Experiment 12 The Valence-Shell Electron-Pair
Repulsion Theory (VSEPR)
Models of Molecular Structure
Purpose:
The purpose of this experiment is to gain an understanding of the spatial arrangement of
atoms in covalently bonded molecules and how the molecular geometry, along with the
electronegativity of the atoms comprising the molecule (polarity), affects many of its chemical
and physical properties such as reactivity, solubility, melting point, and boiling point.
Background:
Most chemical reactions, especially biochemical processes, depend heavily on the
relationship between chemical functionality (does the reaction actually occur?) and the shapes
(geometry) of the interacting molecules. Unfortunately, molecules and the spatial arrangement of
the atoms in the molecule are far too small to be seen directly. Molecular shapes can be deduced
experimentally using sophisticated methods such as spectroscopy. Another approach, to be used
is this experiment, is to use a model to predict the shapes of molecules.
Figure 1 shows three different representations of the simple molecule methanol, CH4O: a line
drawing (2-demensional), and two 3-dimensional models  a ball-and-stick model and a spacefilling model.
Figure 1: Molecular models
In todays’ lab students will study the geometry of molecules by constructing ball-and-stick
models based on Lewis Dot Structures and the Valence-Shell Electron-Repulsion theory
described below. Model building is an important aspect of science. Often the best way to
understand a complicated, real system is to study an idealization of it. In this case simple, solid
balls will be us as the central atom of the molecule and sticks will be used to represent the
directional covalent bonds. This sort of model is useful, not only to help the student visualize the
three dimensional shapes of simple molecules, but even to the experienced chemist approaching
an unfamiliar, complex molecule.
The first step in the process is to use the molecular formula of a molecule to depict a 2dimensional view of the electron configuration of the bonded and unbonded valence electrons in
the molecule using Lewis structures. The electron configuration, along with an assumption
regarding spatial distribution of the electron groups is then applied to the Valence-Shell
Electron-Pair Repulsion Theory, which provides a mathematical means of predicting the
geometric shape of the molecule as well as its polarity.
Lewis Structures:
A Lewis Structure (or Lewis Dot Diagram, Electron Dot Diagram) is a diagram used to
visualize the bonding between atoms of a molecule and any lone pairs of electrons that may be
needed in the molecule for completing the outer valence shell. A Lewis structure can be drawn
for any covalently bonded molecule, as well as coordination compounds. The bonding electron
pairs can be represented as either lines or dots. Double or triple bonds would show 2 or 3 lines or
2 or 3 pairs of dots. Excess electrons that form lone pairs are represented as pairs of dots and are
placed next to the atoms.
Writing a Lewis Structure involves the following steps:
1. Calculate the number of valence electrons for all atoms in the molecule:
a. take the group number for each atom (1-8) and multiply by number of each atom
b. compute the total number of valence electrons all atoms
c. add the charge if molecule is an anion
d. subtract the charge if molecule is an cation
2. Put the atom with the lowest group number and lowest electronegativity as the central
atom
3. Arrange the other elements (ligands) around the central atom
4. Distribute electrons as pairs (dots) to the atoms surrounding the central atom to
satisfy the octet rule for each atom
5. Distribute the remaining electrons as pairs (dots or lines) to the central atom
6. If the Central atom is deficient in electrons to complete its octet; borrow electron
pairs from surrounding atoms to complete central atom valence electron needs, that is,
form one or more double bonds (possibly triple bonds) around the central atom
Example 1: Dichlorodifluoromethane (CCl2F2)
1. Choose central atom – Carbon (C) has the lowest group number and lowest
electronegativity.
2. Place the other elements (ligands - Cl & F) uniformly around the central atom.
3. Determine total number of valence electrons
1 x C(4) + 2 x Cl(7) + 2 x F(7) = 32
4. Draw single bonds from central atom to each ligand and subtract 2 e for each single
bond
4 x 2 = 8.  32 – 8 = 24 remaining
5. Distribute the remaining 24 electrons to the atoms (ligands) surrounding the central
atom to satisfy the octet rule for each atom
Example 2: Nitrate Ion (NO3-) with double bond formation and computation of net ionic
charge from formal charges
1. Select central atom – Nitrogen (N) has the lowest group number and lowest
electronegativity.
2. Total valence electrons - 3 x 6 (O) + 1 x 5 (N) + 1 (ion charge) = 24
3. Add 1 pair electrons between central atom and each other atom – 3 x 2 = 6
4. Add electrons to oxygen atoms to complete octet
5. Nitrogen still missing 2 electrons to complete octet
6. Borrow 2 electrons from one oxygen to form double bond
7. Formal Charge – nitrogen: 5 – (0 + ½*8) = 5 – 4 = +1
8. Formal Charge – single bonded oxygen: 6 – (6 + ½*2) = 6 – 7 = -1 x 2 = -2
9. Formal Charge – double bonded oxygen: 6 – (4 + ½*4) = 6 – 6 = 0
10. Net Charge of ion is: +1 +(-2) = -1
Formal Charge:
As seen in the above example, a molecule or polyatomic ion can contain one or more double
(even triple) bonds. In these cases, it is possible to write more than one Lewis structure and a
decision must be made as to which structure is preferred. The nitrate ion above is shown in
its 3 “resonance” forms. From an electronegativity perspective they are all the same and any
one is just as preferable as any of the others.
In other molecules, the relative placement of the multiple bonds and electronegativity of the
elements does can in resonance forms where each element will have a different net or formal
charge. Evaluation of the formal charge forms the basis of selecting the preferable resonance
form.
Formal charge is defined as:
Formal Charge = no. of valence e-  (no. unshared e- + ½ no. shared e-)
Criteria: a. Smaller formal charges (Pos or Neg) are preferable to larger ones
b. The same nonzero formal charges on adjacent atoms are not preferable
c. A more negative formal charge should reside on a more electronegative atom
Example: The cyanate ion – CNOThe cyanate ion can be presented in 3 resonance forms:
The formal charges for each atom in each form are calculated as follows:
I
FCN = 5 – (6 + ½*2) = -2
FCC = 4 – (0 + ½*8) = 0
FCO = 6 – (2 + ½*6) = +1
II
FCN = 5 – (4 + ½*4) = -1
FCC = 4 – (0 + ½*8) = 0
FCO = 6 – (4 + ½*4) = 0
III
FCN = 5 – (2 + ½*6) = 0
FCC = 4 – (0 + ½*8) = 0
FCO = 6 – (6 + ½*2) = -1
Preferred Form:
Eliminate I – Higher formal charge on nitrogen than carbon & oxygen
Positive formal charge on oxygen, which is more electronegative than
nitrogen
Eliminate II – Forms II & III have the same magnitude of formal charges, but
form III has a -1 charge on the more electronegative oxygen atom
Forms II & III both contribute to the resonant hybrid of the cyanate ion, but form III
is the more important
Note: Net formal charge in form III is the same as the ionic charge (-1)
Extended Valence Shells:
Many molecules (and ions) have more than eight electrons around the central atom. Nonmetal elements from period 3 or higher can expand their valence shell by utilizing empty “d”
orbitals to hold the additional electron pairs.
Phosphorus pentachloride (PCL5) has the following Lewis structure
1. Select central atom – Phosphorus (P) has the lowest group number and lowest
electronegativity.
2. Total valence electrons - 1 x 5 (P) + 5 x 7 (Cl) = 40
3. Distribute 6 e- to each of the 5 chlorine atoms 5 x 6 =30
4. Establish bonding pairs between phosphorus and chlorine atoms 5 x 2 = 10
5. Remaining electrons: 40 – 30 - 10 = 0
6. Phosphorus has “zero’ non-bonding pairs
7. Phosphorus has 10 valence shell electrons, thus it violates the “octet” rule
8. Since phosphorus is in Period 3, PF5 is a “hypervalent” molecule
9. The phosphorus utilizes electrons from vacant “d” orbitals to create a valence
shell with more than 8 electrons
Valence-Shell Electron-Repulsion Theory:
Lewis structures provide 2-dimensional views of molecules, but since molecules are 3dimensional, the Lewis views are inherently misleading. The Valence-Shell Electron-Repulsion
Theory (VSEPR) is used to predict the shape of a molecule from knowledge of its chemical
formula and its electron distribution as depicted in the Lewis structure. The distribution of the
atoms in 3-dimensional space is based on the principle of minimizing the electron repulsion
between the valence electrons by arranging each group of valence electrons around the central
atom as far as possible from the other groups. Groups of electrons include single bonds, double
bonds, triple bonds, unbonded pairs and even lone electrons. These arrangements give rise to the
distinctive molecular shapes (molecular geometry) that depend solely on the positions of the
atomic nuclei relative to each other, i.e. the angular arrangement of the electron pairs around the
central atom. The electron-group arrangement is defined by the bonding and non-bonding
electron groups, but the molecular shape is defined by the spatial arrangement of the groups
around the central atom.
The VSEPR process (includes the Lewis Dot process):
1. Central Atom
1. Place atom with Lower Group Number (least electronegativity) in center
2. If atoms have the same group number, place the atom with the “Higher Period
Number” in the center
Example: Sulfur (S) for SO3; Chlorine (Cl) for ClF3
2. Draw the Lewis structure
a. Determine how many bonding electron pairs are around the central atom.
b. Determine the number of non-bonding electron pairs
c. Count a multiple bond as “one pair”
d. Arrange the electron pairs as far apart as possible to minimize electron repulsions
e. Note the number of bonding and lone pairs
3. VSEPR Molecular Notation:
AXaEb
A – The Central Atom (Least Electronegative atom)
X – The Ligands (Bonding Pairs)
a
– The Number of Ligands
E – Non-Bonding Electron Pairs
b
– The Number of Non-Bonding Electron Pairs
Double & Triple Bonds count as a “single” electron bonding pair
The geometric arrangement is determined by:
sum (a + b)
Group Arrangements and Molecular Shapes:
Two to six maximally spaced ligands can be attached to a central atom resulting in five
geometric patterns. With the exception of the Linear pattern, each arrangement consists of two,
three, or four shapes representing the unique combination of bonded electron pairs and nonbonding electron pairs. The possible geometric arrangements of the electron pairs around the
central atom and the number of electron pairs required to give rise to a particular arrangement are
summarized in tables 2 and 3.
Linear:
The linear arrangement consists of two ligands attached by two pairs of bonding electrons
to a central atom in opposite directions. This arrangement does not have any unbonded
electron pairs. The shape is linear with a bond angle of 180o.
Examples: CO2, CS2, HCN, BeF2, NO2+
Central atom is carbon because it has lower group number
Double bond counts as 1 bonding pair
Carbon has 2 bonding e- pairs and 0 non-bonding e- pairs
AX2E0
a + b = 2 + 0 = 2  Linear
Trigonal Planar:
The Trigonal Planar arrangement consists of up to three ligands attached to a central atom
comprising two molecular shapes:
1. Trigonal Planar  Central atom bound to three ligands (3 bonding electron pairs) and
no non-bonding pairs with bond angles of 120o
Examples: COCl2, SO3, BF3, NO3-, NO2-, CO32Central Atom is carbon with 3 bonding electron pairs (double
bond counts as 1 bonding e- pair) and 0 non-bonding e- pairs.
AX3E0 a + b = 3 + 0 = 3 Trigonal Planar
2. Bent V-shaped  Central atom bound to two ligands (2 bonding electron pairs) plus
one non-bonding electron pair with bond angles of <120o.
Shows the effects of one pairs and double bonds on bond angle.
Examples: SO2, O3, PbCl2, SnBr2
Ozone (O3) has two bonding electron pairs about the central
oxygen (double bond counts as 1 bonding e- pair) and 1 nonbonding e- pair.
AX2E1 a + b = 2 + 1 = 3  Bent (V-shaped)
Tetrahedral:
The Tetrahedral arrangement consists of up to 4 ligands attached to a central atom
comprising three molecular shapes:
1. Tetrahedral  Central atom bound to four ligands (4 bonding e- pairs) with bond
angles of 109.5o.
Examples: CH4, CCl4, SiCl4, SO42-, ClO4Central atom is carbon because it has lower electronegativity.
Carbon has 4 bonding e- pairs and 0 non-bonding e- pairs.
AX4E0 a+b = 4 + 0 = 4  Tetrahedral
2. Trigonal Pyramidal  Central atom bound to 3 ligands (3 bonding electron pairs) plus
one non-bonding electron pairs with bond angles of <109.5o.
Examples: NH3, PF3, CLO3, H3O+
Central atom is nitrogen because it has lower group number.
-
-
Nitrogen has 3 bonding e pairs and 1 non-bonding e pair
AX3E1 a + b = 3 + 1 = 4  Trigonal Pyramidal
3. Bent (V-Shaped)  Central atom bound to two ligands (2 bonding electron pairs)
plus 2 non-bonding electron pairs with bond angles of <109.5o.
Examples: H2O¸ OF2, SCL2
Central atom is oxygen because it has lower group number.
Oxygen has 2 bonding e pairs and 2 non-bonding e pairs.
AX2E2 (a+b) = 2 + 2 = 4  Bent V-Shaped
Trigonal Bipyramidal:
The Trigonal Bipyramidal arrangement consists of up to 5 ligands attached to a central
atom comprising four molecular shapes. Note that trigonal bipyramidal geometry has two
angle sets (axial and equatorial).
1. Trigonal Bipyramidal  Central atom bound to five ligands (5 bonding e- pairs)
with bond angles of 90o and 120o.
Examples: PF5, ASF5, SOF4
Central atom is phosphorus because it has lower electronegativity.
Phosphorus has 5 bonding e- pairs and 0 non-bonding e- pairs.
AX5E0 a + b = 5 + 0 = 5  Trigonal Bipyramidal
2. SeeSaw  Central atom bound to four ligands (4 bounding electron pairs) and one
non-bonding electron pair with bond angles of <90o and <120o.
Examples: SF4, XeO2F2, IF4+, IO2F2The central atom is sulfur because it has lower electronegativity.
Sulfur has 4 bonding e pairs and 1 non-boning e pair.
AX4E1 (a + b) = 4 + 1 = 5  SeeSaw
3. T-shaped  Central atom bound to 3 ligands (3 bonding electron pairs) and two nonbonding electron pairs with bond angles of <90o.
Examples: ClF3, BrF3
The central atom is chlorine because it has lower electronegativity.
Chlorine has 3 bonding e- pairs and 2 non-bonding e- pairs.
AX3E2
(a + b) = 3 + 2 = 5  T-Shaped
4. Linear  Central atom bound to 2 ligands (2 bonding electron pairs) and
3 non-bonding electron pairs with bond angles of 180o.
Examples: XeF2, I3-, IF2
The central atom is Xenon because it has lower electronegativity.
Xenon has 2 bonding e- pairs and 3 non-bonding e- pairs.
AX2E3 (a + b) = 2 + 3 = 5  Linear
Octahedral:
The Octahedral arrangement consists of up to 6 ligands attached to a central atom
comprising three molecular shapes:
1. Octahedral  Central atom bound to six ligands (6 bonding electron pairs) and no
non-bonding pairs with bond angles of 90o.
Examples: SF6, IOF5
Central atom is sulfur because it has lower electronegativity.
Sulfur has 6 bonding e- pairs and 0 non-bonding e- pairs.
AX6E0 (a + b) = 6 + 0 = 6  Octahedral
2. Square Pyramidal  Central atom bound to five ligands (5 bonding electron pairs) and
one non-bonding pair with bond angles of <90o.
Examples: BrF5, TeF5-, XeOF4
Central atom is bromine because it has lower electronegativity.
Bromine has 5 bonding e- pairs and 1 non-bonding e- pairs.
AX5E1 (a + b) = 5 + 1 = 6  Square Pyramidal
3. Square Planar  Central atom is Xenon because it has lower electronegativity.
Xenon is bound to four ligands (4 bonding electron pairs) and
two non-bonding pairs with bond angles of <90o.
Example: XeF4
Central atom is Xenon because it has lower electronegativity.
Xenon has 4 bonding e- pairs and 2 non-bonding e- pairs.
AX4E2 (a + b) = 4 + 2 = 6  Square Planar
Table 1: VSEPR Arrangements
Total # of
Electron Pairs
On Central
Atom (a + b)
Valence-Shell Electron-Repulsion (VSEPR)
Geometric Arrangement of Electron Pairs around the Central Atom
AXaEb
2
Linear
3
Trigonal Planar
4
Tetrahedral
5
Trigonal
Bipyramidal
6
Octahedral
Table 2: VSEPR Arrangements and Shapes
Valence-Shell Electron-Repulsion Theory (VSEPR)
Molecular Shapes within Electron Group Arrangement
Total #
e- pairs
Bonding
e- pairs
NonBonding
e- pairs
Group
(Shape)
AXaYb
Designation
Angles
2
2
0
Linear
AX2
180o
3
3
0
Trigonal
Planar
AX3
120o
AX2E1
<120o
3
2
1
Trigonal
Planar
(Bent
V-Shaped)
4
4
0
Tetrahedral
AX4
109.5o
1
Tetrahedral
(Trigonal
Pyramidal)
AX3E1
<109.5o
4
3
<109.5o
Tetrahedral
4
2
2
(Bent
V-Shaped)
AX2E2
Table 2 (con’t):
Valence-Shell Electron-Repulsion (VSEPR)
Molecular Shapes within Electron Group Arrangement
Total
e pairs
5
Bonding
e pairs
5
NonBonding
e pairs
Group
(Shape)
0
Trigonal
Bypyramidal
AXaYb
Designation
Angles
AX5
120o
90
AX4E1
<120
<90
o
5
4
1
Trigonal
Bypyramidal
(Seesaw)
5
3
2
Trigonal
Bypyramidal
(T-Shaped)
AX3E2
<90
5
2
3
Trigonal
Bypyramidal
(Linear)
AX2E3
180
6
6
0
Octahedral
AX4
90
6
5
1
Octahedral
(Square
Pyamidal)
AX5E1
<90
2
Octahedral
(Square
Planar)
AX4E2
<90
6
4
o
o
o
o
o
o
Isomers:
Isomers are compounds that have the same molecular formula, but different arrangements of
the atoms and functional groups in space. There are several different classes of isomers that
determine whether they have similar or dissimilar physical and chemical properties. Isomers are
characterized as to whether they are chiral or achiral; mirror images of each other;
superimposable on their mirror images; and their ability to rotate a plane of polarized light.
Figure 2 shows the subdivision of the isomer groups.
Figure 2: Isomer Categories.
ISOMERS
Different compounds with the same molecular formula
Constitutional (Structural) Isomers
Isomers whose atoms have a different
. arrangement (connectivity) of atoms
 Chain (Skeletal) isomerism
 Position isomerism
(Regioisomerism)
 Functional group isomerism








Chiral Molecules
Not identical (asymmetric) on both
sides of a central plane
Not superposable on mirror image
Chiral center with four different
substituents
Generally similar properties
Optically active
Enantiomers (optical)
Mirror images of each other
Not superposable on each other.
Identical physical & chemical
properties
Stereoisomers
Structural isomers) that differ only in
the three-dimensional orientations of
their atoms (functional groups) in space.


Achiral Molecules
Identical (symmetrical) on both sides of
a central plane of symmetry)
Identical to mirror image (superposable)
Optically inactive




Diastereomers
Not an enantiomer
Not mirror images of each other.
Not Superposable
Different chemical & physical properties


Geometric (Cis/Trans)
Differ arrangement of groups at a single
atom, at double bonds, or in rings
Constitutional (structural) isomers:
Constitutional isomers have the same molecular formula, but having different arrangements
of the atoms in space. They usually do not share similar properties. Constitutional isomers can be
subdivided into three categories:
1. Chain (Skeletal) Isomers  Same molecular formula but the hydrocarbon chain has
variable amounts of branching.
2. Position Isomers  Same molecular formula but a functional group or
other substituent changes position on a parent structure.
3. Functional Group Isomers  The same number atoms of the same elements, but the atoms
are connected in different ways so that the groupings form different functional groups.
Stereoisomers (or Stereomers):
Stereoisomers are structural isomers of each other with the same molecular formula and
sequence (connectivity) of the bonded atoms, but they differ in the three-dimensional
orientations of their atoms in space.
Chirality – Chirality is a property of handedness, whether or not a stereoisomer can be
superimposed on its mirror image. The left hand cannot be superimposed on the right hand.
Images of the left and right hands do not have a plane of symmetry, thus their mirror images are
not identical. Such molecules are referred to as being “chiral.” “Achiral” molecules do have a
plane of symmetry and their mirror images are identical (superposable). Examples of chiral and
achiral molecules are presented in figure 3, which substitute simple skeletal diagrams for Wedge
(
) & Dash (
) diagrams as an aid in visualizing the 3-dimensional forms of the molecules.
Atoms with the wedge bonds are in front of the plane of paper and atoms with dashed bonds are
behind the plane of paper. The student is encouraged to construct ball and stick models of the
examples below as an even better method of visualizing the chirality, mirror images, superimposability of the molecules.
Figure 3: Wedge & Dash diagrams for chiral and achiral molecules
Chiral stereoisomers can be subdivided into two general categories:
Enantiomers (also called optical isomers)
Molecules that differ three-dimensionally by the placement of substituents around one or
more atoms in a molecule. They are asymmetric molecules that are mirror images of each
other but cannot be superimposed on each other. The molecules are not necessarily locked
into their positions, but cannot be converted into one another, even by a rotation around a
single bond. They are called optical isomers because they rotate the plane-polarized light.
Figure 4: Enantiomer (optical) stereoisomer
Note the asymmetry of the central carbon atom in the example in figure 4 as it is attached
to 4 different groups. This carbon is a “stereocenter” and in organic chemistry is referred to
as a “chiral center.” The exchange of any two of the groups results in a stereoisomer.
Diastereoisomers
Stereoisomers that have the same molecular formula and the same bonding or
connectivity of the atoms, but they are not mirror images of each other and have different
physical properties and chemical reactivity.
Figure 5 shows how the same molecular formula can represent both an enantiomer
isomeric form or a disastereoisomer form. The difference is the relative positioning of the
chlorine, hydroxyl and proton groups, either behind (
) or in front (
) of the plane of
the paper. The student should construct ball and stick models of the A, B, and C molecules to
better visualize why molecule B is a mirror image of A and not superimposable on A, while
molecule C is not a mirror image of A and is not superposable on A.
Figure 5: Diastereoisomers vs Enantiomers:
Cis-Trans Stereoisomers
Included in the disastereoisomer group are the cis-trans isomers (also known as geometric
isomers) where rotation of a functional group around a chemical bond is impossible such as
in the case of double and triple bonds and in some ring structures. Cis-trans isomers occur
both in organic molecules and in inorganic coordination complexes.
Figure 6 shows the cis and trans isomers of 2,3-dichloro-2-butene and 1,2dimethylcyclopentane. In neither case, can the chlorine or methyl groups be interconverted
because of the double bond and ring structure, respectively. They are also not mirror images
of each other. Thus, the cis and trans forms of each compound are diastereoisomers, not
enantiomers.
Figure 6: Cis-Trans isomers:
The Experiment:
Gain familiarity with VSEPR theory and develop an ability to visualize the three-dimensional
structure of molecules using a ball and stick modeling kit.
Materials & Equipment:
Equipment:
Molecular Model Set
Molecular Model Set - Organic
Procedure:
Part A: 5 Basic VSEPR Group Arrangements
1. Obtain a ball & stick model set from the instructor and examine the pieces included.
Notice that there are plastic balls with varying number of holes depending on the color
of the ball to be used as either the central atom or ligands
There are also sticks representing bonding or non-bonding electron pairs and several
flexible plastic bonds to be used to connect the central atom to the ligands.
2. Construct three-dimensional models of the five basic VSEPR group arrangements:
Linear, Trigonal Planar, Tetrahedral, Trigonal Bipyramidal, and Octahedral.
3. Use the included protractor to confirm that the various angles are as indicated in
Table 1.
4. The laboratory report should include a sketch of each model and a description of the
process you used to match the model with the specified molecular arrangement.
5. For the trigonal bypyramidal, differentiate between the axial and equatorial positions
and relevant angles.
6. For the octahedral, differentiate between the two sets of angles.
7. Record the sketches and applicable angles used for each model in the procedure results
block.
8. When the 5 models have been completed, have the instructor check them.
Part B: VSEPR Arrangements of select compounds
1. Develop the Lewis Dot structures and VSEPR models for the molecules listed below:
BeH2
BH3
CH2Li2
H2O
BiF5
SF5Cl
For each case:
a. Calculate the total number of valence electrons in the molecule
b. Draw a Lewis Dot structure for the molecule following the instructors provided
above.
c. Set up the AXaEb notation for the molecule
d. Determine the values of “’a” and “b”
e. State the geometric arrangement of electron pairs corresponding to (a + b)
f. Build the molecular model using different colored balls to represent different
atoms or ligands.
g. Represent the single bonds by using a single connecting rod between the ligand
and the central atom.
h. Use a rod, but no ball to represent non-bonded electron pairs on the central atom.
i. Draw the molecule in its correct geometry using the figures in table 2 as a guide.
j. Explicitly write the symbol for each atom or non-bonding pair of electrons in
you drawing.
Part C: Investigate geometrical isomerism
Build ball and stick models of the indicated compounds in the following exercises. In
your lab report, describe the isomer properties of each pair. Depict their 2-dimensional
structure using simple skeletal diagrams and then their 3-dimensional structures using
Wedge & Dash diagrams.
1. Build molecular models of the two possible isomers of PCl4Br. Which atom should be
used as the central atom? Why are only two isomers possible? Are the two molecules
superimposable? Are they the same or different compounds?
2. Build two models of the alanine molecule (CH3CH(NH2)COOH) showing that one is
the enantiomer isomer of the other, i.e., they are mirror images of each other and not
superimposable.
3. Build two models of the triose molecule (CHOCOH(H)CH(OH)CH2O) showing that
one is the diastereoisomer of the other, i.e. they are not mirror images of each but are
superposible on each other.