chapter 8: covalent bonding - wunder...
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Chapter 8
Chapter 8: Covalent Bonding
Chapter 8
• Two methods to gain or lose valence electrons:
– Transfer of Electrons = Ionic Bonding – Sharing of Electrons = Covalent Bonding
Bonding
Ionic Bonding - attracted
to each other, but not fully
committed
Covalent Bonding - fully
committed, and shares
everything
Chapter 8
• Prefixes used to denote number of atoms of element
• The suffix -ide is added to the more electronegative
element
• For the least electronegative (on the left), mono- is not
used to denote one atom.
Naming Covalent Compounds
Numerical Prefixes
mono- 1 hexa- 6
di- 2 hepta- 7
tri- 3 octa- 8
tetra- 4 nona- 9
penta- 5 deca- 10
Chapter 8
Molecules
• Molecule = covalently bonded atoms
• Diatomic molecule = two atoms
• Molecular compounds – lower melting and
boiling point than ionic compounds
• Molecular formula
– Doesn’t have to be lowest whole-number ratio
– Does represent structure
Chapter 8
Molecular Representations
Figure 8.5, Pg. 215, Text
Chapter 8
Lewis dot structures of covalent
compounds
• In covalent compounds atoms share electrons.
We can use Lewis structures to help visualize
the molecules.
Lewis structures
•Multiple bonds must be considered.
• Will help determine molecular geometry.
• Will help explain polyatomic ions.
Chapter 8
Types of electrons
• Bonding pairs
• Two electrons that are shared between two
atoms. A covalent bond.
• Unshared pairs
• A pair of electrons that are not shared
between two atoms. Lone pairs or nonbonding
electrons.
H Cl
oo
oo
oo
oo
Bonding pair
Unshared
pair
Chapter 8
Single covalent bonds
H H
H
C H H
H
Do atoms (except H) have octets?
F F
Chapter 8
Covalent Bonding
• When two similar atoms bond,
none of them wants to lose or
gain electrons
– Share pairs of electrons to each
obtain noble gas e-
configuration.
• Each pair of shared electrons
= one covalent bond
• Unshared Pairs = Pairs of e-
not shared by all atoms
– Show unshared pairs as dots
Visual, Pg. 219, Text
N H
H H
Chapter 8
Sharing of Electrons
• Water - H2O - Covalent Bonding – H = 1e-, O = 6e-
– Each H shares their electron
with O O = 8e- = [Ne]
– O shares 1e- with each H
H = 2e- = [He]
– Use dashes or dots to show
covalent bonds
Visual, Pg. 218, Text
H H O
Chapter 8
Multiple Covalent Bonds
• Elements can share more
than two electrons - creates
double and triple bonds
– Carbon Dioxide (CO2) - 4
electrons shared between the
carbon and the oxygens
(double bond)
O C O
O C O
Chapter 8
– Triple bond = Six electrons shared between the carbon atoms
– Ethyne (C2H2), a.k.a acetylene
– Bond length decreases from single to triple bonds
– Bond strength increases from single to triple bonds
Triple Covalent Bonds
C C H H C C H H
Chapter 8
Coordinate Covalent Bonds
• Coordinate covalent bond = one atom
contributes both bonding electrons
• Polyatomic ions contain coordinate covalent
bonds
C O O = 8e-
C = 6e-
O must contribute one of its pairs C O O = 8e-
C = 8e-
Chapter 8
O
OO
Resonance Structures
• Structures with multiple bonds can have similar
structures with the multiple bonds between different pairs
of atoms = Resonance Structures
• Example: Ozone has two identical bonds whereas the
Lewis Structure requires one single (longer) and one
double bond (shorter).
Chapter 8
Resonance Structures
Resonance Structures
Chapter 8
• Sometimes we can have two or more
equivalent Lewis structures for a molecule.
O - S = O O = S - O
• They both - satisfy the octet rule
- have the same number of bonds
- have the same types of bonds
• Which is right?
Resonance structures
Chapter 8
• They both are!
O - S = O O = S - O
O S O
This results in an average of 1.5 bonds between
each S and O.
Resonance structures
Chapter 8
• Three classes of exceptions to the octet rule for
molecules
• Odd number of electrons
• One atom has less than an octet
• One atom has more than an octet
• Odd Number of Electrons
• Few examples.
• Molecules such as ClO2, NO, and NO2 have an odd
number of electrons.
N O N O
Exceptions to Octet Rule
Chapter 8
• Less than an Octet
• Relatively rare.
• Typical for compounds of Groups 13.
• Most typical BF3.
• More than an Octet
• Atoms from the 3rd period onwards can accommodate
more than an octet.
• The d-orbitals are low enough in energy to participate in
bonding and accept extra electron density.
Exceptions to Octet Rule
Chapter 8
Atoms with fewer than eight electrons
• Beryllium and boron will both form
compounds where they have less than
8 electrons around them.
:Cl:Be:Cl: : :
: : :F:B:F:
:F:
: :
: :
: :
Chapter 8
Species with an odd
total number of electrons
• Example - NO
• Nitrogen monoxide is an example of a
compound with an odd number of electrons.
• It has a total of 11 valence electrons: six
from oxygen and 5 from nitrogen.
• The best Lewis structure for NO is:
:N::O:
: .
Chapter 8
Drawing Lewis structures
• Write the symbols for the elements in the correct
structural order.
• Calculate the number of valence electrons for all atoms
in the compound.
• Put a pair of electrons between each symbol, the bond
between each.
• Beginning with the outer atoms, place pairs of electrons
around atoms until each has eight (except for hydrogen).
• If an atom other than hydrogen has less than eight
electrons, move unshared pairs to form multiple bonds.
Chapter 8
Lewis Dot Structures: SO42-
• 1. Write a possible
– arrangement.
• 2. Total the electrons. • 6 from S, 4 x 6 from O
• add 2 for charge
– total = 32
• 3. Spread the electrons
– around.
S O
O
O
O
- - |
|
S O
O
O
O
Chapter 8
Lewis structures
• Example CO2
• Step 1
– Draw any possible structures
C-O-O O-C-O
You may want to use lines for bonds.
Each line represents 2 electrons.
Chapter 8
Lewis structures
• Step 2
– Determine the total number of valence electrons.
– CO2 1 carbon x 4 electrons = 4
2 oxygen x 6 electrons = 12
Total electrons = 16
Chapter 8
Lewis structures
• Step 3
– Try to satisfy the octet rule for each atom
• - all electrons must be in pairs
• - make multiple bonds as required
Try the C-O-O structure
No matter what you
try, there is no way
satisfy the octet for
all of the atoms.
C O O
Chapter 8
Lewis structures
O=C=O
This arrangement needs
too many electrons.
How about making some double bonds?
That works!
O C O
= is a double bond,
the same as 4 electrons
Chapter 8
Ammonia, NH3
H
H N H
Step 1 Step 2 3 e- from H
5 e- from N
8 e- total
Step 3 N has octet
H has 2 electrons
(all it can hold)
H
H N H
Chapter 8
Electronegativity
• Electronegativity = Ability to attract
electrons in a chemical bond
– Decrease as you move down a group
– Increase as you move from left to right
across a period.
Decreasing
Electro-
negativity
Increasing
Electro-
negativity
Chapter 8
Polar Covalent Bonds
• Sharing of electrons in a covalent bond does not imply equal sharing of those electrons.
• In some covalent bonds - electrons located closer to one atom than the other
• Unequal sharing polar bonds.
• Sharing based on electronegativity of elements in bond
Chapter 8
Electronegativities
Chapter 8
Electronegativity and Polarity
• Difference in electronegativity between atoms is
a gauge of bond polarity
Difference Type of Bond Electrons
0 – 0.4 Nonpolar Covalent Equal sharing
0.4 – 1.0 Moderately Polar
Covalent
Unequal sharing
1.0 – 2.0 Very Polar
Covalent
Unequal sharing
3 Ionic Transfer
Nonpolar and polar covalent bonds
• Nonpolar: When two atoms share a pair of
electrons equally.
• H H Cl Cl
• Polar: A covalent bond in which the electron
pair in not shared equally.
H Cl
• Note: A line can be used to represent a
shared pair of electrons.
d+ d-
oo
oo
oo
oo
oo
oo
oo o
o
oo
oo
oo
oo
Chapter 8
• The positive end (or pole) in a polar
bond is represented d+ and the
negative pole d- = dipoles (partial
charges)
• Polar molecules placed between
electric (+/-) plates – become
aligned with plates
Polar Dipoles
H H O
δ+ δ+
δ-
Chapter 8
Attractions between Molecules
• Intermolecular attractions weaker than ionic or
covalent bonds
• these are attractions BETWEEN molecules not
inside molecules like ionic or covalent bonds
(intramolecular)
Chapter 8
Types of Intermolecular Attractions
– van der Waals forces
• Dipole interaction
– δ+ end of one molecule attracted to δ- end of another
• Dispersion forces
– Caused by movements of e-
– Hydrogen Bonds
• Hydrogen bonded with very electronegative element is also
weakly bonded to an unshared pair on another
electronegative atom
Chapter 8
Characteristics of Ionic and
Covalent Compounds Characteristic Ionic Compound Covalent Compound
Representative Unit Formula Unit Molecule
Bond Formation Transfer of e- Sharing of e-
Type of elements Metal + Nonmetal Nonmetals
Physical State Solid Solid, liquid, or gas
Melting Point High (usually above
300°C)
Low (usually below
300°C)
Solubility in water High High to low
Electrical conductivity
of aqueous solution
Good conductor Poor to
nonconducting
Table 8.4, Pg 244, Text
Chemistry
Molecular Geometries
Adapted from
John D. Bookstaver
St. Charles Community College
St. Peters, MO
2006, Prentice-Hall, Inc.
Molecular Shapes
• The shape of a
molecule plays an
important role in its
reactivity.
• By noting the number
of bonding and
nonbonding electron
pairs we can easily
predict the shape of
the molecule.
What Determines the Shape of a Molecule?
• Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.
• By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.
Valence Shell Electron Pair Repulsion
Theory (VSEPR)
“The best
arrangement of a
given number of
electron domains is
the one that
minimizes the
repulsions among
them.”
Molecular Geometries
• The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.
Linear Electron Domain
• In this domain, there is only one molecular
geometry: linear.
• NOTE: If there are only two atoms in the
molecule, the molecule will be linear no
matter what the electron domain is.
Trigonal Planar Electron Domain
• There are two molecular geometries:
– Trigonal planar, if all the electron domains are
bonding
– Bent, if one of the domains is a nonbonding pair.
Polarity
• In Chapter 8 we
discussed bond dipoles.
• But just because a
molecule possesses
polar bonds does not
mean the molecule as a
whole will be polar.
Polarity
By adding the
individual bond
dipoles, one can
determine the
overall dipole
moment for the
molecule.
Overlap and Bonding
• We think of covalent bonds forming through the
sharing of electrons by adjacent atoms.
• In such an approach this can only occur when
orbitals on the two atoms overlap.