Chemical Bonding
Chapter 7
Chemical bonding
• ionic bond: an electrostatic attraction
between ions of opposite charge
• covalent bond: “sharing” electrons
between atoms
• metallic bond: bonding occurs throughout
the substance, electrons flow freely
throughout the metal
– (metals make good conductors)
Ionic bonding
• ions group together in a lattice formation.
There is no “single” molecule of an ionic
compound.
• What kind of energy changes are involved
in formation of an ionic compound?
• Na (g) + Cl (g)  NaCl (g)
What trends are there in lattice
energies?
• the potential energy
between two particles.
– Q=charge on particle 1 or 2,
d=distance between the
particles, k=constant.
• So, if magnitude of Q1
and Q2 increases, the
energy _______.
• If d increases, the energy
______, but this change
is not as much as when
the charges are changed.
Q1Q2
Eel  k
d
Crystal lattice energies
• Determine the order of the crystal lattice
energies for the following ionic
compounds, from lowest to highest
• MgO, KBr, ScN
• KBr (671 kJ/mol) < MgO (3795 kJ/mol) <
ScN (7547 kJ/mol)
Lewis symbols
• a pictoral representation of the valence electrons of an atom
– valence electrons vs. core electrons
• examples:
Ionic compounds
• In binary ionic compounds (i.e., no polyatomic
ions involved), the atom to be the anion is
formed by “swiping” the electrons of the atom to
be the cation.
• Example: NaCl
Na
.
More examples:

:
..
Cl
.
:


Na   :

..
Cl
..

: 


Covalent bonds
• Instead of losing or gaining electrons,
atoms “share” electrons so that a bond is
formed between them.
Lewis formulas in covalent bonding
• Lewis formulas show how atoms “share”
electrons between each other to form a
bond.
• Examples: H2O, CH4, NH3
Octet rule
• Atoms tend to gain, lose, or share
electrons until they are surrounded by
EIGHT valence electrons.
• Why eight?
– The octet rule is not a law, and there are
several exceptions we will discuss later.
Rules for drawing Lewis structures
1.
2.
Sum all valence electrons from all atoms
Arrange the atoms
–
–
–
3.
complete octets around all atoms bonded to central atom
–
4.
5.
•
linear molecules (normally two or three), formula sometimes shows
atoms from right to left
central-grouped atoms; center atom normally written first.
Least electronegative element in the center
remember H has only 2 electrons
place all leftover electrons on central atom (even if more than an
octet results)
try multiple bonds if the central atom doesn’t have an octet using
single bonds
Examples: H2O, CO, NH3
Formal charge
• A bookkeeping tool used to help determine the best
structure between two or more structures that follow the
octet rule. (NOT the same as oxidation numbers)
• valence electrons =
• assigned electrons = all the unshared electrons
(nonbonding electrons) + ½ all the bonding electrons
• FC = Valence electrons – assigned electrons
• Sum of FC of all atoms must = charge.
• Rule: The structure resulting in lowest formal charge on
each atom “wins”.
• Example: structures of N2O
Resonance structures
• structures meeting both octet rule and
formal charge requirements.
• Resonant structures are equivalent, but
they are also equally wrong.
• The “actual” structure is in between all
resonance structures.
• Examples: O3, NO3-, SO3
Exceptions to the octet rule
1. Odd number of electrons in molecule.
2. Molecules where an atom has less than
an octet.
3. Molecules where atom has more than an
octet.
• Exception 1 examples: NO, ClO2
• Exception 2 examples: BeCl2, BF3
Exception 3 examples (most
common):
• PCl5, SF6, PO43• Central atom expands into its d shell
orbitals.
• For this to occur, central atom must be 3rd
row or higher.
Polar and non-polar bonds
• Not all atoms share electrons equally in
the atom.
• Electronegativity is an indicator of how
much one atom will “dominate” the
possession of electrons.
– What is electronegativity?
• The difference in electronegativities (EN)
lets us know how polar a molecule is…
Dipole moments
• Dipole moments occur in polar molecules.
• The larger the EN, the larger the dipole
moment.
• The dipole moment points toward the more
electronegative element.
• Molecules with more than two atoms are more
complicated (we’ll talk about that in Chapter 8).
• Examples: