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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 Lewis dot ionic form

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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, CO32-

Total 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 2 3 4 5 6

Drawing − X − − X − X− − X X

= X =

− X ≡ = X

Geometry linear trigonal tetrahedral trigonal octahedral

planar pyramidal

Example CO2 BF3 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 Examples

Hybrid

Orbitals

(central

atom)

Polar?

2 Linear AX2 Linear CO2 sp No

3 Trigonal Planar

(3 electron domains) AX3 Trigonal planar BF3, AlCl3 sp

2 No

AX2E Angular SnCl2 Yes

4 Tetrahedral

(4 electron domains) AX4 Tetrahedral CH4, SiCl4 sp

3 No

AX3E Trigonal pyramidal NH3, PCl3 Yes

AX2E2 Angular (Bent) H2O, SCl2 Yes

5 Trigonal bipyramidal

(5 electron domains) AX5 Trigonal bipyramidal PCl5, AsF5 sp

3d No

AX4E Seesaw SF4 Yes

AX3E2 T-shaped ClF3 Yes

AX2E3 Linear XeF2 No

6 Octahedral

(6 electron domains) AX6 Octahedral SF6

sp3d

2

No

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

Bonding pair

electrons

Lone pair electrons

Lewis Dot Structure

Electron Geometry and Shape

Bond Angle

Hybridization

NH3

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3

1

.. N

H H H

Tetrahedral

Trigonal Pyramidal

107⁰ 107⁰

sp3

CF4

FNO2

10

HCN

IBr2

OH—

H3O+

ClO3—

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POCl3

ClO2—

CCl4

NCl3

12

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+ 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) 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.