference ≥ 0.5 and ≤
2.1
• Ionic: Electronegativity Difference > 2.1
Types of Covalent Bonds
1. Nonpolar Covalent Bond
• Bonding electrons shared evenly
• Difference in electronegativity is less than 0.5
• Examples:
C–H
H–H
Types of Covalent Bonds
2. Polar Covalent Bond
• Bonding electrons not shared evenly
• Difference in electronegativity is 0.5 to 2.1
• Tend to be stronger than nonpolar covalent bonds
• Electrons tend to be closer to more electronegative atoms
• Example: H – F
Polar Covalent Bonds
• Polar covalent bonds have a dipole.
• Dipole: one end has a partial positive charge and the other end
has a partial negative charge.
• δ+ is used to indicate a partial positive charge
• δ- is used to indicate a partial negative charge
• Examples:
Ionic Bonds
• Metal + Nonmetal
• Electronegativity difference is greater than 2.1
• Example: Li – F
Summary
• Nonpolar Covalent: Electronegativity Difference < 0.5
• Polar Covalent: Electronegativity Difference ≥ 0.5 and ≤
2.1
• Ionic: Electronegativity Difference > 2.1
Determine if the following compounds are
ionic, polar covalent, or nonpolar covalent:
1.
Na – F
2.
Ca – O
3.
C – Cl
4.
Al – Cl
5.
Br – Br
Determine if the following compounds are
ionic, polar covalent, or nonpolar covalent:
1.
2.
3.
4.
5.
Na – F = ionic
(electronegativity difference = 3.1)
Ca – O = ionic
(electronegativity difference = 2.4)
C– Cl = polar covalent
(electronegativity difference = 0.6)
Al – Cl = polar covalent
(electronegativity difference = 1.6)
Br – Br = nonpolar covalent
(electronegativity difference = 0)
Properties of Covalent Compounds
1.
2.
3.
4.
5.
Low melting and boiling points
Soft and squishy
Tend to be flammable
Do NOT conduct electricity in water
Are NOT very soluble in water
6.2 Drawing Covalent Compounds: Lewis
Structures
Gilbert N. Lewis (1875 – 1946)
• Gilbert N. Lewis created Lewis structures
• Lewis structures (electron-dot diagrams)
• Structural formula for drawing molecules
• Dots = electrons
• Lines = electron pairs/bonds
• Only valence electrons are shown
Lewis Structures
• These are examples of Lewis structures
H
Be
C
O
1. Lewis Structures of Atoms
• A Lewis diagram consists of the element symbol in the
center and dots going around it representing the valance
electrons.
• The dots representing the valance electrons must be evenly
distributed around all 4 sides of the symbol.
• Group 2 elements and Helium are the exception to the rule
of even distribution.
• When placing dots around the element symbol, you should
start on the left side and go clock wise until you run out of
electrons.
• Nitrogen has 5 Valence electrons
N
N
N
N
N
Practice Questions
Draw a dot diagram of the following elements:
• Potassium
• Boron
• Chlorine
• Neon
• Lead
• Oxygen
K
Pb
B
Cl
Ne
O
Group 2 and Helium
-Group 2 and Helium are the exceptions because their
valence electrons are in the first shell.
-Therefore the 2 valence electrons for Helium and Group 2
elements need to be put on the same side of the element
symbol.
He
Why do Lewis structures help us?
• These diagrams allow us to understand how atoms will
bind based on their valance electron configuration.
Bonding of atoms
• When you have unpaired electrons on a side of an atom that
atom will want to bond/share electrons with other atoms
• This is because all atoms want a full valance shell
• Oxygen has two unpaired electrons
• Hydrogen has one free electron so 2 H can bond with O, this
make H O
2
O
H
O H
H
2. Drawing Lewis Structures of
Compounds
• Step 1: Count the total number of valence electrons in the
compound.
Example: CH3I
• Step 2: Draw a Lewis structure for each atom in the
compound.
Example: CH3I
14 valence electrons
• Step 3: Arrange the atoms
• Halogen and Hydrogen atoms often bond to only 1 other atom and are
usually at the end of the molecule
• Carbon is often in the center of the molecule
• With the exception of carbon, the atom with the lowest electronegativity is
often the central atom
Example: CH3I
14 valence electrons
• Step 4: Distribute the dots
• Distribute the electron dots so that each atom (except H, Be, and
B) satisfy the octet rule
Example: CH3I
14 valence electrons
• Step 5: Draw the bonds
• Change each pair of dots that represent a shared pair of electrons
to a long dash.
Example: CH3I
14 valence electrons
• Step 6: Verify the structure
• Count the electrons in your structure
Practice
1.) H2S
2.) NH3
3.) BF3
More Practice
4.) C2H6
5.) CH2Cl2
6.) CH3OH
7.) SCl2
8.) CHF3
3. Lewis Structures with Polyatomic Ions
2
4
SO
H3O

OH

More Practice
NH

4

4
PCl

IO
CH

3
3. Lewis Structures of Compounds with
Multiple Bonds
C, O, N often form double bonds by sharing 2 pairs of
electrons.
C, N may even form triple bonds by sharing 3 pairs of
electrons.
Practice
1.) O2
2.) C2H4
3.) CO2
4.) N2
More Practice
5.) CO
6.) C2H2
7.) HCN
8.) HC2Cl
9.) N2F2
4. Resonance Structures
-When a molecule has two or more possible Lewis
structures, the two structures are called resonance
structures.
-Place double-headed arrows between the structure to
show that the actual molecule is an average of the two
possible states.
Practice
1.) SO2
2.) O3
3.) NO2-
6.3 Naming Covalent Molecules
WRITING FORMULAS
Covalent Compounds
Covalent Compounds
• Formed from 2 nonmetals or
1 metalloid and 1 nonmetal.
Prefix
Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
Number of Atoms
1
2
3
4
5
6
7
8
9
10
Writing Formulas for Covalent
Compounds
1.
Write the symbol for the cation.
2.
From the prefix, write the number of atoms as
a subscript behind the symbol.
3.
Write the symbol of the anion.
4.
From the prefix, write the number of atoms as
a subscript behind the symbol.
5.
DO NOT use charges.
Examples
1.
Carbon Monoxide = CO
2.
Tetraphosphorus decaoxide = P4O10
3.
Carbon dioxide
4.
Sulfur dioxide
Examples
1.
Carbon Monoxide = CO
2.
Tetraphosphorus decoxide = P4O10
3.
Carbon dioxide = CO2
4.
Sulfur dioxide = SO2
NAMING COMPOUNDS
Covalent Compounds
Covalent Compounds
•
Formed from a metalloid and a nonmetal or two
nonmetals.
Prefix
Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
Number of Atoms
1
2
3
4
5
6
7
8
9
10
1. Use prefixes: mono, di, tri, tetra, penta, hexa, hepta,
octa, nona, and deca.
2. Add a prefix to the cation (from the number of atoms).
Exception: the prefix mono is not added to the cation.
4. Add a prefix to the anion (from the number of atoms).
5. Change the ending of the anion to –ide.
Examples
1.
P4S10 = tetraphosphorus decasulfide
2. N2O = dinitrogen monoxide
3. P2O3
4. IF5
Examples
1.
P4S10 = tetraphosphorus decasulfide
2. N2O = dinitrogen monoxide
3. P2O3 = diphosphorus trioxide
4. IF5 = iodine pentafluoride
6.4 Molecular Shapes: VSEPR
• Objective:
1.
To predict the shape of a molecule using VSEPR
theory.
VSEPR
•V
•S
•E
•P
•R
Valence
Shell
Electron
Pair
Repulsion
• VSEPR: theory that predicts some molecular shapes
based on the idea that pairs of valence electrons
surrounding an atom repel each other
How do we figure out the shape of a
molecule?
• The shape of a molecule is determined by the valence
electrons surrounding the central atom.
The Molecular Shapes…
Tetrahedral
Bent
Trigonal Pyramidal
Linear
Trigonal Planar
Examples
How do we determine the molecular
shape?
Step 1: Draw the Lewis structure for the
molecule.
Example: NH3
Step 2: Count the number of “things” on
the atom you’re interested in (the center
atom)
• Count atoms and lone pairs, NOT BONDS!
Example: NH3
three atoms
one lone pair
+________________
4 total “things”
Step 3: Count the number of lone pairs
that are on the atom you’re interested in.
• Lone pairs on other atoms aren’t important; only count
what is directly stuck to the atom you’re interested in.
Example: NH3
one lone pair
VSEPR Flowchart
Practice
1.) CH4
2.) CO2
3.) carbon tetrabromide
4.) phosphorus trichloride
5.) boron trichloride
6.) sulfur difluoride
Practice
1.) CH4
tetrahedral
2.) CO2
linear
3.) carbon tetrabromide: tetrahedral
4.) phosphorus trichloride: trigonal pyramidal
5.) boron trichloride: trigonal planar
6.) sulfur difluoride: bent
Polarity
• A molecule is polar if it has an overall dipole.
How does shape affect polarity?
• Examples:
• CO2: linear
• No molecular dipole: Not Polar
• H2O: bent
• Molecular dipole: Polar
Are the following molecules polar or
nonpolar?
1.) CF4
2.) SF2
3.) NF3
4.) BeCl2
Are the following molecules polar or
nonpolar?
1.) CF4
Nonpolar
2.) SF2
Polar
3.) NF3
Polar
4.) BeCl2
Nonpolar