Covalent Bonding
Chapter 9
Honors Chemistry
Glencoe
Section 9-1: The Covalent Bond
I.
Atoms bond so they
can achieve a stable
valence e- configuration
(8e-, except for H & He,
which need only 2e-)
II. Lewis Dot Structures:
element surrounded by
dots indicating the
number of valence
electrons. Examples.
III. Covalent Bond: shared e- form a
bond between nonmetals to form a
molecule
IV. Electronegativity and Covalent
Bonding: Sharing: Equally or
Unequally? (Oreos)
A. Nonpolar covalent
bond: electrons
are equally shared
(equal or almost
the same
electronegativity or
pull). Ex: Cl2, O2.
B. Polar covalent bond: electrons are
shared unequally
(electronegativities are different)
Section 9-5
1. Dipoles – poles with + and –
charges are formed when the e- are
pulled. Designated with “δ”.
2. The side with the stronger pull will
have a - dipole. The weaker pull will
have a + dipole.
IV. Electronegativity and Bond Types –
find the difference (subtract) the
electronegativities
A. Nonpolar Covalent bond: if the
difference is below 0.5
B. Polar Covalent bond: if the
difference is between 0.5 and
1.7
C. Ionic bond: if the difference is
greater than 1.7
Section 9-2: Naming Molecular
Compounds and Acids
I.
Naming and Writing Formulas for
Covalent Compounds
A. Naming covalent (molecular
compounds)
1. Name the element in the same
order that they appear in the
formula.
2. Change the last syllable of the
last element to –ide.
3. Add prefixes to each element’s name
to indicate how many atoms of that
element are present in the formula.
4. Mono- is optional, especially for the
first element in the formula.
5. Ex: CO2, CO, BF3, NI7, P4O5, P8 Cl9,
N6O, CaCl2
B. Writing formulas of covalent (molecular)
compounds (nonmetals with nonmetals)
1. Write the symbol of the elements in
the order they appear.
2. Give them the subscript as indicated
by prefix.
3. Examples: carbon monoxide,
dicarbon tetrafluoride, triphosphorus
hexasulfide, nonacarbon heptaiodide,
copper oxide
II.
Acid Names and Formulas
A. Acids are combinations of the hydrogen ion
(H+) and an anion (negatively charged ion).
Examples:
H+ and SO42- form H2SO4
H+ and Cl- form HCl
H+ and PO43- form H3PO4
B. Naming of acids are based on the endings
of the anion. There are 3 endings: -ide, -ite
and -ate.
1. Anions with –ide ending are named
hydro___ic acid.
Examples: H+ and Cl- form HCl
hydrochloric acid
H+ and Br- form HBr
hydrobromic acid
2. Anions with –ite ending are named
___ous acid.
Examples: H+ and ClO2- form HClO2
chlorous acid
H+ and PO33- form H3PO3
phosphorous acid
3. Anions with –ate ending are named ___ic
acid.
Examples: H+ and NO3- form HNO3
nitric acid
H+ and C2H3O2- form HC2H3O2
acetic acid
H+ and CO32- form H2CO3
carbonic acid
Section 9-3: Molecular Structures
I.
Lewis Dot Structures
A. Dots (or crosses) represent valence
e-.
B. Lewis Structures represent
covalently bonded molecules.
1. Bonded pairs: pairs of e- that are
shared are shown as a line
2. Unshared or lone pairs – e- that are not
shared are shown as a pair of dots
C. Drawing Lewis Structures for Molecules
1. Count the total number of valence e-.
2. Arrange the atoms. Carbon is
usually in the center. Hydrogens and
halogens are normally on the ends.
3. All elements have 8 e-, except for H
and He, which have 2 e-. Share
electrons to accomplish this. Place 8
around central atom to start. Hint: Edo NOT need to stay with initial
element!
4. Draw lines to represent the shared
bonds.
5. Check your drawing: Is the number of
e- correct? Is everyone happy?
D. Examples:
1. HCl
2. PH3
3. CF4
4. H2O
E. Structures for ions: extra or missing
electrons are accounted for; brackets with
a charge is used. Ex: NH4+
II.
Multiple Bonds: sometimes more than
one pair of electrons are shared resulting
in double and triple bonds. Ex: HCN,
C2H4
III. Limitations of the Octet Rule
A. Atoms with less than an octet.
1. Central atoms can be stable
with less than an octet. (Usually
boron is the central atom.)
2. Example: BF3
B. Atoms with more than an octet.
1. Noble gases can have more
than an octet on the central
atom.
2. Examples: XeF4 and SF6
IV. Equivalent dot structures
A. Resonance structures: different dot
structures without changing the
arrangement of the atoms. Only epositions change.
B. Resonance is represented by
double-ended arrows (↔).
C. Example: Carbonate ion, CO32-
Section 9-4: Molecular Shape
I.
II.
VSEPR (Valence Shell E- Pair
Repulsion)
A. E- will determine the shape of the
molecule.
B. The e- pairs will orient themselves so
that they are as far away from each
other as possible.
Shapes: See Handout
Molecular Shapes
1. Linear – atoms can be connected in
a straight line. The central atom
has no unshared pairs of electrons.
Molecules with only 2 atoms will
always be linear. Bond angle =
180°. Examples: O2, HCl, CO2.
2. Trigonal Planar – central atom is
bonded to 3 other atoms with no
unshared pairs of e-. Does not
follow the Octet Rule. Boron is
typically the central atom. Bond
angle = 120°. Example: BCl3.
3. Tetrahedral – central atom is
bonded to 4 other atoms. The
valence e- are arranged equally
in 3-d. Bond angle = 109.5°.
Examples: CH4, CF4.
4. Pyramidal – central
atom is bonded with
3 other atoms, but
also has an unshared
pair of valence e-.
Bond angle = 107°.
Examples: NH3, PCl3.
5. Bent – central atom is bonded
to 2 other atoms and has 2
unshared pairs. Bond angle =
105°. Example: H2O.
6. Trigonal bipyramidal – 5 atoms
attached to the central atoms
which has no lone pairs. The
bond angles are 90° and 120°.
Example: PCl5.
7. Octahedral – 6 atoms around the central
atom, which has no unshared pairs. The
bond angles are 90°. Example: SF6.
Examples:
1. What is the shape of the molecule PI3?
2. What is the shape of HCN?
What is the shape of
the molecule to the
right?
Section 9-5: Polarity
I.
Molecular Polarity
A. Polar molecules – net pull in one
direction; polarity is indicated with
arrows. Be sure to use the Ball and
Stick models! Example: HF, CHCl3 &
H2O.
B. Non-polar molecules – overall
nonpolar because all the pulls cancels
each other out, but may still have
individual polar bonds. Example: BF3
and CH4.
II.
Solubility
A. “Like will dissolve like”: polar (ionic)
will dissolve in polar and nonpolar
will dissolve in nonpolar.
1. Water (polar) will not mix with oil
(nonpolar)
2. Salt (ionic) will dissolve in water
(polar), but salt will not dissolve in
oil (nonpolar)…unless the
temperature is changed!
III. Covalent Network Solid
A. Strong covalent bonds,
therefore these solids have
very high temperatures.
B. Examples: quartz,
diamonds, asbestos,
graphite