MOLECULAR SHAPE
Chapter 8
THE SHAPE OF SMALL
MOLECULES
8-1
Molecular Geometry
• The behavior of atoms is determined chiefly by
their electron configurations
• The behavior of molecules also depends on their
structural characteristics
• In this section you will look at the shapes of
molecules and what characteristics of their bonds
produce those shapes
Molecular Geometry
• Two ways of looking at the structure of molecules
to account for their shapes:
• The first model takes into account the repulsive forces
of electrons pairs around an atom
• The second model considers ways in which atomic
orbitals can overlap to form orbitals around more than
one nucleus.
• The electrons in these combined orbitals then serve to bind the
atoms together
Molecular Geometry
• In order to describe the shape of a molecule or polyatomic
ion, it is useful to draw a Lewis electron dot diagram
• For all atoms tat form covalent bonds, except hydrogen, eight
electrons represent a full outer level.
Molecular Geometry
• Take a water molecule for example:
It is the only arrangement of electrons in which all
three atoms can achieve a full outer level.
Notice, that 2 electrons in the outer level of Oxygen
are involved in bonding the Hydrogens.
These are called shared pairs.
The other 2 pairs of electrons are not involved in the
bonding, they are called unshared pairs or lone
pairs.
Molecular Geometry
• Notice that the shared electrons contribute to a full outer
level for both atoms sharing the electrons
Molecular Geometry
• Electron Pair Repulsion:
• One way of looking at molecules is to consider
electron repulsion.
• Each bond and each unshared pair in the outer
level of atom form a charge cloud which repels all
other charge clouds.
• In part, this repulsion is due to all electrons
having the same charge.
Molecular Geometry
• Another more important factor is the Pauli exclusion
principle.
• Although electrons of opposite spin may occupy same volume
of space, electrons of the same spin may not do so.
Molecular Geometry
• The repulsions resulting from the Pauli principle are
greater than the electrostatic ones at small distances.
• Because of these repulsions, atoms cannot be
compressed.
Molecular Geometry
• The repulsions between the charge clouds in the outer
level of atoms determine the arrangement of the orbitals.
• The orbital arrangement, in turn, determines the
shape of molecules.
Molecular Geometry
• Structural formula: ammonia
• Gives number of atoms and unpaired electrons.
• Does not indicate the shape of an ammonia molecule.
• Only identifies bonds, but not how bonds are arranged in
space.
Molecular Geography
• Molecular models
• Spheres = nucleus and inner-level electrons
• Sticks = bonds
• Three dimensional
• Symmetrical
Molecular Geometry
• Why are atoms arranged
symmetrically?
• Valence electrons are found in pairs
• Valence electrons repel other electron pairs
because of similar electric charges.
• VSEPR theory – valence-shell
electron pair repulsion theory.
• In a small molecule, the pairs of valence
electrons are arranged as far apart from
each other as possible.
Molecular Shapes
• As result, the following rule may be stated:
• Electron pairs spread as far apart as possible to
minimize repulsive forces.
• If there only two electron pairs in the outer level,
they will be on opposite sides of the nucleus.
• The arrangement is called linear.
Molecular Shapes
• Linear
• Atoms are in a straight line
• All molecules that contains only 2 atoms.
• Bond angle – 180°
• Because atoms are arranged as far apart as possible.
• Example: CO2
Molecular Shape
• Trigonal Planar
• Triangular, flat shape
• Bond angle = 120
• Usually have a central atom
that is bonded to three other
atoms and the central atom
has no unshared pairs of
electrons.
• Example: Boron trichloride
(BCl3)
Molecular Shape
• Tetrahedral
• Tetra = 4
• A shape with 4
surfaces
• Three dimensional
• Bond angle = 109.5
• Example: Methane
(CH4)
Molecular Shape
• Pyramidal
• Represents a shape with unshared
pair of electrons.
• All pairs of valence electrons repel
each other equally.
• Unshared pairs exert a greater
repulsion force (take up more room).
• Bond angle = 107°
• Usually have a central atom bonded
to 3 other atoms and an unshared
pair of valence electrons.
• Example: Ammonia (NH3)
Molecular Shape
• Bent
• Example: Water H2O
• Oxygen in the central atom with two bonds to hydrogen
and two pair of unshared electrons.
• The two unshared pairs around the oxygen atom exert a
greater repulsion force that the two electron pairs in the
bonds.
• 105° bond angles
Molecular Shape
• The bonds and unshared electron pairs
determine the shape a molecule.
• An unshared pair is acted upon by only one
nucleus.
• It's charge cloud is like a very blunt pear, Figure
13-2, with its stem end at the nucleus.
Molecular Shape
• A shared pair of electrons moves
within field of two nuclei.
• The cloud is more slender.
Molecular Shape
• The electron pair repulsions in a
molecule may not all be equal.
• The repulsion between two unshared
pairs is greatest when they occupy the
most space.
Molecular Shape
• The repulsion between shared pairs is
least because they occupy the least
space.
• The repulsion between an unshared
pair and a shared pair is an
intermediate case.
Molecular Shape
• unshared-unshared repulsion >
unshared-shared repulsion >
shared-shared repulsion
• Electron pair repulsion strengths may not be equal.
Molecular Shape
• Let us look at the molecular shapes of the
compounds CH4, ,H2O, and NH3 to illustrate this
repulsion.
• In each of these compounds, the central atom
has four clouds around it.
• We expect the axes of all four charge clouds to
point approximately to the corners of a
tetrahedron.
Molecular Shape
• In methane molecules all clouds are
shared pairs, so their sizes are equal
and each bond angle is in fact 109.5o
• Methane is therefore a perfect
tetrahedron
Molecular Shape
• In NH3 molecules, there are one
unshared pair and three shared pairs
• The unshared pair occupies more
space than any of the other three, so
the bond clouds are pushed together
and form an angle of 107o with each
other
Molecular Shape
• Although the electron clouds form a
tetrahedron one cloud is not involved
in bonding.
• Therefore, the atoms composing the
molecule form a trigonal pyramid
Molecular Shape
• In H2O molecules, two unshared pairs
are present
• Both of these clouds are larger than
the bond clouds
• This additional cloud size results in a
still greater reduction in the bond
angle which is, 104.5o
Molecular Shape
• Note that the electron clouds are
tetrahedral but the molecule is “V”
shaped, or bent
Molecular Shape
• Note that in the 3 molecules
discussed, each has 4 electron clouds.
• The differences in molecular shape
result from the unequal space
occupied by the unshared pairs and
the bonds
Molecular Shape
• In most compounds, the outer level is considered
full with four paris or 8 electrons
• The outer level in some atoms can contain more
than eight electrons
• If the outer level is the third or higher level
• Some nonmetals, but mainly halogens form
compounds in which the outer level is expanded
to 10, 12, or 14 electrons
Molecular Shape
• Other shapes:
• T-shaped
• Square Planar
• Trigonal bipyramidal
• Octahedral
Practice Problems
• What is the molecular shape of nitrogen trifluoride? (NF3)
• What are the bond angles?
• What is the molecular shape of carbon tetrachloride?
(CCl4)
• What are the bond angles?
Hybrid Orbitals
• Electrons are found in orbitals around the nucleus
• 1s, 2s, 2p….
• Orbitals do not explain the electrons in bonds of a
molecule.
• When atoms bond, the electrons are found in hybrid
orbitals.
• Atomic orbitals of different atoms “mix” together.
• They have a combination of the properties of the atomic orbitals
that formed them.
Hybrid Orbitals
• In nature there are many different possibilities.
• Linear – mix of s and p orbital = sp orbital
• Trigonal planar – mix of s and 2 p orbitals = sp2 orbital
• Hybrid orbitals are often used to categorize molecular
shape.
Hybrid Orbitals
• Methane is the bonding of 4 H to 1 C
• The bonds involve the overlap of the s
orbital of each H atom with one of the
sp3 hybrid orbitals of a C atom
Hybrid Orbitals
• There is an angle of 109.5o between
each carbon-hydrogen bond axis
Hybrid Orbitals
• When 2 carbon atoms bond their sp3
overlap. The 3 remaining sp3 orbitals
may bond with the s orbital of 3
hydrogen atoms
Hybrid Orbitals
• A covalent bond is formed when an
orbital of one atom overlaps an orbital
of another atom and they share the
electron pair the bond.
• For example, a bond may be formed
by the overlap two s orbitals.
Hybrid Orbitals
• A bond formed by the direct overlap of
two orbitals is called a sigma bond, and
is designated σ.
Hybrid Orbitals
• A sigma bond is also formed by the
overlap of an s orbital of one atom with
a p orbital of another atom,
• the overlap of 2 p orbitals,
• the overlap of 2 hybrid orbitals,
• or the overlap of a hybrid orbital with an s orbital
Hybrid Orbitals
• Because p orbitals are not spherical,
when 2 half-filled p orbitals overlap,
one of two types of bonds can form
• 1. If 2 p orbitals overlap along an axis in
and end-to-end fashion, a sigma bond
forms
Hybrid Orbitals
• 2. If the 2 p orbitals overlap sideways
(parallel), they form a pi bond,
designated π.
Hybrid Orbitals
• Pi bonds are always formed by the
sideways overlap of unhybridized p
orbitals
Hybrid Orbitals
• Double bonds are 2 pairs of electrons that
are shared between the bonding atoms.
• A double bond always consists of one
sigma bond and one pi bond
Hybrid Orbitals
• In a triple bond 3 pairs of electrons are
shared between the bonded atoms
• 2 sp hybrid orbitals, one from each carbon,
overlap to form 1 sigma bond
• The 2 p orbitals from atom overlap to form 2 pi
bonds
Hybrid Orbitals
• Both double and triple bonds are less
flexible than single bonds are, and they are
also shorter
• Pi bonds are more easily broken than
sigma bonds are because the electrons
forming pi bonds are farther from the nuclei
of the 2 atoms
Hybrid Orbitals
• So molecules containing multiple bonds are
usually more reactive than are similar
molecules containing only single bonds
• Compounds that contain double or triple
bonds between carbon atoms are called
unsaturated compounds
Hybrid Orbitals
• If atoms share more than one pair of
electrons, all atoms in the molecule
can have full outer levels
Hybrid Orbitals
• How does the electron-pair repulsion
theory predict the shapes of molecules
containing multiple bonds?
Hybrid Orbitals
• Remember that double bonds consist of 4
electrons occupying the space between the
bonded atoms
• The resulting cloud will occupy more space
than a single bond
• The triple bond occupies still more space
than the double bond because it has 6
electrons being shared
Hybrid Orbitals
• How is molecular shape affected by
the presence of multiple bonds?
Hybrid Orbitals
• The methanal molecule below has a
double bond and 2 single bonds
• This would form a trigonal planar
shape
Hybrid Orbitals
• Bond types
• H–C–H
• H–C=O
• C=C=O
• H–C=O
•
effect the bond angles:
116
122
180
120
180
Hybrid Orbitals
• Because the pi electrons are shared
equally among all the carbon atoms and
not confined to one atom or bond, they are
delocalized.
• This delocalization of pi electrons among
the carbon atoms in benzene results in
greater stability of the compounde
Hybrid Orbitals
• Whenever multiple p orbital overlap
can occur, the molecule is said to
contain a conjugated system
• This can also occur in rings
Bond Length
• Different pairs of atoms form bonds of different lengths.
• Trends:
• Moving down a group in the periodic table – atoms form
longer bonds.
• Atoms become larger as you move down a group.
• Multiple bonds are shorter than single bonds.
• The more electrons in a bond, the stronger that bond attracts the
positively charged nuclei of the bonding atom.
HOMEWORK
Pg. 265 1-4
POLARITY
8-2
Polarity
• In a polar bond electrons are
shared unequally between 2
atoms.
• Electrons are pulled closer to the
more electronegative atom
giving it a slight negative charge
and the other atom a slight
positive charge.
• In a nonpolar bond, electrons
are shared equally.
Polarity
• Molecules can also be polar or
nonpolar.
• A polar molecule has one end with a
positive charge and another end with
a negative charge.
• Dipoles – polar molecules
• Polarity gives molecules different
properties:
• Align in electric fields
• Attracted to or deflected by a magnetic
field
Determining Polarity
• Any molecule that is composed of only one kind of atom is
a nonpolar molecule.
• Only have nonpolar bonds.
• H 2 , O2
• A molecule that contains polar bonds is not necessarily a
polar molecule.
• Example: CO2
Determining Polarity
• To determine whether a molecule is polar,
you need to look at its shape.
• The shape of a molecule and the polarity
of its bonds together determine whether
the molecule is polar or nonpolar.
Determining Polarity
• Because no 2 elements have exactly
the same electronegativities, in a
covalent bond between different
elements, one of the atoms attracts
the shared pair more strongly than the
does the other.
Determining Polarity
• The resulting bond is said to be polar covalent.
• In this bond, the atom with higher
electronegativity attracts the electrons more
strongly, and that end of the bond will have a
partial negative charge.
• The bond at the other end of the bond will have a
partial positive charge
Water
• Very important molecule
• Liquid state at room temperature
and part of almost every liquid on
earth.
• Liquid because:
• Positive hydrogen end of one water molecule
attracts to the negative oxygen end of
another water molecule.
• Loosely bonds molecules together.
• Only compounds found in nature as
solid, liquid, or gas.
Determining Polarity
• Partial charges within a molecule are indicated by δ
(delta)
• Water molecules are a good example of this.
Determining Polarity
• Polar bonds may produce polar molecules
• To be polar the charges must be unequal.
• To be a nonpolar molecule the charges
must be pulling in equal strength and
therefore, cancel each other out
Determining Polarity
• What about the water molecule, is it polar
or nonpolar?
• What about the carbon dioxide molecule, is
it polar or nonpolar
Determining Polarity
• In carbon dioxide molecules the carbon-
oxygen bond is polar because oxygen has
a greater electronegativity than carbon
does.
• However, the polarities of the two bonds
are in exactly opposite directions and so
they cancel each other out.
• This does not occur with water.
Carbon Dioxide
• CO2
• Linear molecule
• Two carbon-oxygen double bonds.
• Carbon-oxygen bonds are polar.
• But carbon dioxide is not a polar molecule:
• Positive charge is concentrated in the center.
• Negative charge is divided equally on both sides.
• Being nonpolar gives carbon dioxide important
properties:
• Molecules have little attraction to each other making
carbon dioxide a gas at room temperature.
Determining Polarity
• Water has a bent/angular geometry so
the bonds aren’t exactly opposite from
each other.
• Therefore, they don’t cancel each
other out.
• So water is polar
Formaldehyde
• Used to preserve biological
specimens.
• CH2O
• Carbon forms bonds with 3 other
atoms.
• Oxygen atom has highest
electronegativity.
• Electrons in the C-O bond are attracted
more towards the oxygen.
• Oxygen becomes partially negative and
carbon partially positive.
• Carbon more electronegative than
hydrogen.
• Difference in negative and positive
partial charges makes molecule polar.
Determining Polarity
• Polar bonds are a necessary but not a
sufficient condition for polar molecules.
• In a polar molecule, the polar bonds cannot
be symmetrically arranged.
Determining Polarity
• Because it has both a positive and a negative
pole, a polar molecule, such as water, is also said
to be dipole, or to have a dipole moment.
• Not to be confused with a shiny hair moment
• A dipole moment is a measure of the strength of
the dipole and is a property that results from the
asymmetrical charge distribution in a polar
molecule
Dipoles
• The dipole moment depends upon the size of the
partial charges and the distance between them.
• μ is the dipole momement, q is the size of the
partial charge in coulombs and r is the distance in
meters between the partial charges.
• The units are in coulomb x meters
Dipoles
• The higher the dipole moment, the stronger the
intermolecular forces; and, consequently, the higher the
melting point and boiling point for molecules of similar
mass.
Dipoles
• Van der Waals forces are sometimes
referred to as weak forces because they
are much weaker than chemical bonds.
• Weak forces involve the attraction of the
electrons of one atom for the protons of
another.
Dipoles
• Intramolecular forces are forces within a
molecule that hold atoms together, that is,
covalent bonds
• Intermolecular forces are forces between
molecules that hold molecules to each
other, that is, van der Walls forces
Dipoles
• The first van der Waals force is the dipole-
dipole force.
• With dipole-dipole forces, two molecules of
the same or different substance that are
both permanent dipoles, will be attracted to
each other.
Dipoles
• A dipole can also attract a molecule that is
ordinarily not a dipole.
• When a dipole approaches a nonpolar
molecule, its partial charge either attracts
or repels the electrons of the other particle.
Dipoles
• For instance, if the negative end of the dipole
approaches a nonpolar molecule, the
electrons of the nonpolar molecule are
repelled by the negative charge.
• The electron cloud of the nonpolar molecule is
distorted by bulging away from the approaching
dipole as shown in
• Figure 14-3.
Dipoles
• As a result, the nonpolar molecule is
itself transformed into a dipole.
• We say it has become an induced
dipole.
• Since it is now a dipole, it can be
attracted to the permanent dipole.
Dipoles
• Interactions such as these are called dipole-
induced dipole forces.
• An example of this force occurs in a water
solution of iodine.
• The I2 molecules are nonpolar while the water
molecules are highly polar.
• The case of two nonpolar molecules being
attracted must also be taken into account.
Dipoles
• For instance, there must be some force
between hydrogen molecules; otherwise it
would be impossible to form liquid
hydrogen.
• Consider a hydrogen molecule with its
molecular orbital including both nuclei.
Dipoles
• We know intuitively that the electrons
occupying that orbital must have a specific
location.
• If they are both away from one end of the
molecule for an instant, then the nucleus is
exposed for a short time.
Dipoles
• That end of the molecule has a partial positive
charge for an instant; a temporary dipole is set
up.
• For that time, the temporary dipole can induce a
dipole in the molecule next to it and an attractive
force results as shown in Figure 14-4.
• The forces generated in this way are called
dispersion forces.
Dipoles
• The various kinds of interactions making van der
Waals forces affect each other, but we are only
interested in the net result.
• The liquid and solid states of many
compounds exist because of these
intermolecular forces.
Dipoles
• These forces are effective only over very
short distances.
• They vary roughly as the inverse of the
sixth power of distance.
• In other words, if the distance is doubled,
the attractive force is only 1/64 as large.
Dipoles
• Of the three contributing factors to van
der Waals force, dispersion forces are
the most important.
• They are the only attractive forces that
exist between nonpolar molecules.
Dipoles
• Even for most polar molecules,
dispersion forces account for 85% or
more of the van der Waals forces.
• Only in some special cases, such as
NH3 and H20, do dipole-dipole
interactions become more important
than dispersion forces.
Large Molecules
• Small molecules – the shape
helps to determine polarity.
• Large molecules – the polarity
often helps to determine its
shape.
• Example: Protein
• Essential to all living things. Build and
repair cells and components of many
cell structures.
• Extremely large molecules.
(thousands of atoms)
• Composed of individual subunits into
a chain.
Large Molecules
• Subunits have polar
sidechains
• Molecule bent and
twisted because polar
sections attracted to each
other.
• Large molecules have
a large variety of
shapes.
• Geometry around
individual atoms is
identical to small
molecules.
HOMEWORK
Pg. 273 1-4