Molecular Geometry: Using molecular models to determine the
structure of covalent molecules and polyatomic ions
OBJECTIVE
In this activity, you will learn how to predict the three-dimensional shapes of small molecules and
polyatomic ions by drawing Lewis structures and by using the principles of the Valence Shell
Electron Pair Repulsion (VESPR) model. You will build ball and stick models of these species so
as to see the shapes and draw them. In addition, the polarity of bonds between atoms in molecules
will be examined and the molecules will be determined to be overall polar or nonpolar.
APPARATUS
Molecular modeling kit
DISCUSSION
Lewis Dot Structures
A Lewis dot structure is a two-dimensional representation that shows the arrangement of atoms in
a molecule, and includes both bonding electrons and nonbonding electrons. In particular, Lewis
structures are useful for showing covalent and polar covalent bonds. The formation of covalent
bonds can be represented for H2 as shown below:
H + H
H
H
The formation of a bond between an H atom and an F atom can be represented similarly:
H + F
H
F
Typically, the bonding electrons are shown by a line, to emphasize that it is a bond, and the
nonbonded electrons (the electron lone pairs) are shown as dots, as thus:
H + F
H
F
By sharing the electrons, each atom obtains the noble gas configuration. For hydrogen, the noble
gas configuration that yields stability is the helium configuration with two valence electrons. For
all other atoms, the noble gas configuration has eight valence electrons; hence, the formation of
bonds so that each atom can resemble a noble gas is referred to as the “octet rule”. Note that the
octet rule often fails: in reality, only a few elements consistently achieve the outer electronic
configuration of the nearest noble gas when they combine to form molecules. These elements are
in the first and second periods of the periodic table, most notably H, C, N, O, and F. Elements in
the third and higher periods can “expand” their capacity to hold more than eight valence electrons
by using empty orbitals in their valence shells, thus allowing them to accommodate 10, 12, or even
14 electrons when they are present as the central atom of a molecule. This allows additional bonds
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and lone pair electrons to be present. Boron is another exception, in that it often forms molecules
such as BF3, etc., where the central boron atom contains only six valence electrons.
Multiple bonds are formed when atoms obtain their complete octets by sharing more than one pair
of electrons between them. When two electron pairs are shared, two lines, representing a double
bond, are drawn. A triple bond corresponds to the sharing of three pairs of bonding electrons
between two atoms. Two examples of multiple bonding are found in CO2 and N2:
O
C
O
N
N
Drawing Lewis structures (using NO3- as an example):
1. Calculate the total valence electrons available on all the atoms in the species (molecule or
polyatomic ion) and…
…if it is an anion, add one electron for each negative charge.
…if it is a cation, subtract one electron for each positive charge.
For NO3- :
N = 5 valence electrons
O = 6 valence electrons
O=6“
O=6“
charge = 1 e_______________________________
= 24 electrons available
2. Draw the skeleton structure of the molecule or ion, joining the bonded atoms with single bonds
(one pair of bonded electrons). In general, the atom that is most likely to be in the center is one
that has a lower electronegativity, so, comparing various atoms, one would first see a C atom in
the middle spot, then N, then O, then a halogen, etc. For NO3-, the nitrogen would be the central
atom:
O
N
O
O
3. Drawing the skeleton structure has used some of the available electrons for bonding electrons.
Calculate how many are left and distribute the remaining electrons as nonbonding electron pairs
about each atom that is connected to the central atom until each is surrounded by eight electrons
(just two electrons for a H atom though). Any extra electrons still left should then be put on the
central atom, even if it means that it will have more than eight valence electrons. [We are not yet
finished – make sure to read #4!]
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For NO3-, there are three single bonds drawn, two electrons per bond, so we have 24 – 6 = 18
electrons left, allowing us to put 3 nonbonding pairs of electrons onto each oxygen:
O
N
O
O
4. Now examine the central atom: if it does not yet have a full octet, convert one or more single
bonds to double or triple bonds by converting one or more nonbonding pairs of electrons to shared
pairs.
O
N
O
O
Note that bonding electrons are counted as belonging to BOTH atoms that share them. For NO3-,
any of the three oxygen atoms can share one of its nonbonding electron pairs with the central N
atom. This means that resonance structures can be drawn for this molecule.
Sharing a pair with the nitrogen gives one of three possible resonance structures for this
polyatomic ion:
O
N
O
O
Predicting molecular shapes using VESPR theory:
Many methods have been used to model the three-dimensional shape of a molecule. Three
commonly used methods are shown below for carbon tetrachloride, CCl4, which is tetrahedral.
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Ball and stick models are in particular a very useful tool for gaining a picture of the threedimensional shape as they can be formed with any of the commonly-found molecular modeling
sets that can be bought as well as with simple tools such as Styrofoam balls and toothpicks. You
will use the latter today to make your molecular models in this lab.
Ball and stick models of the eleven most common molecular structures for small molecules and
polyatomic ions are shown below.
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Once an accurate Lewis structure is drawn for a molecule or polyatomic ion, it can be used to
predict the shape of the species. The technique for doing this was developed in 1957 by a
Canadian chemist, R. J. Gillespie, called the Valence Shell Electron-Pair Repulsion (VESPR)
Theory. He realized that the natural electrostatic repulsion that exists between electron pairs, both
bonding and nonbonding, will cause them to be as far apart as possible. In this theory, it is the
domains of electron density that have the effect on neighboring electron domains, so a single bond
is treated the same as a double bond or triple bond when predicting the electronic geometry of the
molecule. By orienting the bonding and nonbonding electron pairs, or the areas of electron
density, as far apart as possible, the molecule will exist in its most stable shape. One can then
determine first the electronic geometry around the central atom where one counts all electron
domains, those with atoms attached as well as lone pairs, and then the molecular geometry around
the central atom can be found, where the shape is determined from the positions of the atoms
around the central atom alone. (Here note that while the atomic placement alone determines the
molecular geometry, the atomic positions are determined still by both the bonding and the
nonbonding electron pairs.)
To predict electronic geometries and molecular shapes:
1. Derive the Lewis structure of the species. If the species displays resonance, choose one
resonance form.
2. Identify the central atom in the species around which you will determine the electronic
geometry and the molecular geometry.
3. Total the number of areas of electron density around the central atom. Each nonbonding pair of
electrons counts as one area, each single bond counts as one area, each double bond counts as one
area, and each triple bond counts as one area. The total of all the areas of electron density predicts
a specific electronic geometry as shown below in Table 1, where each shape has the areas of
electron density arranged as far apart as possible.
4. To determine molecular shape, look at the same arrangement of electron density areas around
the central atom but only take into account for the structure the ones that are shared between two
atoms. Then refer to the pictures on the previous page to determine the shape. This bears
restating: The shape of the molecule or ion is determined by the arrangement of the atoms in
space. Although the nonbonding electron pairs are very important in determining this arrangement
of atoms, only the geometric arrangement of the atoms is used to describe molecular shapes.
Thus a molecule like NH3, which has 4 areas of electron density around it and possesses
tetrahedral electron geometry, has only three atoms attached to the central atom (but one spot still
taken by a lone pair of electrons), and it will have a trigonal pyramidal shape (NOT trigonal
planar):
N
H
H
H
In the more complicated shapes such as trigonal bipyramidal and octahedral arrangements of areas
of electron density, nonbonding electron pairs (lone pairs) always occupy specific locations,
shown in Table 2 below.
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TABLE 1: The five arrangements of electron domains about a central atom
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TABLE 2: Location of bonding and nonbonding electron pairs in trigonal bipyramidal and
octahedral geometries
Bond polarity, molecular polarity and dipole moments
A covalent bond between two atoms of different elements is usually a polar bond, due to the
different electronegativities of the two elements, which causes the bonding electrons to be
unequally shared between the two atoms. This creates a bond dipole moment, µ, which is the
charge separation (Q) multiplied by the distance (r):
µ = Qr
The bonding electrons have an increased attraction to the more electronegative atom, thus creating
an excess of electron density (!-) near that atom and a deficiency of electron density (!+) near the
less electronegative atom. We symbolize the bond dipole (a vector, with magnitude and direction)
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with an arrow with a cross on the end, with the point of the arrow pointing towards the negative
end of the dipole and the cross representing the positive end (
).
To determine if a molecule is polar, one first looks at the individual bond to determine if any have
polarity. If bond polarity exists, then the molecule may be polar, if the bond dipoles do not cancel
each other out. The molecular dipole moment, if it exists, is the vector sum of the individual bond
dipoles. As examples, consider the molecules H2O and CCl4 (see Figure below). In both of the
molecules, every bond in each molecule is between two atoms of differing electronegativity and is
thus polar. In H2O, the bond dipoles give a vector sum that is nonzero, making the molecule polar
(and thus it possesses a molecular dipole moment). In CCl4, although each bond has a dipole
moment, the vector sum is zero. There is no molecular dipole for CCl4 since the symmetry of the
molecule ensures that the bond dipoles cancel. Therefore, CCl4 is a nonpolar molecule.
Bond length and bond angle
We have discussed thus far the three-dimensional arrangement of atoms in a molecule, which is
called its structure, shape or spatial geometry. Structures of small molecules can also be
described in terms of bond lengths and bond angles. A bond length is the average distance
between the nuclei of two bonded atoms and is usually expressed in units of picometers (1 pm =
10-12 m) or Angstroms (1Å = 10-10 m, 1 Å = 100 pm). A bond angle is the angle of arc between
any two bonds that join the same atom. If all bond angles and bond lengths are known, then the
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complete three-dimensional shape of that molecule can be reproduced. Ideal bond angles for a few
of the different molecular geometries are as follows:
Molecular geometry
Ideal bond angle
Linear
180˚
Trigonal planar
120˚
Tetrahedral
109.5˚
Trigonal bipyramidal
90˚; 120˚
Because lone pairs (nonbonding electron pairs) are larger (take more space) than bonding electron
pairs, lone pair-lone pair electrostatic interactions are greater than lone pair-bonding pair
interaction, which in turn are greater than bonding pair-bonding pair interactions. Using this idea,
one can see that the typical bond angles seen will be distorted when lone pairs are present, making
the bond angles other than the nondistorted, ideal values. For example, in NH3 the electron
geometry is tetrahedral and the molecular shape is trigonal planar. Simple VESPR theory predicts
the bond angles for each "H-N-H to be 109.5˚. However, since one of the four tetrahedral
positions is taken up by a lone pair of electrons, the three bonding electron pairs are “squeezed
together” resulting in bond angles less than the ideal (the experimentally measured bond angle for
NH3 is "H-N-H = 107.3˚).
The bond lengths and bond angles for molecules that contain double or triple bonds are also
slightly different than ideal. As an example, the values for a formaldehyde molecule are shown
below:
Carbon is the central atom and the other atoms are bonded to it. It is the orientation of the
hydrogen and the oxygen atoms about the carbon atom that determines the structure of the
molecule. Formaldehyde is a flat, planar molecule; the fact that the bond angles total 360˚ verifies
this.
REPORT SHEET
MOLECULAR GEOMETRY LAB
1. Using an appropriate set of models, make molecular models of the compounds listed below and
complete the table. On the next page you will draw the Lewis dot structure of each molecule. Use
your models to guide you in understanding the 3-D geometry of these compounds.
Molecular
formula
BeCl2
BF3
ClF3
NO
NH3
H2O
POCl3
SF4
BrF3
XeF2
SF6
IF5
XeF4
# of
bonding
pairs
# of
nonbonding
pairs on
central
atom
Electronic
geometry
Molecular
shape
Bond
angle(s)
Molecular
dipole
moment
(yes or no)
2
Molecular formula
BeCl2
BF3
ClF3
NO
NH3
H2O
POCl3
Lewis dot structure, indicating any polar bonds and overall
polarity of the molecule
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SF4
BrF3
XeF2
SF6
IF5
XeF4
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2. Draw the Lewis dot structure of the following ions, including any resonance structures, and
predict the electronic geometry and the molecular geometry.
Ion
N3 -
CO32-
NH4+
NO3-
Lewis dot structure(s)
Electronic
Molecular
geometry
geometry
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H3O+
BF4-