Molecular Geometry
Objectives
The objectives of this laboratory are to:

Write Lewis structure representations of the bonding and valence electrons in molecules.

Use the VSEPR model to predict the molecular geometries (shapes) of molecules.

Determine whether a molecule is polar (has a dipole moment) or not.
Background
The physical and chemical properties of a molecule are a direct consequence of the way in which the
atoms are bonded together and the three-dimensional geometry, or shape, of the molecule.
For example, water is a polar molecule, due in part to it’s bent geometry:
The polarity of a molecule (or lack thereof) influences physical properties such
O
solubility, boiling point, freezing point, vapor pressure, and chemical reactivity,
H
H
to name a few. An understanding of molecular geometry is central in
determining the chemical and physical properties of molecules. The first step in
determining molecular geometry is to write the Lewis structure for a molecule, which gives the
arrangement of valence electrons in terms of bonding pairs and lone pairs of electrons. The Valence
Shell Electron Pair Repulsion (VSEPR) model is then applied to determine a molecule’s threedimensional geometry (or shape). The basic premise of the VSEPR approach is that the electron pairs
surrounding an atom in a molecule repel each other, and tend to minimize the repulsions by positioning
themselves as far away from each other as possible. The application of VSEPR permits the prediction of
molecular geometries and bond angles. This information can be used to subsequently determine other
properties, such as molecular polarity.
For example, consider ammonia, NH3, which has the Lewis structure:
H
N
H
The nitrogen atom has four pairs of valence electrons, 3 bonding pairs
H
(shared with H atoms) and 1 lone pair (unshared). The VSEPR model
predicts that the four pairs minimize repulsions by adopting a tetrahedral arrangement around the
nitrogen. The arrangement of all electron pairs – bonding pairs and lone pairs – around the central atom
is referred to as the electron pair arrangement. By contrast, the molecular geometry does not include the
lone pairs, but only describes the geometric arrangement of the atoms in the molecule. Thus, in the
example of NH3, the electron pair arrangement is tetrahedral while the molecular geometry is trigonal
pyramidal (shown at the right).
VSEPR allows the prediction of ideal bond angles; for the tetrahedral electron
pair arrangement of the NH3 molecule, the ideal H-N-H angle would be
predicted to be 109.5°, the tetrahedral angle. Table 1 lists the molecular
geometries predicted via the application of the VSEPR model.
N
H
H
H
Before coming to lab please review these sections in the textbook (Zumdahl 8.1-8.3 & 8.7-8.11) with
particular focus on Lewis structures, molecular geometry & VSEPR, and molecular polarity.
Procedure for Drawing Lewis Structures:
1. Add up the total # of valence electrons for all the atoms. Account for charge: If the species has a
negative (–) charge: add one valence electron for each negative charge; for a positively charged (+)
species, subtract one electron for each positive charge.
2. Draw the molecular skeleton and connect the atoms with one bond. The central atom is generally
the atom with the lowest electronegativity, but never H.
3. Satisfy the Octet rule: Distribute the remaining valence electrons by adding lone pairs to complete
the octets of the outer atoms first (H only requires 2 electrons), then place any remaining e–’s on the
central atom.
4. If there are too few valence electrons to give each atom an octet, multiple bond(s) may be required. In
this case, convert outer atom lone pairs to bond pairs to form multiple bonds.
5. There are some exceptions to the “octet rule”: 3rd period or heavier elements may more than 8
electrons around them if needed (expanded valence shells). Atoms such as Be, B and Al may have less
than an octet of electrons. In species with an odd number of electrons the least electronegative atom
carries the odd electron.
VSEPR Procedure:
1. Draw the Lewis structure
2. Determine the number of electron pair domains on the central atom. This is is equal to the number
of atoms plus the number of lone pairs attached to the central atom. (Electron domains are lone pairs or
bonding pairs of electrons, but multiple bonds count as one domain).
3. Determine the geometrical arrangement of electron domains (also see Table 1):
2 domains
Linear
3 domains
Trigonal Planar
4 domains
Tetrahedral
5 domains
Trigonal Bipyramidal
6 domains
Octahedral
4. Re-draw the molecule in the correct electron domain geometry place the lone pairs in their proper
positions in order to minimize repulsions (refer to Table 1).
5. Determine the molecular geometry (shape) of the molecule (the geometric arrangement of just the
atoms – ignore the lone pairs )
Table 1.
Total
Electron
Domains
Molecular Geometries (Shapes) Predicted by VSEPR
No Lone
Pairs
One Lone
Pair
Two Lone
Pairs
Three Lone
Pairs
2
Geometry:
Angles:
Linear:
180°:
••
3
Geometry:
Angles:
Trigonal Planar
Bent
120°:
~120°:
••
4
Geometry:
Angles:
Tetrahedral
Trigonal-pyramid
bent
109.5°
~109.5°
~109.5°
Trigonal
Bipyramidal
See-saw
T-shaped
Linear
90°, 120°
90°, 120°
90°
180°
5
6
••
••
Geometry:
Angles:
••
Octahedral
Square-pyramid
Square-planar
90°
90°
90°
•• ••
Angles:
••
•• ••
••
Geometry:
Prediction of Polarities of Molecules from VSEPR and Electronegativities
A polar covalent bond is formed when a pair of electrons is shared between two atoms with different
electronegativities. The atom with the greater electronegativity exerts a greater attraction for the
electrons in the bond than does the atom with the smaller electronegativity. The electron density is
displaced toward the more electronegative atom, with the net effect that the more electronegative atom
has a partial negative character and the less electronegative atom has a partial positive character. The
HCl molecule is an example of a molecule possessing a polar covalent bond:
–
Cl
+
H
The Cl atom has a greater share of the bonding electrons than H atom due to the larger electronegativity
of the Cl atom. The HCl molecule is said to possess a dipole moment due to the separation of the
partial positive and negative charges. The dipole moment (or polarity of the bond) is indicated by an
arrow pointing towards the negative end of the bond ( and molecule in this case since there is only one
bond).
In molecules containing two or more bonds, the overall polarity of the molecule (or lack thereof)
depends on the net sum of the individual bond polarities. In some cases the individual bond polarities
may cancel, giving rise to no net dipole moment and an overall non-polar molecule. In the case of H2O,
the individual O-H bond polarities do not cancel due to the bent shape of H2O, and the result is a polar
molecule.
O
H
O
H
H
bond dipoles
H
molecular diople
In the carbon dioxide (CO2 ) molecule, the linear geometry of the molecule results in the individual bond
polarities exactly canceling one another and the result is that CO2 is a non-polar molecule:
O
C
O
In general, a molecule will be non-polar when there is a symmetrical arrangement of the same type of
bonds, such that the individual bond polarities cancel one another. The five basic electron domain
geometries from VSEPR (linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral) plus
the square planar geometry are symmetrical arrangements. Polar molecules result when there is an
asymmetric arrangement of bonds, such that individual bond polarities do not exactly cancel one
another.
Experimental Procedures
During this lab period you will determine the Lewis Structures molecular geometries, and predicted
polarities for a number of molecular species beginning only with the molecular formula. As well, you
will need to create a 3-dimensional model of the molecule with ~ correct bond angles using a set of
toothpicks and Styrofoam balls.
You may not be as familiar with the geometries that include 5 or 6 electron domains, so be sure to refer
to the table of geometries at the end of the lab. You will be assigned a set of molecular species (see list
below). You are to do the following tasks for each species, and write the answers in the space provided
on the worksheets:
1. Determine the number of valence electrons in the molecule (or polyatomic ion).
2. Determine and draw the Lewis structure for the molecule (or polyatomic ion).
3. Determine the molecular geometry (shape) using VSEPR
4. Construct a model of the molecule using the materials provided
5. Draw a correct 3–D representation of the molecule from your results and the model
6. Determine whether the molecule (or polyatomic ion) is polar or non-polar.
AXE Method
The "AXE method" of electron counting is commonly used when applying the VSEPR theory. The
A represents the central atom and always has an implied subscript one. The X represents the
number of sigma bonds between the central atoms and outside atoms. Multiple covalent bonds
(double, triple, etc.) count as one X. The E represents the number of lone electron pairs
surrounding the central atom. The sum of X and E is known as the steric number.
The number of lone pair electrons can often be determined by multiplying the number of sigma
bonds (X) times 8 and subtracting this from the total number of electrons. The number of lone
pairs is this difference divided by 2. In other words:
E = (# electrons – 8X)/2
Molecular Species List
1. SiCl4
2. PCl3
3. NO34. SF6
5. PBr5
6. IF3
7. XeF5+
8. CO2
9. BrF210. SO2
11. H2O
12. SF4
13. XeF4
Molecular Geometry Worksheet
Name
Section
Lab Partner
Molecule List #
1. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
2. Formula:
Lewis Structure:
Polar or Non-Polar?
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles)
Polar or Non-Polar?
3. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles)
4. Formula:
Lewis Structure:
3–D Drawing (include bond angles)
Polar or Non-Polar?
# of valence electrons =
Molecular Geometry =
AXE = ________________________
Polar or Non-Polar?
5. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
6. Formula:
Lewis Structure:
3–D Drawing (include bond angles):
Polar or Non-Polar?
# of valence electrons =
Molecular Geometry =
AXE = ________________________
Polar or Non-Polar?
7. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
8. Formula:
Lewis Structure:
Polar or Non-Polar?
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
Polar or Non-Polar?
9. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
10. Formula:
Lewis Structure:
Polar or Non-Polar?
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
Polar or Non-Polar?
11. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
12. Formula:
Lewis Structure:
Polar or Non-Polar?
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
Polar or Non-Polar?
13. Formula:
Lewis Structure:
# of valence electrons =
Molecular Geometry =
AXE = ________________________
3–D Drawing (include bond angles):
Polar or Non-Polar?
Linear
2 electron groups, 0 lone pairs
Many linear molecules exist, prominent examples include CO2, HCN, and xenon difluoride. Linear
anions include N3− and SCN−. Linear cations include NO2+.[1]
Trigonal Planar
3 electron groups, 0 lone pairs
Examples of molecules with trigonal planar geometry include boron trifluoride (BF3), formaldehyde
(H2CO), phosgene (COCl2), and sulfur trioxide (SO3). Some ions with trigonal planar geometry include
nitrate (NO3−), carbonate ion (CO32−), and guanidinium C(NH2)3+. In organic chemistry, planar, threeconnected carbon centers that are trigonal planar are often described as having sp2 hybridization.[2][3]
Bent
3 electron groups, 1 lone pair
SO2
Tetrahedral
4 electron groups, 0 lone pairs
side from virtually all saturated organic compounds, most compounds of Si, Ge, and Sn are tetrahedral.
Often tetrahedral molecules feature multiple bonding to the outer ligands, as in xenon tetroxide (XeO4),
the perchlorate ion (ClO4−), the sulfate ion (SO42−), the phosphate ion (PO43−). Thiazyl trifluoride, SNF3
is tetrahedral, featuring a sulfur-to-nitrogen triple bond
Trigonal Pyramidal
4 electron groups, 1 lone pair
Some molecules and ions with trigonal pyramidal geometry are ammonia (NH3), xenon trioxide, XeO3,
the chlorate ion, ClO3−, and the sulfite ion, SO32−.
Bent
4 electron groups, 2 lone pairs
Nonlinear triatomic molecules and ions are common for compounds containing only main group
elements, prominent examples being water, nitrogen dioxide, SCl2, and the CH2
Trigonal bipyriamidal
5 electron groups, 0 lone pairs
PCl5
See saw
5 electron gropus, 1 lone pair
SF4, SeF4, IF2O2-, IOF3, ClF4+, TeF4, XeO2F2, AsF4-
T shaped
5 electron groups, 2 lone pairs
Example of a T-shaped molecules are the halogen trifluorides, such as ClF3
Linear
5 electron groups, 3 lone pairs
XeF2
Octahedral
six electron groups, 0 lone pairs
sulfur hexafluoride SF6 and molybdenum hexacarbonyl Mo(CO)6. The term "octahedral" is used
somewhat loosely by chemists, focusing on the geometry of the bonds to the central atom and not
considering differences among the ligands themselves. For example, [Co(NH3)6]3+, which is not
octahedral in the mathematical sense due to the orientation of the N-H bonds, is referred to as octahedral
Square pyramidal
6 electron groups, 1 lone pair
Some molecular compounds that adopt square pyramidal geometry are XeOF4,[2] and XF5 (X = Cl, Br, I)
also have square pyramidal geometries.[3][4] Complexes of vanadium(IV), such as [VO(acac)2] are
square pyramidal (acac = acetylacetonate, the anion of 2,4-pentanedione that has lost a proton)
Square planar
6 electron groups, 2 lone pairs
The noble gas compound XeF4 adopts this structure as predicted by VSEPR theory. The geometry is
prevalent for transition metal complexes with d8 configuration, which includes Rh(I), Ir(I), Pd(II), Pt(II),
and Au(III). Notable examples include the anticancer drugs cisplatin [PtCl2(NH3)2] and carboplatin