Chapters 7 and 8
Chemical bonding
Valence Electrons
 Elements with similar chemical behavior have the same
number of valence electrons.
 For the representative elements (1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A)
the group number corresponds to the number of valence
electron in each group (with the exception of He)
 When examining electron configurations, the electrons that
are present in the highest principle energy level represent the
valence electrons of those atoms.
Br: [Ar]4s2 3d10 4p5
Bromine has 7 valence electrons
Valence Electrons and Electron Dot
Structures
 Valence electrons are the electrons that participate in
chemical bonds
 Electron dot structures consist of the atom symbol and its
valence electrons represented as dots.
Br: [Ar]4s2 3d10 4p5
The Octet Rule - Metals
 Metals tend to lose electrons to gain stability. When they
lose electrons they attain a stable electron configuration
where the next lowest-energy level has a full octet.
The Octet Rule – Non-metals
 Non-metals tend to gain electrons to obtain a full octet in
their highest occupied p sublevel
Formation of Ionic Compounds
 Cations and anions combine to form ionic compounds
+
+
NaCl
The Formation of Calcium Chloride
CaCl2
Properties of Ionic Compounds
 The chemical formula for an ionic compound is called a formula
unit which represents the lowest who number ration of atoms in
that compound.( ex. NaCl)
 Most ionic compounds are crystalline solids at room temperature
 Ionic compounds generally have very high melting points
 Ionic compounds can conduct electricity when molten or
dissolved in water
Coordination Number = The
number of ions of opposite charge
surrounding the ion in a crystal.
Metallic Bonding
 Positive metal ions floating in a sea of electrons.
• Metals are excellent conductors
of electricity
• Metals can be easily deformed
(They are malleable and ductile)
Know that metals are arranged in crystalline
structures, but do not worry about the
arrangement of these structures: face-centered,
body-centered, hexagonal.
Alloys
 Alloys are mixtures of metals
 Properties of alloys are usually superior to those of their
component elements.
Substitutional Alloy
Interstitial Alloy
Covalent Bonding
 Covalent bonds occur between two or more non-metals
 Unlike ionic bonds where electrons are transferred from one
atom to another, electrons are shared between atoms in a
covalent bond.
 Atoms joined together by covalent bonds are called
molecules
 A compound composed of molecules is called a molecular
compound
Molecular and Structural Formulas
 A molecular formula indicates the types and numbers of each
atom in a molecule
 The structural formula indicate the arrangement of the atoms
in the molecule
H2O
Covalent Bonds and the Octet Rule
 Atoms share electrons in a covalent bond so that each atom
has enough electrons to satisfy the octet rule
Varieties of Covalent Bonds
 Single bonds (sigma bonds) en
 One pair of electrons is shared between two atoms
Lone pair
Varieties of Covalent Bonds
 Double Bonds (1sigma bond, 1 pi bond)
 Atoms share two pairs of electrons
 Triple Bonds (1sigma bond, 2 pi bonds)
 Atoms share three pairs of electrons
Coordinate Covalent Bonds
 A covalent bond in which one atom contributes both bonding
electrons.
Resonance Structures
Def: A condition when more than one valid Lewis structure
can be written for a molecule or ion.

Exceptions
to
the
Octet
Rule
Too few electrons surrounding the central atom (ex: BH )
3
Boron will not have a full octet, only 6 electrons. It can only achieve a full octet when
another atom shares an entire pair of electrons with it (Coordinate covalent bonding)
 Too many electrons surrounding the central atom (ex: PCl5)
 An odd number of electrons
How to Draw a Lewis Structure for Molecules

Predict the location of atoms
1.
If there are more than two atoms, place the least electronegative atom in the center and
surround it by the remaining atoms.
2.
Hydrogen is always terminal (outside) because it can only make one bond

Determine the total number of electrons if each atom had a full set of valence electrons (2
for H, 8 for all others)

Add up the number of valence electron that you have to work with

Subtract total valence electrons from total electrons and divide by two. This is the number
of bonding pairs that are needed to put together the molecule.

Connect the atoms with the number of bonds that you calculated above

Add lone pairs where needed so that each atom has a full octet (except for hydrogen which
can only have two electrons)
Molecule
HCN
Total Electrons
Valence Electrons Bonding Pairs
Polyatomic Ions
Polyatomic ions are a cluster of non-metals that carry a charge.
To draw the structure of a polyatomic ion, follow the procedure for
drawing ordinary molecules but add or subtract the number of
electrons gained or lost to the total number of valence electrons in
your structure as indicated by the charge on the ion.
Molecule
IO3-
Total Electrons
Valence Electrons
Bonding Pairs
Molecular Shape (VSEPR)
Valence Shell Electron Pair Repulsion – minimizes the repulsion of shared and
unshared pairs of electrons around the central atom.
 The shape of a molecule determines many of its physical and chemical
properties.
 The VSEPR is based on the arrangement of bonding and lone electrons
around a central atom to minimize repulsion.
 The repulsion of electrons creates a specific bond angle between a central
atom and two terminal atoms.
 Lone pairs of electrons occupy more space than bonding pairs of electrons
Molecular Geometry
or 2
Electronegativity and Polarity
Recall: Electronegativity is the ability of an atom to attract an electron.
Chemical bonding is like “Tug-o-War”
Bond Type
Non-polar Covalent
Polar Covalent
Ionic
Electronegativity
Difference
0-0.4
0.5-2.0
>2.0
Molecular Polarity
Molecules are either polar or non-polar
Both polar and non-polar molecules may contain polar bonds. What determines whether
a molecule is polar or non-polar is the symmetry of the molecule
PolarBonds
Present
Symmetr Polar/
y
NonPolar
Examples
No
No
Non-Polar
NO2
No
Yes
Non-polar
SiH4
Yes
No
Polar
NH3
Yes
Yes
Non-polar
CO2
VSEPR shapes that can demonstrate symmetry are:
Linear
Trigonal Planar
Tetrahedral