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 A elements, the number of valence electrons is the same as the group number in the periodic table.

Element
Na
Cl
Group
IA
VIIA
Valence electrons Lewis dot
1
.
Na
7 ionic form
Na
+
..
: Cl :
.
..
[ : Cl
. .
: ] -
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:
Cl
2
..
..
: Cl – Cl : and HCl
. .
. .
..
H – Cl :
.
.
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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:
O = C = O H
 – C ≡ N :
Carbon dioxide Hydrogen cyanide
H

C ≡ C

Acetylene
H
Rules for drawing Lewis Dot Structures
I. Determine the total number of valence electrons and valence pair electrons 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 are less than the total valence electrons pairs calculated earlier, add a pairs to the central atom to match the total number of valence pairs.
(c) the number of pairs are 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.
Example: Carbonate ion , CO
3
2-
Total valence pair electrons = 1 (C) + 3 (O) + 2 e
-
= 1 (4 e
-
) + 3 ( 6e
-
) + 2 e
-
= 24 e
-
= 12 pair electrons
Drawing:
..
..
O :
: O – C ..
O = C
O :
2-
O
: O – C
:
2-
O
2-
: O – C
O : O : O O :
(13 pe
-
) (12 pe
-
) (12 pe
-
) (12 pe
-
)
The carbonate ion 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 ( BF
3
) or more than eight electrons ( PCl
5
, SF
6
,
XeF
4
, etc). For most of these compounds, the central atom and each outer atom are bonded by single bonds consisting of one pair of electrons. If there are any extra electrons on the central atom, they are grouped as shared pairs on the central atom.
For example:
BF
3
= 1(3e
-
) + 3(7 e
-
) = 24 e
-
= 12 p e
-
PCl
5
= 1 (5 e
-
) + 5 ( 7 e
-
) = 40 e
-
= 20 p e
-
F :
: F – B
F :
( 12 pe
-
)
:
: Cl :
| Cl
Cl P
| Cl
Cl :
( 20 pe
-
)
:
:
Valence Shell Electron Pair Repulsion(VSEPR) Theory
Electron Pair and Molecular Geometry
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, CO
3
2 has three regions, CO
2
has two regions, and PCl
5
has five regions. The following electronic geometries are expected for the following numbers of regions of high electron density. Some examples follow.
Regions 2 3
Drawing − X −
= X =
− X
− X ≡ = X
4
− X−
5
− X
6
X
Geometry linear trigonal tetrahedral trigonal octahedral
planar pyramidal
Example CO
2
BF
3
CCl
4
PCl
5
SF
6
Attach the other atoms to the regions of e density so that these other atoms are as far as possible. The relative positions of the central atom and the other atoms can be described as the molecular geometry
(shape). Some examples follow.
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Electron geometry Molecular
Geometry (shape)
Linear (Symmetrical) 1. Linear
2. Trigonal planar Bent (Unsymmetrical)
3. Trigonal planar
4. Tetrahedral t-planar (Symmetrical)
Bent (Unsymmetrical)
Examples of
compounds
BeCl
SO
SO
H
2
2
3
O t-pyramidal (Unsymmetrical) NH
T d
(Symmetrical)
5. Trigonal bipyramidal Linear (Symmetrical)
CH
XeF
3
4
2 t-shaped (Unsymmetrical) ClF
3
2
Angles
180
120
120
105
107
109
90 o
90 o o o o o o o
, 120 o , 180 o
6. Octahedral
Seesaw (Unsymmetrical) t-bp (Symmetrical)
SCl
PCl
4
5
Sq. planar (Unsymmetrical) XeF
4
Sq. pyramidal (Unsymmetrical) IF
5
Octahedral (Symmetrical) SCl
5
Molecular Polarity (Dipole Moment)
90
90
90
90
90 o o o o o
, 120
, 120
, 180 o o o
, 180
, 180 o o
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 reenforce or oppose the effect of the lone pairs of electrons. Some examples are as follow:
O = C = O
N
|
H H H
Symmetrical (non-polar)
Unsymmetrical (polar)
Molecular Hybridization
Once the Lewis structure is drawn, it is possible to apply either the VSEPR 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.
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Electro Pair and Molecular Geometry total number bond e- pair molecular of e
-
pairs shared e- pairs unshared e- pairs angles geometry geometry examples
2 2 0 180 o linear linear BeCl
2
3
3
3
2
0
1
120
120 o o
t -planar
t -planar t -planar t-planar
BF
SO
3
3
4 0 109 o
tetrahedral tetrahedral CCl
4
4
4 3 1 107 o
tetrahedral t -pyramidal NF
3
4 2 2 105 o tetrahedral bent H
2
O
Obtain a molecular kit and draw a Lewis dot structure for each compound in table 1. From your dot structure, and using the summary table as a guide, fill in all the blanks for each compound. Then build a model of the molecule. Use the model to verify the geometry you predicted. If the model and your prediction do not agree, find out why. DO NOT proceed to the next compound until you thoroughly understand each structure.
Part I - Determine the Lewis dot structure for the following molecules in the table below and predict their molecular geometry. Using a molecular kit, construct the molecules and confirm your prediction. Lastly, use electronegativity to predict whether the molecule is polar or nonpolar.
Part II . For each species listed below, draw the Lewis structure first, and then complete the rest of the information requested for the species. “Build” the molecule or ion with your molecular kit. Verify from your model the geometry you predicted. Complete each line in the table, including the model, before proceeding to the next line. If there are more than two atoms in the “center”, you can reduce the geometry separately for each one. Continue this procedure for each species listed here.
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Part I.
Molecular Lewis formula structure
Molecular
geometry
Polar/
Nonpolar ____
H
2
O
O
H H Bent (unsymmetrical) polar
_____________________________________________________________________________________
CCl
4
_____________________________________________________________________________________
NF
3
_____________________________________________________________________________________
CH
2
O
_____________________________________________________________________________________
HCN
_____________________________________________________________________________________
OF
2
_____________________________________________________________________________________
NH
4
+
_____________________________________________________________________________________
NO
3
-
_____________________________________________________________________________________
CS
2
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Part II
Molecular Lewis e Pair Molecular Bond Polar/
Formula Structure Geometry Geometry Angle Nonpolar_____________
H

CH
4
C T d
H

H
T d
109 o
Non-polar
(Symmetrical)
H
_____________________________________________________________________________________
SO
3
_____________________________________________________________________________________
H
2
O
2
_____________________________________________________________________________________
SO
4
2-
_____________________________________________________________________________________
HCl
_____________________________________________________________________________________
H
3
O +
_____________________________________________________________________________________
NO
2
-
_____________________________________________________________________________________
C
2
H
4
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Due before lab begins. Answer in space provided.
1. Write the Lewis dot structure for the following atoms or ions.
Al Cl K
+
O
2-
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) Two regions of electron density: _________________________
3. Draw Lewis structures for BF
3
and NF
3
. Compare the geometries, angles, and polarities.
Give the electron pair geometry and the molecular geometry of BF
3
and NF
3
according to
VSEPR.
4. Define resonance and give an example.
5. Are BF
3
and NF
3
symmetrical or unsymmetrical molecules? Explain your reasoning.
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Due after completing lab. Answer in space provided.
1. Describe polar and non-polar bond. Give an example of each.
2. Draw Lewis structure for CS
2
, H
2
Se, and ClO
3
-
.
3. Are BF
3
and NF
3
polar or non-polar molecules? Explain your reasoning.
4. Define dipole moment and give an example.
5. Draw the Lewis structures for ethanol, C
2
H
5
OH, and dimethylether(CH
3
OCH
3
). Determine the
electron pair geometry and molecular geometry around the oxygen atom in each.
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