Chapter 6 Notes
Chemical Bonding
Section 1:Introduction to
Chemical Bonding

Atoms seldom exist as independent
particles in nature. Most substances
consist of combinations of atoms held
together by chemical bonds.
 Chemical Bond – A mutual electrical
attraction between the nuclei and valence
electrons of different atoms that binds the
atoms together.
Section 1:Introduction to
Chemical Bonding
 You might ask why most particles bond to
others? This is because as independent
particles most atoms have relatively high
potential energies. Nature, however,
favors arrangements in which potential
energy is minimized. This means that
most atoms are less stable existing by
themselves than when they are combined.
By bonding with each other, atoms
decrease their potential energy, thereby
creating more stable arrangements.
Section 1:Introduction to
Chemical Bonding
Types of Chemical Bonding
 When atoms bond, their valence electrons
are redistributed in a way that makes the
atoms more stable. The manner in which
the electrons are redistributed determines
what type of bond will be made.
Section 1:Introduction to
Chemical Bonding
 Ionic Bonding – Chemical bonding that
results from the electrical attraction
between cations and anions.
 In purely ionic bonds, an atom will give up
electrons to another atom in the
compound. In contrast, a covalent bond
is made when the atoms share electrons.
 Covalent Bonding – Chemical bonding that
results from the sharing of valence
electrons between atoms.
Section 1:Introduction to
Chemical Bonding
 Bonding between atoms is rarely purely
ionic or covalent. It usually falls
somewhere in the middle depending upon
how strongly the atoms are attracted to
the electrons of the other atoms in the
compound.
Section 1:Introduction to
Chemical Bonding
 Remember that electronegativity is a
measure of an atom’s ability to attract
electrons. The degree to which bonding
between atoms of two elements is ionic or
covalent can be estimated by calculating
the difference in the element’s
electronegativities.
Section 1:Introduction to
Chemical Bonding
 Ionic – any difference that is greater than
1.7
 Polar-Covalent – any difference that is
greater than or equal to .3 and less than
or equal to 1.7
 Non-polar covalent – any difference that is
less than .3
Section 1:Introduction to
Chemical Bonding
 Non-polar covalent bond – A covalent
bond in which the bonding electrons are
shared equally by the bonding atoms,
resulting in a balanced distribution of
electrical charge.
 Polar – An uneven distribution of charge.
 Polar-covalent bond – A covalent bond in
which the bonded atoms have an unequal
attraction for the shared electrons.
Section 1:Introduction to
Chemical Bonding
 Hydrogen is an element that is never
found alone. Even when it is just an
element, hydrogen is at least bonded to
itself as H2. In this sort of situation the
electronegativities between the two
elements in the compound would be 0.
What that means is that the electrons that
are shared between the two atoms would
be shared equally because the atoms are
pulling on the electrons with the same
force.
Section 1:Introduction to
Chemical Bonding
 However, in water, hydrogen is bonded to
a different atom. Oxygen has a
significantly higher electronegativity than
hydrogen does. Therefore, the oxygen
nucleus pulls harder on the electrons that
are being shared. The electrons are still
shared, just not equally.
Section 1:Introduction to
Chemical Bonding
 They are more likely to be found around
the oxygen nucleus than the hydrogen
nucleus because oxygen pulls harder on
them. That, in turn, makes the water
molecule polar. All polar means is that it
has a positive side and a negative side.
Because the electrons spend more time
toward the oxygen, that side is slightly
negative and the hydrogens are slightly
positive.
Section 2:Covalent Bonding and
Molecular Compounds
 Many chemical compounds found in living
things, and produced by living things, are
composed of molecules.
 Molecule – A neutral group of atoms that
are held together by covalent bonds.
 Molecular compound – A chemical
compound whose simplest units are
molecules.
Section 2:Covalent Bonding and
Molecular Compounds
 A single molecule of a chemical compound
is an individual unit capable of existing on
its own. It may consist of two or more
atoms of the same element, such as H2, or
of two or more different atoms as in CO2.
Section 2:Covalent Bonding and
Molecular Compounds
 The composition of a compound is given
by its chemical formula.
 Chemical formula – Indicates the relative
numbers of atoms of each kind in a
chemical compound by using atomic
symbols and numerical subscripts.
 We use the term chemical formula to
describe any compound. If we are
specifically talking about a covalent
compound, however, we would refer to it
as a molecular formula.
Section 2:Covalent Bonding and
Molecular Compounds
 Molecular formula – Shows the types and
numbers of atoms combined in a single
molecule of a molecular compound.
 The molecular formula for water, for
example, would be H2O. This reflects the
fact that there are two hydrogen atoms
that are covalently bonded to an oxygen
atom.
Section 2:Covalent Bonding and
Molecular Compounds
 Formation of a Covalent Bond
 So how does a covalent bond form? Well,
remember that we said that all atoms
have a certain energy associated with
them? When those atoms are bonded,
usually it lowers the amount of potential
energy that they have. Keep this in mind.
Section 2:Covalent Bonding and
Molecular Compounds
 Picture two isolated hydrogen atoms
separated by a distance large enough that
they do not influence each other. At this
distance, the overall potential energy of
them is arbitrarily set at 0. Now, consider
what must happen as those two atoms
approach each other. The electrons will
repel each other and the nuclei will too.
Section 2:Covalent Bonding and
Molecular Compounds
 However, the electron from one hydrogen
atom will be attracted to the other
hydrogen nucleus. As the get closer to
each other the potential energy must be
decreasing. At some point, the nuclei will
be at the optimum distance from each
other, shown by the lowest possible
potential energy for those two atoms.
Section 2:Covalent Bonding and
Molecular Compounds
 As you continue to push the atoms
together, the potential energy will increase
because the nuclei are too close to each
other and they want to repel. Like
pushing two positive ends of magnets
together.
Section 2:Covalent Bonding and
Molecular Compounds
Section 2:Covalent Bonding and
Molecular Compounds
 The relative strength of attraction and
repulsion between the charged particles
depends on the distance separating the
atoms. When the atoms first begin to
interact, the electron-proton attraction is
stronger than the electron-electron and
proton-proton repulsion.
Section 2:Covalent Bonding and
Molecular Compounds
 As they get closer, eventually there is a
distance where that ratio shifts the other
direction. At the point where those two
forces equal each other is where the
atoms are at the lowest potential energy
Section 2:Covalent Bonding and
Molecular Compounds
 Characteristics of the Covalent Bond
 The electrons of each hydrogen atom of
the hydrogen molecule are shared
between the nuclei. They can be pictured
as occupying overlapping orbitals, moving
about freely in either orbital. The bonded
atoms vibrate a bit, but as long as their
potential energy stays close to the
minimum, they will remain covalently
bonded.
Section 2:Covalent Bonding and
Molecular Compounds
 This distance between two bonded atoms
at their minimum potential energy is
known as the bond length. The bond
length between two hydrogen atoms is 75
pm. We saw earlier that the hydrogen
atoms lowered their energy to form H2.
That same amount of energy that was
released to form the hydrogen molecule
would be required in order to break the
bond as well.
Section 2:Covalent Bonding and
Molecular Compounds
 Bond energy – the energy required to
break a chemical bond and form neutral
isolated atoms.
 The energy relationships described here
for the formation of a hydrogen-hydrogen
bond apply generally to all covalent
bonds. However, bond lengths and bond
energies differ depending upon the atoms
in the bond.
Section 2:Covalent Bonding and
Molecular Compounds
 The Octet Rule
 Unlike other atoms, the noble gas atoms
exist independently in nature. They
possess a minimum of energy existing on
their own because of the special stability
of their electron configurations. Their
stability comes from the fact that their
outer energy level is full of electrons.
Other atoms can effectively fill their
valence shell by sharing electrons through
covalent bonding. Such bond formation
follows the octet rule.
Section 2:Covalent Bonding and
Molecular Compounds
 Octet rule – Chemical compounds tend to
form so that each atom, by gaining,
losing, or sharing electrons, has an octet
of electrons in its highest occupied energy
level.
Section 2:Covalent Bonding and
Molecular Compounds
 Electron Dot Notation
 Covalent bond formation usually involves
only the electrons in an atom’s valence
shell. To keep track of these electrons it is
helpful to use electron dot notation.
 Electron-dot notation – An electron
configuration notation in which only the
valence electrons of an atom of a
particular element are shown, indicated by
dots placed around the element’s symbol.
Section 2:Covalent Bonding and
Molecular Compounds
Section 2:Covalent Bonding and
Molecular Compounds
 Lewis Structures
 Electron dot notation can also be used to
represent molecules. For example, a
hydrogen molecule, H2, is represented by
combining the notations of two individual
hydrogen atoms, as follows:
 The pair of dots represents the shared
electron pair of the hydrogen-hydrogen
covalent bond.
Section 2:Covalent Bonding and
Molecular Compounds
 When we use fluorine, you can see the
electrons that are involved in the covalent
bond. You can also see the unshared
pairs of electrons that are surrounding the
fluorine nuclei. The shared electrons most
often are replaced with a dash to
represent the bond.
Section 2:Covalent Bonding and
Molecular Compounds
 These representations are known as Lewis
structures.
 Lewis structures – Formulas in which
atomic symbols represent nuclei and inner
shell electrons, dot pairs or dashes
between two atomic symbols represent
electron pairs in covalent bonds, and dots
adjacent to only one atomic symbol
represent unshared electrons.
Section 2:Covalent Bonding and
Molecular Compounds
 It is common to write Lewis structures that
only represent the electrons that are involved
in bonding. This is called structural formulas.
 Structural formula – indicates the kind,
number, arrangement, and bonds but not the
unshared pairs of electrons in the atom.
 Single bond – a covalent bond in which one
pair of electrons is shared between two
atoms.
Section 2:Covalent Bonding and
Molecular Compounds
 Multiple Covalent Bonds
 Atoms of some elements can share more
than one electron pair. A double covalent
bond, or more simply a double bond, is a
bond in which two pairs of electrons are
shared between two atoms. You can
either show this by having two pairs of
dots or two lines.
Section 2:Covalent Bonding and
Molecular Compounds
Section 2:Covalent Bonding and
Molecular Compounds
 A triple bond, therefore, would be the
sharing of three pairs of electrons
between two atoms. You would represent
this by 6 dots or 3 lines.
 Multiple bonds – double or triple bonds
Section 2:Covalent Bonding and
Molecular Compounds
 Resonance Structures
 Some molecules and ions cannot be
represented adequately by a single Lewis
structure. One such molecule is ozone,
O3.
Section 2:Covalent Bonding and
Molecular Compounds
 Notice that each structure indicates that
the ozone molecule has two types of
bonds, one single and one double.
Scientists used to think that ozone split its
time between these two states, resonating
back and forth between having a single
and double bond. However, experiments
have shown that the bonds in ozone are
identical. Therefore scientists now say
that ozone has a single structure that is
the average of these two structures.
Section 2:Covalent Bonding and
Molecular Compounds
 Resonance – Refers to bonding in
molecules or ions that cannot be correctly
represented by a single Lewis structure.
Section 3: Ionic Bonding and
Ionic Compounds
 Most of the rocks and minerals that make
up Earth’s crust consist of positive and
negative ions held together by ionic
bonding. A familiar ionic compound is
sodium chloride, table salt. A sodium ion,
Na+, has a charge of 1+. A chloride ion,
Cl -, has a charge of -1.
Section 3: Ionic Bonding and
Ionic Compounds
 There is an electrical force of attraction
between oppositely charged ions. In ionic
compounds, the atoms combine in such a
way as to cancel out the charges that
exist on the ions. For this example, it will
take one sodium ion to cancel out the
charge of one chloride ion.
Section 3: Ionic Bonding and
Ionic Compounds
 Ionic compound – Compound composed of
positive and negative ions that are
combined so that the numbers of positive
and negative charges are equal.
 Most ionic compounds exist as a 3
dimensional network of positive and
negative ions. Not as individual molecules
like covalent compounds.
Section 3: Ionic Bonding and
Ionic Compounds
 Therefore, the chemical formula of the
ionic compound does not represent
individual, neutral units, but rather, the
simplest ratio of the compound’s
combined ions that gives electrical
neutrality.
 Formula unit – The simplest collection of
atoms from which an ionic compound’s
formula can be established.
Section 3: Ionic Bonding and
Ionic Compounds
 The ratio of ions in a formula unit depends
upon the charges of the ions combined.
For example, if sodium combined with a
fluoride ion, that would be a one to one
ratio because they both have a magnitude
of 1 in their charge. However, if
magnesium combined with fluoride, it
would take two fluorides to balance out
one magnesium because the charge on a
magnesium ion is 2+.
Section 3: Ionic Bonding and
Ionic Compounds
 When the sodium atom comes into contact
with the chlorine atom, the difference in
electronegativity is very large. Chlorine
rips an electron from the sodium atom,
thus causing the atom to become a
positive ion. With the gaining of another
electron, chlorine becomes a chloride ion.
 Now that we have ions, they are
electrically attracted to one another and
an ionic bond is formed.
Section 3: Ionic Bonding and
Ionic Compounds
 Formation of Ionic Compounds
 Electron dot notation can be used to
demonstrate the changes that take place
in ionic bonding. Ionic compounds do not
ordinarily form from the combination of
isolated ions, but let’s pretend they do for
a minute. This will be easier to
understand.
Section 3: Ionic Bonding and
Ionic Compounds
 Characteristics of Ionic Bonding
 Remember that nature works to lower the
potential energies of the atoms involved in
bonding. In an ionic crystal, ions
minimize their potential energy by
combining in an orderly arrangement
known as a crystal lattice.
Section 3: Ionic Bonding and
Ionic Compounds
Section 3: Ionic Bonding and
Ionic Compounds
 The attractive forces at work within an
ionic crystal include those between
oppositely charged ions and those
between the nuclei and electrons of
adjacent ions. The repulsive forces
include those between like charged ions
and those between electrons of adjacent
ions. The placement of ions in a crystal
lattice represents a balance of these
forces.
Section 3: Ionic Bonding and
Ionic Compounds
 To compare bond strengths in ionic
compounds, chemists compare the
amounts of energy released when
separated ions in a gas come together to
form a crystalline solid.
 Lattice energy – The energy released
when one mole of an ionic crystalline
compound is formed from gaseous ions.
Section 3: Ionic Bonding and
Ionic Compounds
 A Comparison of Ionic and Molecular
Compounds
 The force that holds ions together in ionic
compounds is a very strong overall
attraction between positive and negative
charges. In a molecular compound, the
covalent bonds of the atoms making up
each molecule are also strong.
Section 3: Ionic Bonding and
Ionic Compounds
 However, the forces holding the individual
molecules to other molecules is not. This
difference in strength of attraction for
molecules and formula units gives rise to
several interesting properties.
Section 3: Ionic Bonding and
Ionic Compounds
 The melting point, boiling point, and
hardness of a compound all are based on
how strongly the base units are attracted
to each other. Because the forces
between individual molecules is not very
strong, covalent compounds tend to have
lower boiling and melting points than ionic
compounds.
Section 3: Ionic Bonding and
Ionic Compounds
 In fact, many covalent compounds are
already a gas at room temperature. Ionic
compounds tend to have very high
melting and boiling points because of the
strong attractiveness between units.
Section 3: Ionic Bonding and
Ionic Compounds
 Ionic compounds are hard, but brittle.
Think about it. You basically have several
sheets of ions all stacked upon each other.
They are held in place because of all of
the forces acting upon them. If one of
those sheets gets shifted by just a small
amount, then all of those forces are lined
up in the total opposite direction.
Section 3: Ionic Bonding and
Ionic Compounds
 This would bring about massive repulsive
charges. That sheet of ions has now been
split from the others. This is why when an
ionic compound breaks, it does so very
cleanly.
Section 3: Ionic Bonding and
Ionic Compounds
 In their solid state, ionic compounds
cannot conduct electricity. However, if the
compound can be dissolved, it will conduct
electricity because the compounds break
up into charged particles. Covalent
compounds usually do not conduct
electricity in either case. Even when the
compound dissolves, they just break up
into separate molecules, not charges.
Section 3: Ionic Bonding and
Ionic Compounds
 Polyatomic Ions
 Some atoms bond covalently with each
other to form a group of atoms that has
both molecular and ionic characteristics.
 Polyatomic ions – A charged group of
covalently bonded atoms.
 Polyatomic ions bond with other ions to
form ionic compounds. The charge on a
polyatomic ion comes from having too
many or too few electrons.
Section 4: Metallic Bonding
 Chemical bonding is different in metals
than it is in ionic and covalent. For one
thing, they are great conductors of
electricity, even better than molten ionic
compounds. This is due to the highly
mobile valence electrons of the metals.
Section 4: Metallic Bonding
 The Metallic Bond Model
 The highest energy levels of most metals
are occupied by very few electrons. This
leaves many orbitals vacant. All of these
vacant orbitals in the outer energy levels
overlap with the atom next to them.
Section 4: Metallic Bonding
 The overlapping of empty orbitals allows
the valence electrons of metals to be
delocalized. That means that they do not
belong to any one atom of metal. It is
sort of thought of as a sea of electrons
that have metal atoms packed together in
a crystal lattice.
 Metallic Bonding – the chemical bonding
that results from the attraction between
metal atoms and the surrounding sea of
electrons.
Section 4: Metallic Bonding
 Metallic Properties
 The freedom of electrons to move in a
network of metal atoms accounts for the
high electrical and thermal conductivity
characteristic of all metals. Because they
contain so many orbitals that are
separated by only a small amount of
energy, they are able to absorb a wide
range of light frequencies.
Section 4: Metallic Bonding
 Those electrons will then become excited
and move to a higher energy level. They
cannot hold onto that extra energy and
will release it in the form of light at a
frequency similar to the one they
absorbed. This is what gives metals the
shine or luster that we are familiar with.
Most metals are easily formed into desired
shapes.
Section 4: Metallic Bonding
 Malleability – ability of a substance to be
hammered or beaten into thin sheets.
 Ductility – the ability of a substance to be
drawn, pulled, or extruded through a
small opening to produce a wire.
Section 4: Metallic Bonding
 These properties are possible in metals
because in their lattice, all of the bonding
is the same in all directions. Unlike ionic
crystals with their build up of repulsive
charges when one sheet is moved, metals
have no such build up. Therefore, the
atoms can slide past each other and take
another form.
Section 5: Molecular Geometry
 The properties of molecules depend not
only on the bonding of atoms but also on
molecular geometry. Atoms are not flat,
they have depth. They are 3 dimensional
particles. Molecular geometry is the 3
dimensional arrangements of a molecule’s
atoms.
Section 5: Molecular Geometry
 A chemical formula reveals very little
about a molecule’s geometry. After many
tests designed to reveal the shapes of
various molecules, chemists developed
two different theories to explain certain
aspects of their findings. One theory
accounts for molecular bond angles. The
other is used to describe the orbitals that
contain the valence electrons of a
molecule’s atoms.
Section 5: Molecular Geometry
 VSEPR Theory
 Diatomic molecules, like those of H2 and
HCl, must be linear because only one
bond has been made. However, to predict
the geometries of more complicated
molecules you must consider the locations
of all electron pairs surrounding the
bonded atoms. This is the basis of the
VSEPR theory.
Section 5: Molecular Geometry
 VSEPR stands for “valence shell electron
pair repulsion.” This refers to the
repulsion between pairs of valence
electrons in the atoms of a molecule.
 VSEPR Theory – repulsion between the
sets of valence-level electrons surrounding
an atom causes these sets to be oriented
as far apart as possible.
Section 5: Molecular Geometry
 Let’s examine the simple molecule BeF2.
The beryllium atom forms a covalent bond
with each fluorine atom. According to
VSEPR theory, the shared pairs will be as
far away from each other as possible.
With atoms being three dimensional, the
farthest that two electron pairs can be
from each other is 180o. Thus all three
atoms would be in a straight line, the
molecule would be linear.
 Draw BeF2:
Section 5: Molecular Geometry
 If we allow the central atom to be represented
by the letter A and we represent the atoms
bonded to the central atom as B, then this
would be an example of a AB2 molecule. All
AB2 molecules are linear.
 What would you expect an AB3 molecule to
look like? How many bonds to the central
atom? What degree would the bonds be from
each other?
 Draw BF3:
Section 5: Molecular Geometry
 This is an AB3 molecule. The three bonds stay
as far apart as possible by being 120o away
from each other. This would be called trigonal
planar.
 An AB4 molecule would have four bonds to the
central atom. Most people automatically think
that the bonds would be 90o apart. However,
because the atoms are 3 dimensional, the
bonds can actually be 109.5o apart. This
shape is called a tetrahedral.
 Draw CH4:
Section 5: Molecular Geometry
 Practice: Use VSEPR to predict the
molecular geometry of the following
molecules:
 HI
 CBr4
 CH2Cl2
Section 5: Molecular Geometry
 VSEPR Theory and Unshared Electron
Pairs
 Ammonia, NH3, and water, H2O, are
examples of molecules in which the
central atom has both shared and
unshared electron pairs.
 Draw Lewis Structure of Ammonia:
Section 5: Molecular Geometry
 As you can see, ammonia has three
hydrogen atoms bonded to a central
nitrogen atom. The nitrogen atom has an
unshared pair of electrons as well. VSEPR
theory says that the lone pair occupies
space around the nitrogen atom just as
the bonding pairs do. Therefore, this isn’t
really an AB3 molecule, it is an AB3E
molecule.
Section 5: Molecular Geometry
 We would not call this a tetrahedral
molecule, however. Our description of
shapes of molecules only refers to the
positions of atoms. Therefore, this would
look like a trigonal planar molecule that
has been bent downward a little. This is
called trigonal pyramidal.
 Draw Ammonia:
Section 5: Molecular Geometry
 Water has two unshared electron pairs. It
is an AB2E2 molecule. The oxygen atom
can be thought of as being in the center of
a tetrahedron, with two corners occupied
by unshared electron pairs. Remember
that the shape is only determined by the
atoms, not the unshared pairs. Therefore,
this is a bent molecule.
Section 5: Molecular Geometry
 Draw Water:
 Not a big deal, but unshared electron pairs
repel more than a regular bond. Therefore,
trigonal pyramidal and bent bonds do not
have 109.5o bonds like tetrahedral, but a little
smaller.
 Finally in VSEPR, double and triple bonds are
treated just as a single bond. Polyatomic ions
are treated similarly to molecules. Use this
table to help you predict the geometries for
other types of molecules.
Section 5: Molecular Geometry
Section 5: Molecular Geometry
Section 5: Molecular Geometry
 Hybridization
 VSEPR theory is useful for explaining the
shapes of molecules. However, it does not
reveal the relationship between a
molecule’s geometry and the orbitals
occupied by its bonding electrons. We use
a different model to explain how the
orbitals of an atom become rearranged
when the atom forms covalent bonds.
This model is called hybridization.
Section 5: Molecular Geometry
 Hybridization – The mixing of two or more
atomic orbitals of similar energies on the
same atom to produce new hybrid atomic
orbitals of equal energies.
 We will practice with methane, CH4. The
orbital notation for a carbon atom is as
shown below.
Section 5: Molecular Geometry
 We know from earlier that methane has a
tetrahedral shape. How does carbon form
four equivalent, tetrahedrally arranged
covalent bonds by orbital overlap with four
other atoms?
 Two of carbon’s valence electrons occupy
the 2s orbital and two occupy the 2p
orbital. Remember that 2s and 2p orbitals
have different shapes.
Section 5: Molecular Geometry
 To achieve four equivalent bonds, carbon’s
2s and three 2p orbitals hybridize to form
four new, identical orbitals called sp3
orbitals. The superscript 3 indicates that
three p orbitals were included in the
hybridization. The superscript 1 on the s
is understood. The sp3 orbitals all have
the same energy, which is more than a 2s,
but less than a 2p.
Section 5: Molecular Geometry
 Draw carbon’s orbital notation with
increasing energy:
 Hybrid orbitals – Orbitals of equal energy
produced by the combination of two or
more orbitals on the same atom.
Section 5: Molecular Geometry
 The number of hybrid orbitals produced
equals the number of orbitals that have
been combined. Hybridization also
explains the bonding and geometry of
many molecules formed by group 15 and
16 elements. The sp3 hybridization of a
nitrogen atom yields four hybrid orbitals.
Three that contain an unpaired electron
and one that contains an unshared pair of
electrons.
Section 5: Molecular Geometry
 Each unpaired electron is capable of
forming a bond. Similarly, two of the four
sp3 hybrid orbitals on an oxygen atom are
occupied by two electron pairs and two
are occupied by unpaired electrons. Each
unpaired electron can form a single bond.
Section 5: Molecular Geometry
 The linear geometry of BeF2 is made
possible by hybridization involving the s
orbital and one available empty p orbital
to yield sp hybrid orbitals. The trigonal
planar geometry of BF3 is made possible
by hybridization involving the s orbital and
2 of the p orbitals to make sp2 hybrid
orbitals. Use the following table to help
with recognizing hybridization .
Section 5: Molecular Geometry
Section 5: Molecular Geometry
 Intermolecular Forces
 As a liquid is heated, the kinetic energy of its
particles increases. At the boiling point, the
energy is enough that the force of attraction
between the liquid’s particles is overcome.
The particles will pull away from each other
and enter into the gas phase. Boiling point is
a great measure of the force of attraction
between particles of a liquid. The higher the
boiling point is, the stronger the forces
between the particles must have been.
Section 5: Molecular Geometry
 The forces of attraction is known as
intermolecular forces. Intermolecular forces,
IMF, can vary in strength, but are usually
weaker than covalent, ionic, and metallic
bonds.
 Molecular Polarity and Dipole-Dipole Forces
 The strongest IMF exist between polar
molecules. They act as tiny dipoles because
of their uneven charge distribution.
 Dipole – IMF created by equal but opposite
charges that are separated by a short
distance.
Section 5: Molecular Geometry
 The negative region in one polar molecule
attracts the positive region in adjacent
molecules, and so on throughout a liquid
or solid. The forces of interaction between
polar molecules are known as dipoledipole forces. Molecules that display this
type of IMF will show a much higher
melting and boiling point because of the
extra attractions.
Section 5: Molecular Geometry
 Polar molecules can also induce a dipole
moment in other molecules that usually
do not exhibit such behavior. For
instance, water can cause O2 to
temporarily display dipole behavior
because it is highly polar. However, the
induced dipole is temporary and very
weak.
Section 5: Molecular Geometry
 Hydrogen Bonding
 Some hydrogen containing molecules
exhibit very high boiling points because of
an unusually strong type of dipole-dipole
force known as hydrogen bonding. We
see hydrogen bonding in such molecules
as water, ammonia, and hydrogen
fluoride. In each of these molecules,
hydrogen is bonded with an atom that is
very electronegative.
Section 5: Molecular Geometry
 This causes the electrons to spend more
time around the central atom than
hydrogen. This leaves a partial positive
charge on what is essentially a proton
hanging off of another atom. Because
there are no other electrons on the
hydrogen, it can get very close to other
molecules. This creates a very strong IMF.
Section 5: Molecular Geometry
 Hydrogen Bonding – The IMF in which a
hydrogen atom that is bonded to a highly
electronegative atom is attracted to an
unshared pair of electrons of an
electronegative atom in a nearby
molecule.
Section 5: Molecular Geometry
 London Dispersion Forces
 Even noble-gas atoms and molecules that
are nonpolar experience a weak IMF.
Electrons are in constant motion. At any
moment the electrons could overload one
side of the atom just by chance. This
causes a very temporary dipole moment.
The molecule or atom would then attract
another molecule or atom that is close to
it.
Section 5: Molecular Geometry
 London Dispersion Forces – The IMF
attractions resulting from the constant
motion of electrons and the creation of
instantaneous dipoles.
 These type of IMF can occur in any
molecule, but they are the only IMF that
occurs when you have noble gas atoms.