EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory
Materials:
Molecular Model Kit
INTRODUCTION
Although it has recently become possible to image molecules and even atoms using a high-resolution
microscope, most of our information about molecular structure comes from often this information enables us to
piece together a 3-dimensional picture of the molecule. On paper, one of the best methods we have of
representing this model is by drawing a Lewis Dot Structure of the molecule or ion. The ability to draw Lewis
structures for covalently bonded compounds and polyatomic ions is essential for understanding of polarity,
resonance structures, chemical reactivity, and isomerism. Molecular molecules are a useful tool to help you
visualize structures, especially when the ion or molecule is not planar. It is not the intention here to teach you all
aspects of drawing Lewis structures or constructing molecular models. The goal here is to supplement your
textbook by guiding you through some examples and providing a few insights and tactics where difficulty is
often encountered. Be sure to refer to your chemistry textbook for more complete instruction and details. You
will use a molecular model kit to construct molecules as they are discussed in this exercise. For each model,
you will draw a Lewis dot structure, including nonbonding electrons.
The Lewis dot structure is a two-dimensional representation that shows the arrangement of atoms in a molecule.
The Lewis dot structure includes both bonding and nonbonding electrons. When drawing covalent molecules,
remember that the electrons are shared between two atoms, forming a covalent bond.
In this experiment you will be asked to determine the polarity of certain bonds and, after considering the
geometry of the molecule, decide whether it is polar or nonpolar.
Lewis Diagram (Electron Dot)
In most stable molecules or polyatomic ions, each atom tends to acquire a noble-gas structure by sharing
electrons. This tendency is often referred to as the octet rule. One way to show the structure of an atom or a
molecule is using dots to represent the outermost s-and p-electrons (the so-called valence electrons). For the
group IA elements, the number of valence electrons is the same as the group number in the periodic table.
Element
Group
Valence electrons
1
Lewis dot
ionic form
In drawing Lewis diagrams we usually do not attempt to show from which atom the valence electrons come we simply
indicate a shared pair of electrons as either two dots or a straight line connecting the atoms. Unshared pairs, also called
lone pairs, are indicated by dots drawn around the elemental symbols. For example:
Atoms in polyatomic ions are held together by covalent bonds. In the ions we will consider in this assignment, all atoms
have a noble-gas structure. Occasionally there are too few electrons available in a species to allow an octet to exist around
each atom with only a single bond. In this instance, multiple (double or triple) bonds will form. For example:
Carbon dioxide
Hydrogen cyanide
Acetylene
Rules for drawing Lewis Dot Structures
I.
Determine the total number of valence electrons and valence electron pairs in all the atoms in the
molecule. For polyatomic ions, add or subtract electrons to arrive at the appropriate charge. For
example, if the charge is( –1), you need to add one electron, and if the charge is( +1), you need to subtract
one electron.
II. Connect the central atom to the surrounding atoms with a single bond. Each bond represents two
electrons.
III. Place electrons about the outer atoms so each (except hydrogen with two electrons) has an octet.
IV. Count the total pairs of electrons; if:
(a) the number of pairs matches the total number of valence electron pairs calculated earlier; this is the
correct Lewis dot structure.
(b) the number of pairs is less than the total valence electron pairs calculated earlier; add a pairs to the
central atom to match the total number of valence pairs.
(c) the number of pairs is more than the total valence electron pairs calculated earlier, form one double
bond (one extra pair) or one triple bond (two extra pairs) around central atom. In the case of a triple
bond, you can place two double bonds instead.
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Example 1 - Methane CH4
1. This compound is covalent. Determine the total number of valence electrons available. One carbon has
4 valence electrons, four hydrogen, each with one valence electron, totals 4. This means there are 8
valence electrons, making 4 pairs, available.
2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer
atoms. Hydrogen is never the central atom. Determine a provisional electron distribution by
arranging the electron pairs (E.P.) until all available pairs have been distributed. The formal charge (F)
on the central atom is zero.
Example 2 - Ammonia NH3
1. This compound is covalent. Determine the total number of valence electrons available. One nitrogen
has 5 valence electrons. Three hydrogen, each with one valence electron, totals 3. This means there
are 8 valence electrons, making 4 pairs, available.
2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer
atoms. Hydrogen is never the central atom. It does not matter which of the three sides you use to put
hydrogens on. Determine a provisional electron distribution by arranging the electron pairs (E.P.)
until all available pairs have been distributed. The formal charge (F) on the central atom is zero.
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Example 3 - Water H2O
1. This compound is covalent. Determine the total number of valence electrons available. One oxygen has
6 valence electrons. Two hydrogen, each with one valence electron, totals 2. This means there are 8
valence electrons, making 4 pairs, available.
2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer
atoms. Hydrogen is never the central atom. It does matter which of the two sides you use to put
hydrogens on. Use sides that are next to each other. DO NOT put the hydrogens 180 degrees apart.
Determine a provisional electron distribution by arranging the electron pairs (E.P.) until all available
pairs have been distributed. The formal charge (F) on the central atom is zero.
Example 4 – Carbonate ion, CO32Total valence pair electrons = 1 (C) + 3 (O) + 2 e- = 1 (4 e-) + 3 ( 6e-) + 2 e- = 24 e- = 12 electron pairs
Drawing:
(13 E.P.)
(12 E.P.)
(12 E.P.)
(12 E.P.)
The carbonate ion, CO3 2- has three resonance structures. Delocalization of pi electrons forms resonance
structures.
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Exceptions to Octet Rule and Lewis Dot Structure
There are compounds that cannot be represented by these rules (octet rules) for Lewis dot structures. The central
atom may have either less than eight electrons ( BF3) or more than eight electrons ( PCl5, SF6, XeF4, etc). For
most of these compounds, the central atom and each outer atom are bonded by single bonds consisting of one
pair of electron. If there are any extra electrons on the central atom, they are grouped as shared pairs on the
central atom. For example:
BF3 = 1(3e-) + 3(7 e-) = 24 e- = 12 p e-
PCl5 = 1 (5 e- ) + 5 ( 7 e- ) = 40 e- = 20 p e-
( 12 p e-)
( 20 p e-)
Valence Shell Electron Pair Repulsion(VSEPR) Theory
(Electron Pair and Molecular Geometry)
VSEPR stands for Valence Shell Electron Pair Repulsion. The whole concept revolves around the idea that the
electrons in a molecule repel each other and will try and get as far away from each other as possible. VSEPR
explains a lot about molecular geometry and structure. The electrons (both in pairs and singles as you will see)
are "attached" to a central atom in the molecule and can "pivot" freely on the atom's surface to move away from
the other electrons.
Electrons will come in several flavors:
a) bonding pairs - this set of two electrons is involved in a bond, so we will write the two dots
BETWEEN two atoms. This applies to single, double, and triple bonds.
b) nonbonding pairs - this should be rather obvious.
c) single electrons - in almost every case, this single electron will be nonbonding.
Almost 100% of the examples will involve pairs, but there are a significant number of examples that
involve a lone electron.
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VSEPR uses a set of letters to represent general formulas of compounds. These are:
a) A - this is the central atom of the molecule (or portion of a large molecule being focused on).
b) X - this letter represents the atoms attached to the central atom. No distinction is made between atoms of
different elements. For example, AX4 can refer to CH4 or to CCl4.
c) Eand e - this stands for nonbonding electron pairs.
Each area where electrons exist is called an "electron domain" or simply "domain." It does not matter how
many electrons are present, from one to six, it is still just one domain. Now a domain with six electrons in it (a
triple bond) is bigger (and more repulsive) than a lone-electron domain. However, it is still just one domain.
The more electrons in a domain, the more repulsive it is and it will push other domains farther away than if all
domains were equal in strength. Keep in mind that the domains are all attached to the central atom and will
pivot so as to maximize the distance between domains.
An important point to mention in this introduction is that an element's electronegativity will play an important
role is determining its role in the molecule.
For example, the least electronegative element will be the central atom in the molecule. The more
electronegative the element, the more attractive it is to its bonding electrons This will play a very important
role, especially in five domains.
The most important domain numbers at the introductory level are 3, 4, 5, and 6. Domains of 1 and 2 exist, but
are simple to figure out.
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The VSEPR theory states that regions of high electron density will arrange themselves as far apart as possible
around the central atom. The regions of high electron density are counted for each single atom. In the following
examples, CO32- has three regions, CO2 has two regions, and PCl5 has five regions.
The following electronic geometries are expected for the following numbers of regions of high electron density.
Some examples follow.
Regions
Drawing
Geometry
2
3
−X−
=X=
−X
−X≡
=X
linear
trigonal
4
− X−
tetrahedral
planar
Example
CO2
BF3
5
6
−X
X
trigonal
octahedral
pyramidal
CCl4
PCl5
SF6
Molecular Polarity (Dipole Moment)
If its molecular geometry is completely symmetrical, a molecule is nonpolar. If the molecular geometry is
unsymmetrical, the molecule is polar because of the lone pair of electrons on the central atom. Polar bonds (due
to differences in the electronegativities) may reinforce or oppose the effect of the lone pairs of electrons. Some
examples are as follow:
Symmetrical(polar)
Unsymmetrical (non-polar)
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Molecular Hybridization
Once the Lewis structure is drawn, it is possible to apply either theVSEPR theory or a hybridization to predict
the shape of the species. Both systems when appropriately used will predict (with very few exceptions the same
shape. The table below summarizes the theories.
Electron Pair and Molecular Geometry
Table of Three to Six Electron Domains
Electron
Domains
Arrangement of Electron
Domains
General
Molecular
Formula Molecular Shape
Hybrid
Orbitals
Examples
Polar?
(central
atom)
2
Linear
AX2
Linear
CO2
sp
No
3
Trigonal Planar
(3 electron domains)
AX3
Trigonal planar
BF3, AlCl3
sp2
No
Angular
SnCl2
Tetrahedral
CH4, SiCl4
AX3E
Trigonal pyramidal
NH3, PCl3
Yes
AX2E2
Angular (Bent)
H2O, SCl2
Yes
Trigonal bipyramidal
PCl5, AsF5
AX4E
Seesaw
SF4
Yes
AX3E2
T-shaped
ClF3
Yes
AX2E3
Linear
XeF2
No
Octahedral
SF6
AX2E
4
5
6
Tetrahedral
(4 electron domains)
Trigonal bipyramidal
(5 electron domains)
Octahedral
(6 electron domains)
AX4
AX5
AX6
Yes
sp3
sp3d
No
No
No
sp3d2
AX5E
Square pyramidal
BrF5
Yes
AX4E2
Square planar
XeF4
No
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EXPERIMENT 17: Lewis Dot Structure / VSEPR Theory
REPORT FORM
Name ___________________________
Instructor ________________________
Date ____________________________
Partner’s Name: ___________________
Formula
Total
Valence
electrons
NH3
8
Bonding
pair
electrons
3
Lone pair
electrons
Lewis Dot
Structure
H
H
H
CF4
FNO2
9
Hybridization
Tetrahedral
..
N
1
Electron
Bond
Geometry Angle
and Shape
Trigonal
Pyramidal
107⁰
sp3107⁰
HCN
IBr2
OH—
H3O+
ClO3—
10
POCl3
ClO2—
CCl4
NCl3
11
SiF4
CO32-
NH4+
I3-
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EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory
Name: ___________________________
Pre- laboratory Questions and Exercises
Due before lab begins. Answer in space provided.
1. Write the Lewis dot structure for the following atoms or ions.
Al
Cl
K+
O2-
Si
2. What is the electronic geometry about a central atom which has the following number of regions of
electron density?
a) Three regions of electron density: _________________________
b) Five regions of electron density: _________________________
3. Draw Lewis structures for CO2 and SO2. Compare the geometries, angles, and polarities.
4. Define resonance and give an example.
5. Are BF3 and NF3 symmetrical or unsymmetrical molecules? Explain your reasoning.
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EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory
Name: ___________________________
Post- laboratory Questions and Exercises
Due after completing lab. Answer in space provided.
1. Describe polar and non-polar bonding. Give an example of each.
2. Draw Lewis structure for SO3 and SO32 –.
3. Are CH3 + and CH3 – polar or non-polar molecules? Explain your reasoning.
4. Define dipole moment and give an example.
5. Draw the Lewis structures for ethanol, (C2H5OH), and dimethylether (CH3OCH3 ). Determine the
electron pair geometry and molecular geometry around the oxygen atom in each.
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