3.06B and 3.07
Concepts
3.06 B Lewis Structures
• Just as we represented valence electrons
around a symbol of an element in a Lewis
structure of an individual atom, we can also
represent bonded and non-bonded valence
electrons around the atoms within a
molecule.
• Ex:
3.06 B Lewis Structures
• In a Lewis dot, or electron dot, structure of
a covalent compound, chemists usually use
a straight line to represent the two
electrons shared in a covalent bond.
3.06 B Lewis Structures
• The other valence electrons not involved in
bonding are represented by dots around
the symbol of the element. These other
valence electrons are called unshared, or
lone pair, electrons.
3.06 B Steps for drawing Lewis Stuctures
• Place the least electronegative element in
the center of the molecule.
– The central atom is usually the first element
written in the molecular formula, except H,
which cannot be a central atom because it only
bonds once and then its valence is “full.”
3.06 B Steps for drawing Lewis Stuctures
• Calculate the total number of valence
electrons.
– Add up the total number of valence electrons
thatshould be in the picture according to the
periodic table. Make a list of the number of
valence electrons (determined by the location
on the periodic table) for each atom in the
molecule, and then add them together. We will
call this the “reality number” because this is
the number of valence electrons that must be
present in the molecule.
3.06 B Steps for drawing Lewis Stuctures
• Write the skeleton structure.
– Attach all other atoms to the center with single
bonds. You may change this later, but we know
that each atom in the molecule must be
attached with at least one set of shared
electrons (a single covalent bond).
3.06 B Steps for drawing Lewis Stuctures
• Complete the valence electrons.
– Fill every atom’s valence by adding electron
pairs until they are all full. Remember that
most atoms need eight valence electrons to be
full, except H, which is full with two electrons.
Also, remember that a single bond (each line
drawn in the model) represents two electrons
being shared by both atoms involved in the
bond. This means that one line counts as two
electrons in the valence of each atom touching
that line.
3.06 B Steps for drawing Lewis Stuctures
• Tally the totals.
– Count up the number of electrons represented
in your drawing. Remember that each line
represents two shared electrons and each pair
of dots represents two unshared electrons
when you are counting the electrons in the
drawing. We will call this counted number the
“picture number” in our examples.
3.06 B Steps for drawing Lewis Stuctures
• Perform a comparison.
– Compare the picture number to the reality
number to see if you need to change your
Lewis structure drawing.
– If picture # = Reality #, then the Lewis structure
drawing is a good representation of the
molecule and you do not need to change
anything.
– If picture # > Reality #, you must fix it by
adding multiple bonds where appropriate in
order to reduce the Picture # to equal Reality
#.
3.06 B Steps for drawing Lewis Stuctures
• Distribute.
– Distribute electrons to atoms surrounding the
central atom to satisfy octet rule
– Atoms that form multiple bonds are C, N, O, S.
Oxygen atoms do not bond to each other
(except in O2 & O3, H2O2, peroxides,
superoxides).
– All atoms must have eight electrons (octet
rule), except hydrogen whose octet is two.
3.06 B Practice
• Share your desktop and show the 3.06
Lewis Structures examples and practice.
3.06 B What to do!
• Share your desktop and show the 3.06B Lab
and where to complete it and submit it.
3.07 Intermolecular Forces
• There are also forces of attraction between
separate molecules, called intermolecular
forces. The prefix inter- comes from the
Latin stem meaning “between,” such as in
words like Internet, interface, and
international. Intermolecular forces,
sometimes called van der Waals forces,
vary in strength, but they are generally
weaker than the ionic and covalent bonds
found within compounds.
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Molecular Geometry
3.07 Predicting Polarity
• Share your desktop and show the steps and
examples.
3.07 Electronegativity and polarity
• The difference in electronegativity values
will affect the polarity (dipole moment) of
the molecule.
3.07 Dipoles
• Many molecules have dipole
moments due to non-uniform distributions
of positive and negative charges on the
various atoms. Partial charges are denoted
as δ+ (delta plus) and δ− (delta minus).
• There are three types of dipoles:
– Permanent dipoles
– Instantaneous dipoles
– Induced dipoles
3.07 Dipoles
• Instantaneous dipoles: These occur due to
chance when electrons happen to be more
concentrated in one place than another in
a molecule, creating a temporary dipole. A
molecule is polarized when it carries an
instantaneous or an induced dipole.
3.07 Dipoles
• Induced dipoles: These can occur when
one molecule with a permanent dipole
repels another molecule's electrons,
"inducing" a dipole moment in that
molecule temporarily.
3.07 Intermolecular forces
• Intermolecular forces are attractive forces
between molecules. These forces exist
between molecules when they are
sufficiently close to each other. They are
responsible for the non-ideal behavior of
gases and for properties of matter such as
boiling point and melting point.
• There are four different types of
intermolecular forces, depending on the
polarity of the molecules involved.
3.07 Intermolecular forces
• In order of increasing strength these inter
molecular forces are:
–
–
–
–
London Dispersion
Dipole-Dipole
Hydrogen Bonding
Ion-Dipole
• Share you desktop and show the different
forces.
3.07 What to do!
• Share your desktop and show the 3.07 Lab
and where to complete it and submit it.