Post on 11-Jul-2015
Covalent bonds form when atoms share
electrons to complete octets.
Covalent bonds are typically between two
nonmetal atoms.
2/14/13
CHM 101 Ch 4: Covalent Compounds
Covalent Compounds
1. Element with lower group number is
named first.
2. If elements are in the same
group, element with higher period
number is named first.
3. Second element named as root + -ide.
4. Greek prefixes are used to designate
the number of atoms.
Binary Covalent Compounds – molecule that contains atoms of only 2 elements
CS2, SO3, PCl5, etc.
General Rules for Naming
You must know these!
Naming Covalent Compounds
What is the name of SO3?
1. The first nonmetal is S sulfur.
2. The second nonmetal is O named oxide.
3. The subscript 3 of O is shown as the prefix tri.
SO3 sulfur trioxide
The subscript 1 (for S) or mono is understood.
Naming Covalent Compounds
1. PCl5
2. CS2
3. P4S3
Give names for the following binary compounds:
Give molecular formulas for the following compounds:
1. Dinitrogen monoxide
2. Selenium hexafluoride
3. Dichlorine heptaoxide
Naming Covalent Compounds
phosphorous pentachloride
carbon disulfide
tetraphosphorous trisulfide
N2O
SeF6
Cl2O7
Use the table below to quiz yourself. Use the formula, write the name. Use
the name, write the formula.
Naming Covalent Compounds
Covalent Bonding
Octet Rule: Main group elements gain, lose, or share e- to achieve a
stable e- configuration with 8 valence e- (except H and He – only need
2 valence e- for stability)
Main group elements
Usually make the number of bonds necessary to have noble gas
configuration
Noble gases (except He) have 8 valence electrons
Lewis Dot Symbol - Representation of the number of valence electrons
in an atom. Usually only used for main group elements. Maximum
number is 8.
X
H He
Li Be B C N O F Ne
Covalent Bonding
Cl2 molecule
Cl: 7 valence electrons, needs to make 1 bond to have octet
Cl Lewis Dot Symbol
Lone pair of e- (unshared)
ClCl
Bonding pair of e- (shared)
ClCl
ClCl
Lewis Dot Structure
Count the “shared” e- for both atoms
Covalent Bonding
Lewis Dot Structure – shows
how the atoms of a molecule
are connected. Shows lone
pairs of electrons and bonding
pairs of electrons.
In carbon dioxide, CO2, the C atom shares 4 electrons with
each O atom in a double bond.
Covalent Bonding
Multiple Bonds – Double & Triple
In nitrogen molecule, N2, each N atom shares 6 electrons
with the other N atom in a triple bond.
Multiple Bonds – Double & Triple
Covalent Bonding
Group 7(A) 7 1 HCl
Group 6(A) 6 2 H2O
Group 5(A) 5 3 NH3
Group 4(A) 4 4 CH4
# Valence e- # Bonds for Octet Example
We can often predict the number of covalent bonds an atom will form
based on the number of valence electrons.
ClH OCl HOH
Covalent Bonding
2. Arrange atoms next to each other and connect with bonds
Central atom is usually written first in the formula and has lower group
number
(If atoms are in same group, central atom is from higher period)
Rules for Drawing Lewis Dot Structures
1. Count the total number of valence electrons in the molecule
NF3 N: 5 e-
F: 7e- x 3 = 21 e-
26 val e- total for the molecule
N F
F
F26 val e-
- 6 bonding e-
20 remaining e-
Covalent Bonding
Rules for Drawing Lewis Dot Structures
3. Place lone pairs around each atom to satisfy octet rule, starting with
terminal atoms
20 remaining e-
-20 lone e-
0 remaining e-
N F
F
F
Make sure each atom has an octet. If it doesn’t, check your work.
Covalent Bonding
Practice: Draw Lewis Structures for CH3Br and H2Se.
If the central atom has < 8 e-…
5. Change single bonds to multiple bonds (double or triple) using lone
pairs from terminal atoms.
H2CO, CH3COOH, HCN
Covalent Bonding
Rules for Drawing Lewis Dot Structures
IF there are electrons remaining…
4. Place leftover e- on the central atom. It is ok to exceed the octet
(have more than 8 electrons on an atom) if the central atom is in
Period 3 or higher.
SF4
Draw the Lewis structure for formaldehyde: H2CO
1. Count the total number of valence electrons in the molecule
H2CO H: 1 e- x 2 = 2 e-
C: 4e- x 1 = 4 e-
O: 6e- x 1 = 6 e-
12 val e- total for the molecule
2. Place atoms relative to each other and connect with bonds
H can only make 1 bond, so it has to be a terminal atom. C is a
very common central atom (lower group # than O).
12 val e-
- 6 bonding e-
6 remaining e-
C HH
O
Covalent Bonding
Draw the Lewis structure for formaldehyde: H2CO
3. Place lone pairs around each atom to satisfy octet rule, starting with
terminal atoms
H can only accommodate 2 electrons, which it has with the bonding
electrons. Never put lone pairs on H.
6 val e-
- 6 bonding e-
0 remaining e-
C HH
O
The central atom (carbon) has < 8 e-
4. Change single bonds to multiple bonds (double or triple) using lone
pairs from terminal atoms.
C HH
O
Carbon and oxygen now both have
“octets”.
Covalent Bonding
Lewis Structures of Ions
For anions – add electrons to total valence equal to charge
NO3– : N: 5e- x 1 = 5
O: 6e- x 3 = 18
+ 1e- (b/c of ion charge)
24 e- total
- 6 e- bonding
18 e-
- 18 e- lone pairs
0 e-O N O
O
O N O
O– Does each atom have an octet?
For ions, always put the Lewis structure in
brackets and write the charge as superscript.
How can we give N an octet?
For cations – subtract electrons from total valence equal to charge
NH4+ : N: 5e- x 1 = 5
H: 1e- x 4 = 4
- 1e-
8 e- total
- 8 e- bonding
0 e-
H N H
H
H+
Examples for you to practice: ClO3–, ClF4
+
Lewis Structures of Ions
Exceptions to the Octet Rule
1. Fewer than 8 electrons: Molecules with Be or B as central atom are often electron deficient. Be usually only needs 4 electrons for stability and B only needs 6.
BeCl2 BF3
Covalent Bonding
2. Odd # of valence electrons:
NO2:
Free radicals: atoms or molecules with unpaired electrons, highly reactive
3. More than 8 valence electrons: “expanded valence shells”
- only allowed for atoms in Period 3 and beyond
H2SO4
Exceptions to the Octet Rule
Covalent Bonding
O S O
O
O
H H
Example for you to practice: PCl5
Draw the Lewis Structure for ozone, O3:
There are 2 possible structures only differing in the location of electrons and the double bond.
Experimental data says that both bonds are identical.
The actual structure of O3 is neither of them, but a composite of the two, called a resonance hybrid. Each Lewis structure is a resonance structure of O3.
Covalent Bonding
Resonance Structures
O
O+
O-
O
O+
O-
Use a double-headed arrow for resonance structures
O
O+
O-
O
O+
O-
• Resonance structures differ only in the assignment of e- pair positions, not in atom positions.
Draw resonance structures for NO3-:
Covalent Bonding
Resonance Structures
Valence-Shell Electron-Pair Repulsion (VSEPR) Theory
Bonding and lone pair electrons surrounding a central atom repel each
other. To minimize the repulsions, electron groups are oriented as far
apart as possible.
Molecular Shape – VSEPR Theory
The properties of a molecule are heavily influenced by molecular shape.
How many electron “groups” are on a central atom? Each of the following
counts as one electron group:
• Lone pair of e-
• Single bond
• Double bond
• Triple bond
Covalent Bonding
Molecular Shape – VSEPR Theory
Covalent Bonding
Ideal bond angles
Covalent Bonding
Electron pair geometry – Arrangement of e- groups (bonds and lone pairs) around the central atom.
Molecular geometry – Arrangement of atoms in space, shape of molecule. This is different than e- group geometry if lone pairs are present
Linear electron pair geometry
2 atoms attached to a central atom
180 bond angle
Can’t have any lone pairs, so electron pair geometry is same as molecular geometry
Examples: BeCl2, CO2, CS2, HCN
If there are 2 electron “groups” on the central atom…
Linear molecular geometry
Covalent Bonding
Molecular Shape – VSEPR Theory
Trigonal Planar electron pair geometry
If 0 lone pairs, electron pair geometry is same as molecular geometry
Examples: BF3, SO3, NO3–, CO3
2–
3 atoms attached to a central atom
120 bond angle
If there are 3 electron “groups” on the central atom and 0 lone pairs…
Trigonal Planar molecular geometry
Covalent Bonding
Molecular Shape – VSEPR Theory
Tetrahedral electron pair geometry
Electron pair geometry
Molecular geometry
Example
4 atoms attached to a central atom
109.5 bond angle
If 0 lone pairs, electron pair geometry is same as molecular geometry
Examples: CH4, SiCl4, SO42–, ClO4
–
If there are 4 electron “groups” on the central atom & 0 lone pairs…
Tetrahedral molecular geometry
Covalent Bonding
Molecular Shape – VSEPR Theory
Electron pair geometry
Molecular geometry
Example
3 atoms & 1 lone pair attached to a central atom
109.5 bond angle
Examples: NH3, PF3, ClO3–, H3O
+
Tetrahedral electron pair geometry
If there are 4 electron “groups” on the central atom & 1 lone pair…
Trigonal pyramidal molecular geometry
Covalent Bonding
Molecular Shape – VSEPR Theory
Electron pair geometry
Molecular geometry
Example
2 atoms & 2 lone pairs attached to a central atom
109.5 bond angle
Examples: H2O, SCl2
Tetrahedral electron pair geometry
If there are 4 electron “groups” on the central atom & 2 lone pairs…
Bent (Angular) molecular geometry
Covalent Bonding
Molecular Shape – VSEPR Theory
Covalent Bonds & Electronegativity
Electrons are shared in covalent bonds, but they usually are not shared
equally.
One atom usually pulls the electrons more strongly than the other.
The bond between H and Cl
in HCl is covalent.
Cl pulls the shared electrons
more strongly than H.
This creates a “partially
negative” region around Cl
and a “partially positive”
region around H.
The bond between H and Cl is a polar covalent bond.
Covalent Bonding
WHY???
Covalent Bonds & Electronegativity
Covalent Bonding
The electronegativity value for an atom indicates the attraction of that atom
for shared electrons (in covalent bonds).
Electronegativity strength
usually increases as the size of
an atom decreases.
A nonpolar covalent bond is an equal or almost equal sharing of electrons.
The atoms involved have almost no electronegativity difference (0.0 to 0.4).
Examples:
Electronegativity
Atoms Difference Type of Bond
N-N 3.0 - 3.0 = 0.0 Nonpolar covalent
Cl-Br 3.0 - 2.8 = 0.2 Nonpolar covalent
H-Si 2.1 - 1.8 = 0.3 Nonpolar covalent
H-C ??? ???
Nonpolar Covalent Bonds
Covalent Bonding
Diatomic molecules are molecules that
contain 2 atoms of the same element.
They are the natural state for elements
H, O, N, Cl, Br, I, and F.
These are the only truly nonpolar covalent
bonds.
Nonpolar Covalent Bonds
Covalent Bonding
A polar covalent bond is an unequal sharing of electrons.
The atoms involved have a moderate electronegativity difference (0.5 to 1.7).
Examples: Electronegativity
Atoms Difference Type of BondO-Cl 3.5 - 3.0 = 0.5 Polar covalentCl-C 3.0 - 2.5 = 0.5 Polar covalent
O-S 3.5 - 2.5 = 1.0 Polar covalent
Polar Covalent Bonds
Covalent Bonding
An ionic bond occurs between metal and nonmetal ions, and is a result of
electron transfer from the metal to the nonmetal.
There is a large electronegativity difference (1.8 or more) between the
atoms.
Examples:
Electronegativity
Atoms Difference Type of Bond
Cl-K 3.0 – 0.8 = 2.2 Ionic
N-Na 3.0 – 0.9 = 2.1 Ionic
S-Cs 2.5 – 0.7= 1.8 Ionic
Ionic Bonds
Covalent Bonding
Covalent Bonding
A covalent bond is polar if it connects atoms with different electronegativity
values, i.e. H-Cl, C-F, C-Cl, etc.
How do you know if a molecule is polar?
- Bond Polarity
- Molecular Shape
Covalent Bonding
Molecular Polarity
Determine the molecular polarity of CO2:
C OO
Each C=O bond is polar, but the molecule is linear and the two
dipoles cancel each other. Therefore, the CO2 molecule is nonpolar.
COS (carbonyl sulfide):
C SO
C and S have the same EN. There is only 1 dipole, pointing towards
O. Therefore, the COS molecule is polar.
Covalent BondingMolecular Polarity
What about molecular shape?
H2O:
O HH
Is the molecule nonpolar?
O
H H
NO! The H2O molecule is not linear, it is tetrahedral!
Covalent Bonding
Molecular Polarity
Indicate the dipole moment (if any) for CF4 and CHF3.
Covalent Bonding
Molecular Polarity
Indicate the dipole moment (if any) for CH2Cl2.
Covalent Bonding
Molecular Polarity