Chapters 8 & 9
IONIC and COVALENT
BONDING
Properties of Ionic Compounds
• Ionic compounds are composed of positive
ions surrounded by negative ions.
• Negative ions are surrounded by positive
ions.
• The resulting structure is called a crystal
lattice which is a regular, repeating, 3-D
arrangement of ions.
• It involves the strong attraction between
oppositely charged ions.
Unit Cells in a Crystal Lattice
Properties of Ionic Compounds
• HIGH melting and boiling points
• Brittleness (shatters)
• Nonconductors of electricity in the solid
state
• Good conductors when melted or
dissolved in water (called electrolytes)
Crystal Lattice
Lattice Energy
• Defined as the energy needed to separate
one mole of the ions of an ionic
compound.
• Expressed as a negative quantity.
• The more negative the lattice energy, the
stronger the force of attraction.
Factors affecting lattice energy
• The smaller the size of the ion, the
GREATER the energy and force of
attraction. Smaller ions are closer to the
nucleus and have a greater attraction for
the valence electrons.
• The higher the charge on the ion, the
greater the energy. Mg2+ is stronger than
K+
Practice
• For each of the following pairs of ionic
compounds, state which would be
expected the have the HIGHER (more
negative) lattice energy:
• LiF or KBr
• NaCl or MgS
• BaCl2 or KI
Metallic Bonds
• Bonding in metals is explained by the
electron sea model (atoms in a metal
contribute their valence electrons to form a
“sea” of electrons that surrounds metallic
cations).
• These electrons are delocalized and can
move easily throughout the solid– very
mobile.
Properties of Metals
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Very high boiling points
Melting points are much lower
Malleable (can be pounded into a shape)
Ductile (drawn into a wire)
Good conductors of heat and electricity
Lustrous; shiny
All of these properties are attributed to the
delocalized sea of electrons.
Electron Sea Model
Chemical Bonds
• Two major types of chemical bonds: Ionic
and Covalent
• Ionic bonds are formed from the transfer of
electrons (which forms ions).
• Produces a strong, electrostatic attraction
which accounts for the fact that ionic
compounds are hard and brittle with very
high melting points.
Ionic Bonding
Ionic Compound
Covalent Bonding
• Formed from the sharing of electrons.
• Most common type of bond.
• Sharing between electron pairs may be
equal or unequal.
Covalent Bonding
Nonpolar Covalent Bonding
• Aka “pure covalent”
• Equal sharing of electrons
• Caused when two atoms have the same
electronegativity and the same attraction
for electrons.
• This causes the molecule to be weaker,
with a LOWER boiling point compared to
stronger molecules.
Polar Covalent Bonding
• Caused by unequal sharing of electrons
• More common, and creates a stronger
force of attraction.
• Polar bonds have a partial positive charge
on one end of the molecule, and a partial
negative charge on the other end.
• δ+ is the abbreviation for partial + charge
• δ- means partial – charge
Cont.
• The more electronegative atom is
assigned the δ- charge and the less EN
atom is δ+.
• The greater the difference in EN (when
you subtract), the STRONGER the force of
attraction (and the higher the boiling point
or melting point).
Water Molecule Example
Example Cont.
• What kind of bond exists between H and
O?_Covalent___
• This means that electrons are _Shared_____
• What do both the arrows and the delta symbols
indicate?_Polarity___
• Which direction is the arrow pointing ( to the less
EN atom or the more EN atom)?
• The “plus” at the end of the arrow indicates the
(less EN atom or the more EN atom).
How to Determine Type of Bond
Use your electronegativity chart
(developed by Linus Pauling).
Electronegativity is defined as the power
of an atom to attract _electrons__ to itself
when bonding.
This is a numerical scale (which element
has the highest EN?).
It is useful in helping to determine if a
bond is polar, nonpolar or ionic.
Continued
In a compound, find the EN difference
between the 2 atoms (the absolute
difference). Don’t worry about the
subscripts. Look up the EN for EACH
atom, and SUBTRACT! Remember, only
the absolute difference is needed.
Guidelines for Bonds
If the difference in EN is
 ≥ 2.0 = ionic
 < 1.7 = covalent
 ≥0.3 = polar covalent
 < 0.3 = nonpolar covalent
If the difference is between 1.7 and 2.0: if a
METAL is involved, it is IONIC; if only nonmetals
are involved, it is polar covalent.
Why Bonds Form
• Two atoms will bond if it lowers their
potential energy.
• Nature seeks a decrease in energy
(downhill).
• The formation of a chemical bond is
usually an EXOTHERMIC process.
• This means the release of energy.
• Breaking a bond usually ABSORBS
energy (ENDOTHERMIC).
Cont.
• When atoms are separated by breaking a
bond, they have a HIGHER potential
energy than when they were bonded.
• Whether or not a chemical reaction will
occur spontaneously depends on whether
or not forming new bonds in the products
produces enough energy to break bonds
in the reactants.
Things to Memorize
• Molecular substances (molecules) have
COVALENT bonds.
• There are 7 diatomic molecules (they exist
as pairs). Since they are identical atoms,
the bonding is NONPOLAR covalent for
each.
• H2, N2, O2, F2, Cl2, Br2, I2
Polyatomic Ions
• A group of charged atoms that are
COVALENTLY bonded (since they are all
nonmetals).
• Examples include hydroxide (OH-),
cyanide (CN-), ammonium (NH4+) and
sulfate (SO42-)
Rules of Thumb
• In general, compounds composed of a
metal & a nonmetal are IONIC.
• Compounds are COVALENT if they are
nonmetal & nonmetal.
• Can you look at the location of the element
on the periodic chart and tell if it is a metal
or nonmetal?
Classify each as ionic or covalent
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CaCl2
CO2
H2O
K2O
NaF
CH4
SO3
LiBr
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MgO
HCl
KI
NO2
CO
FeCl3
P2O5
N2O3