Chapter 9
Covalent Bonding
• This chapter is hard
• You must do your homework and
study every day
• You must know your polyatomics
and be able to write chemical
formulas
• You must learn what I tell you to
learn
Covalent Bonding
• In ionic bonding, we talked about the
transferring of electrons. Covalent bonding
involves the sharing of pairs of electrons.
• Atoms that bond covalently are called
molecules.
• Covalent bonds occur nonmetal to nonmetal.
It can involve a metalloid.
• In covalent bonding all atoms involved
contribute.
Covalent Bonding and the Periodic
Table
• Because of the number of valence e- in each
group, nonmetals in certain groups will form
a certain number of bonds.
• Group 17 will share a pair of e-.
– Ex. F2
• Group 16 will share 2 pair of e-.
– Ex. O2
• Group 15 will share 3 pair of e-.
– Ex. NH3
Types of Bonds
• Single bond
–Also called a sigma bond(
)
–Sharing of a single pair of electrons
–In a sigma bond, the s orbital holding
the shared e- will be centered
between two atoms and overlap end
to end.
• Double bond
–2 shared pair of e–1 sigma bond and 1 pi bond
• Pi bond (π)
–The orbitals containing the
shared electrons overlap above
and below the plane of the sigma
bond.
–Pi bonds cannot be alone; They
must accompany a sigma bond.
•Triple bond
–3 shared pair of e–Contains 2 pi
bonds and one
sigma bond.
•Bonding examples
–Draw and model
•C2H4
•C2H2
Bond Length and Strength
• The shorter the length the stronger
the bond.
• Triple bonds are the shortest and
have the greatest strength.
• Double is smaller than single.
• Single is the longest and weakest.
Bond Association Energy
• The total energy of a chemical reaction is
determined by bonds formed and bonds
broken.
• Determining type of reaction
– Endothermic vs Exothermic
• Formula: energy of bonds broken – energy of bonds
formed = net energy of products.
• If the value is positive it is an endothermic reaction.
• If the value is negative it is an exothermic reaction.
Naming Binary Molecular Compounds
• Use prefixes listed on page 248
–You must learn these
–Name root-ide
–Add prefixes
• Do not add mono- to the first word
• Do not use double vowels
–No i-i, o-o, a-a
Examples
•
•
•
•
•
CO2 – Carbon dioxide
CO – Carbon monoxide
CCl4 – Carbon tetrachloride
As2O3 – Diarsenic trioxide
Pg. 249
– H2O – Water
– NH3 - Ammonia
Acids and Bases
• Acids
–Arrhenius acid – H+
–Bronsted-Lowry acid – H+ donor
• Bases
–Arrhenius base – OH–Bronsted-Lowry base – H+ acceptor
Naming Acids
• Acids always contain H and they are in an
aqueous solution
– They are different as a gas
• There are 2 types
– Binary acids that do not contain oxygen
• H and a monoatomic anion in aqueous solution
• Exception
– Polyatomic ions without O
– Oxyacids
• H and an O
Naming Binary Acids
• Hydro- root of anion –ic acid
– Exception
• You always use the full name for sulfur
– HBr(aq) – hydrobromic acid
– HCl(aq) – hydrochloric acid
– H2S(aq) – hydrosulfuric acid
– Exception: Polyatomic without O
• HCN(aq) – hydrocyanic acid
Naming Oxyacids
• H and an oxyanion
• DO NOT USE HYDRO
• Polyatomic anions end in –ate or –ite. This
indicates the number of O atoms in the polyatomic.
You will use this for naming oxyacids.
• If the polyatomic ends in –ate you will name the
acid ending with –ic. If the polyatomic ends in
–ite you will name the acid ending with –ous.
• Example
– HNO3 – nitric acid
– HNO2 – nitrous acid
Acids that contain Halogens
• Polyatomics that contain a halogen have 4
names.
• Example: perchlorate, chlorate, chlorite,
hypochlorite.
• When naming these you will apply the same
rules of –ate to –ic and –ite to –ous AND indicate
per- and hypo-.
• Example
–
–
–
–
HClO4 – perchloric acid
HClO3 – chloric acid
HClO2 – chlorous acid
HClO – hypochlorous acid
Name These Acids
•
•
•
•
•
•
•
HIO2
HBr
HBrO3
H2SO3
H2CO3
HF
HIO4
•
•
•
•
•
•
•
Iodous acid
Hydrobromic acid
Bromic acid
Sulfurous acid
Carbonic acid
Hydrofluoric acid
Periodic acid
Naming Hydroxides
• Cation and hydroxide
–Most are metals; ammonium is not.
–Examples
• NaOH – Sodium hydroxide
• NH4OH – Ammonium hydroxide
• Mg(OH)2 – Magnesium hydroxide
•Groups
•Pg 250
–18-22
Naming Covalent Compounds Practice
• Worksheets
– Covalent compounds and Acids
• Pg 272 88-98
• Pg 274 128
Molecular Structures
• Fact you must know
– Lone or nonbonding pairs of electrons require more
space than bonding pairs.
• Rules for depicting structural formula
• Example: NH3
1) Predict atom location
– Terminal atoms – comes off of the central atom(s)
• H is always terminal
– Central atoms – usually located closest to the middle
of the periodic table. (Only nonmetals – We are
talking covalent bonding)
• Everything will stem off of it
2. Get the sum of all valence e– NH3
• 5e- + 3(1e-) = 8e• Divide this number by two in order to get number of
pairs of electrons
• NH3 has 4 pair
3. Use one pair to bond each terminal atom
to the central atom
H
N
H
H
4. Equally distribute remaining
pairs to the terminals. Then,
add any remaining pairs to
the central atom.
–Make sure each atom has 4
pair. Why?
–What about H?
H
..
N
H
H
5. If the central does
not have 4 pairs of
e- then make
multiple bonds.
•Example: CO2
Polyatomic Lewis Structures
• Follow the same steps given for
structural formula. The only difference
is you must indicate charge by using
[ ] or [ ]+.
• Examples
– ClO4– NH4+
Resonance Structures
• What is resonance?
– Guitar string
– Something vibrating back and forth very
quickly
– Resonance structures occur when there is
more than one Lewis structure for a
molecule or compound.
• Example: NO3- and SO2
• In NO3-, one bond resonates between 3
places.
• You must draw each resonating structure.
Exceptions to the Octet Rule
• Some atoms have odd numbers of
electrons
– Example: NO2 has 17 e-
• Boron tends to form 3 bonds when it is the
central atom.
– Example: BH3
• Expanded Octets
– Some atoms form more than 4 bonds on the
central atom.
• Almost always involves halogens
• PCl5
Positions to know when looking at
shape
Cl
P
Cl
Axial Position
Cl
P
Cl
Equatorial Position
• You must learn page 260.
• All of it
• Group Work
– Pg 255 #30-34
• Pg. 273
–99-104
• Naming Sheet
– No book
– No notes
Molecular Shape
• Linear
– Usually formed by atoms with multiple bonds
• Example: CO2
• Trigonal Planer
– Flat
– All angles are equal (120)
• Example: BH3
• Tetrahedral
– There are four equal bonding sites
– All angles are equal (109.5)
• Example: CH4
• The next four are deformation of
the tetrahedral.
• Trigonal Pyramidal
–Involves a lone pair. Lone pairs take
up more space than bonded pairs.
–The angles between the terminal
atoms decreases from the angle of the
tetrahedral (107.3).
–Example: NH3
• Bent
– 2 lone pairs
– The angle decreases more (104.5)
– Example: H2O
• Trigonal Bypyramidal
– What does bi mean
• 2 pyramids
– Only occurs with elements in period 3 or
greater
– Must have d orbitals
– Has 2 different angles: 90 and 120
• Octahedran
–All bond angles are equal
(90)
–Forms 6 bonds and 8 (octa-)
faces
–Example: SF6
Hybridization
• Hybrid orbitals
– Mixing of orbitals to form new orbitals
• Example: C
– Show with electron configuration and
orbital notation
– Carbon tends to form 4 bonds
– The outer most energy level is 2
– Forms an sp3 hybrid
Bond Polarity
• Bonds have oppositely charged ends
• Related to electronegativity in a chemical
bond
– The higher the electronegativity value the more
strongly the atom pulls electrons to itself in a
chemical bond.
• Pg. 263
• In a bond between F and H which atom will
pull more on an electron?
– F – 3.98 and H – 2.20
• Because F has a greater pull on the
electron it creates polarity.
• How to represent polarity.
–H
–
F or
Is the symbol for indicating a
partial charge
• Compare the electronegativity values of the
bonded atoms one bond at a time.
• Rules for determining polarity
– If the difference is less than or equal to 0.4 then the
bond is nonpolar.
– If the difference is between 0.5 and 1.7 then the
bond is polar.
– If the bond is greater than 1.7 then the bond is ionic.
• The only true covalent bond exist between
bonded identical atoms.
• In order for a molecule to be polar you must
have positive and negative ends.
• How does symmetry of a molecule relate to
polarity?
Polarity Example Problem
• H2O
Homework
• Pg 273 105-114