Chapter 9: Covalent Bonding
Review
• Noble gases are the most stable
– Have full outer energy level
– Do not react with other elements to form bond
• Metals and Nonmetals react
– form binary ionic compounds
– electrons are transferred
– attain a noble gas configuration
• Sometimes two atoms that BOTH need to
gain electrons to become stable will have
a similar attraction for electrons
• These atoms SHARE electrons
• RECALL: Octet rule states that atoms will
gain, lose, or share electrons in order to
become stable
What is a covalent bond?
• Covalent bond- results from the sharing of
valence electrons
• The shared electrons are considered to
be part of the complete outer energy level
• Forms between two non-metals
Molecules
• Forms when two or more atoms bond
covalently
Ex. Carbohydrates, proteins, DNA, and fats
Formation of a Covalent Bond
• Hydrogen (H2), Nitrogen (N2), Oxygen (O2),
Fluorine (F2), Chlorine (Cl2), Bromine (Br2),
and Iodine (I2) are all diatomic molecules
• Diatomic molecules form because they are
more stable than the individual atoms
• Consider two individual atoms of Fluorine...
• Fluorine’s electron configuration is as follows:
1s22s22p5
• Each fluorine atom has 7 valence electrons
• As two atoms of Fluorine approach each other,
a repulsive force occurs
• The electrons repulse each other but there is
also an attraction between the nuclei of each
atom and the electrons
• A maximum point of attraction occurs and the
force of attraction balance out the forces of
repulsion (most stable)
• At this point, the covalent bond forms
Covalent Vocabulary
• Lone pairs-the unshared pair of electrons
• Bonding pair- the shared pair of electrons
Ex. Look at the dot structure of F2
Single Covalent Bonds
• Formed by the sharing of one pair of
electrons
• Hydrogen and the halogen group elements
will form single covalent bonds with an
identical atom
• Group 6A elements will form two single
covalent bonds (ex. H2O, Water)
• Group 5A will form three single covalent
bonds (ex. NH3, Ammonia)
Page 244, Practice Problems
Draw the Dot Structures for these molecules:
1. PH3
2. H2S
3. HCl
4. CCl4
5. SiH4
Sigma Bonds
• Sigma bonds are single covalent bonds
• Represented by the Greek letter sigma (σ)
• Occurs when electron pair is shared in an
area centered between two atoms and
there is an overlap in the valence orbital
Multiple Covalent Bonds
• Double bonds- two pairs of electrons are shared
(ex. O2)
• Triple bonds – three pairs of electrons are
shared (ex. N2)
• Multiple bonds consist of a sigma bond and at
least one pi bond (π)
• Double bond has one sigma and one pi bond
• Triple bond has one sigma and two pi bonds
Strength of Covalent Bonds
• Strength depends on how much distance
separates the nuclei
• Bond Dissociation energy
•
•
•
•
•
•
•
Extra Credit Assignment – Due Dec. 1
READ PAGES 246-247
DO page 247 6-11 (6)
READ PAGES 252-258
DO page 258 42-47(6)
READ PAGES 259 – 262
DO page 259 54-58 (5)
9.2 Naming Binary Molecules
1. First element is always named first, using
entire elements name
2. The second element is named with the
ending –ide
3. Prefixes are added to each name to
indicate the number of atoms present
4. Exception: if there is only one of the first
element omit prefix
Number of Atoms
Prefix
1
Mono-
2
Di-
3
Tri-
4
Tetra-
5
Penta-
6
Hexa-
7
Hepta-
8
Octa-
9
Nona-
10
Deca-
Page 249, 13-17
13. CCl4
14. As2O3
15.CO
16.SO2
17.NF3
Page 249, Table 9-2
Formulas and Names of
Some Covalent Compounds
Formulas
Common Name
Molecular Compound name
H 2O
Water
Dihydrogen monoxide
NH3
Ammonia
Nitrogen trihydride
N2H4
Hydrazine
Dinitrogen tetrahydride
N 2O
Nitrous oxide
Dinitrogen monoxide
NO
Nitric oxide
Nitrogen monoxide
Naming Acids
• Water solutions of some molecules are
acidic (ACIDS)
• If in solution, the compound produces H+
ions, then it is acidic
• Example: HCl
hydrochloric acid
• Two Types of Acids:
– 1. Binary Acids
– 2. Oxyacids
Binary Acids
• Contains hydrogen and one other element
• When naming, use prefix hydro- to
indicate the presence of hydrogen
• Add the ending –ic, and then write the
word acid at the end of the name.
• Ex. HBr
hydrobromic acid
Oxyacids
• Acids containing an oxyanion such as SO4
• In naming, do not use the hydro prefix, you
will need to change the suffix on the anion
– If the ending is –ate change it to –ic
– If the ending is –ite change it to –ous
HNO3  nitric acid
HNO2  nitrous acid
Page 250, 18-22
18. HI
19. HClO3
20. HClO2
21. H2SO3
22. H2S
9.3 Molecular Structures
• Structural formulas use the letter symbols and
the dots to illustrate relative positions of atoms
• The central atom is the one with the least
attractive force
• Hydrogen is always terminal atom due to only
sharing one possible electron
Ex. NH3
So, How do polyatomic ions
form anyway?
• Example:
ClO4-1
PO4-3
NH4+1
Resonance Structures
• It is possible to have more than one
correct Lewis Structure
• These are called resonance structures
• Ex. Nitrite ion NO2-1
• These are common molecules with more
than one structure: ozone (O3), nitrate,
nitrite, sulfite, and carbonate
9.4 Molecular Shape
• Molecular geometry is important ot
understanidng the stregth of bonds
• VSEPR Model determines shape
• Valence Shell Electron Pair Repulsion
– Based on minimizing repulsion of shared and
unshared pairs of electrons
– Memorize shapes on page 260
9.5 Electronegativity and Polarity
• Polar bonds occur when electrons are not
shared equally, resulting in unequal
distribution of charge and the formation of
a dipole
• spatial arrangement of polar bonds in a
molecule determines the overall polarity of
a molecule