Bonding
By John Patrick Fahy III of Galway
Coulomb’s Law
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Attractive force is proportional to (+q)(-q)/r^2
+q = magnitude of the positive charge
-q = magnitude of the negative charge
r = distance between the charges
All bonds occur because of electrostatic
attractions, electrostatic forces are governed
by this law.
What does this mean?
• Bigger charges mean stronger bonds, smaller
charges mean weaker bonds
• Charges close together means stronger bonds,
charges far apart mean weaker bonds
Bonds Within Molecule
• Atoms join to form molecule because atoms
like to have a full outer shell of electrons. This
usually means having eight electrons in the
outer shell. Atoms with too many or too few
electrons either give up, gain, or share
electrons with other atoms.
Types of Bonding
Ionic bonds
• Occur between atoms with
very different
electronegativities; this
means usually metals and
nonmetals
• one atom gives up electrons
and becomes a cation while
the other the accepts
electrons and becomes an
anion
Covalent Bonds
• Occurs when 2 atoms share
electrons
• Each atom counts the
shared electrons as part of
their valence shell
Drawing Lewis Structures
• 1. Count the valence electrons in the molecule
• 2. If it is an ion either subtract or add
electrons depending on its charge
• 3. Draw the skeletal structure of the molecule
and place one single bond between each pair
of bonded atoms, the least electronegative
atom usually serves as the central atom
• 4. Add electrons to the surrounding atoms
until each has a complete outer shell
• 5. Add the remaining electrons to the central
atom
• 6. Look at the atom
• (a) if the central atom has less than eight
electrons remove an electron pair from an
outer atom and add another bond between
that and the central atom. Do this until the
central atom has a complete octet
• (b) if the central atom has a complete octet,
you’re finished
• (c) if the central atom has more than 8, that’s
okay too
Molecular geometry
• When atoms come together to form a
molecule, the molecule will assume the shape
that keeps its different electron pairs as far
apart as possible, the model based on this
idea to predict molecular geometries is known
as the valence shell electron-pair repulsion
theory (VSEPR)
Geometry
• If the central atom has 2 electron pairs than
its hybridization is sp and its shape is linear
• Ex. BeCl2, CO2
• If the central atom has three electron pairs
then it has sp2 hybridization and its shape is
trigonal planar, if one of the pair is a lone pair
then the shape is categorized as bent
• Ex. BF3, SO3, SO2
• If the central atom has 4 electron pairs then it
has sp3 hybridization and its shape is
tetrahedral, if one of the pairs is a lone pair
then its shape is trigonal planar, if 2 of the
pairs are lone pairs then it shape is bent
• Ex. CH4, NH3, P
• Cl3, AsH3, H2O, OF2
• If the central atom has 5 electron pairs then it
has dsp3 hybridization and it is trigonal
bipyramidal, 1 lone pair = seesaw, 2 lone pairs
= T-Shaped, 3 lone pairs = linear
• Ex. PCl5, PF5, SF4, ClF3, ICl3, XeF2
• If the central atom has 6 electron pairs then it
has d2sp3 hybridization and is octahedral
• 1 lone pair = square pyramidal
• 2 lone pairs = square planar
• Ex. SF6, BrF5, IF5, XeF4
Intermolecular Forces
Dipole-Dipole Forces
• Occur between neutral
polar molecules
• The positive end of one
molecule is attracted to the
negative end of another
molecule
• Molecules with greater
polarity have greater dipoledipole attractions
London Dispersion Forces
• Occur between neutral,
non-polar molecules
• Very weak, occur because of
the random motions of
electrons on atoms within
molecules