Valence Shell Electron Pair Repulsion Theory Lab VSEPR

Student Kit Materials:

Lab Kit:

Item	Item #	Price
Molymod® Individual Organic Pack	840172	13.00

Student Responsibility:

None

EXPERIMENT 2

 VSEPR- Structure and Shape 

Materials and Equipment: Molecular Model Kit.

Introduction

This exercise provides procedures to determine the structure and shape of molecules.  This information is important because the properties of molecules are dependent upon their structure.  The first step in determining the structure (Lewis structure) of a molecule is to draw a structure accurately showing the location of all valence electrons.  From the Lewis structure, you can use a method called valence shell electron pair repulsion (VSEPR) theory to predict the shape of a molecule or ion.

In order to use VSEPR, you need to be able to determine the number of electron groups bonded to the central atom and the number of atoms bonded to the central atom.  Lastly, after you have determined the molecule’s shape, you can determine whether electron density in the molecule is arranged symmetrically (a nonpolar molecule) or asymmetrically (a polar molecule).  In a polar molecule, one end of the molecule has a partial positive charge, one end has a partial negative charge. The polarity of a molecule has important implications for the properties of molecules.

 

Theory of VSEPR

In order to use VSEPR, it is necessary to have a completed Lewis structure for the molecule.  VSEPR is based on the principle that electron groups in a molecule tend to stay as far apart from each other as possible due to the repulsive forces that exist between like charges (the electrons).  An electron group could be a lone pair of electrons, a single bond, a double bond or a triple bond around the central atom.  The most probable arrangement of two, three, or four electron groups around a central atom are given in the table below.  This arrangement allows groups to spread out as far as possible.

Table 1.  Electron Group Geometries

 

# of Electron Groups	Electron Group Geometry
2	Linear
3	trigonal planar
4	Tetrahedral

 

As an example, let’s consider methane, CH4.  The Lewis structure for methane is given below:

 

 

 

 

In this case we can see that there are four electron groups (4 single bonds) surrounding the carbon atom, hence the geometric arrangement of the electrons about the carbon atom is tetrahedral.

Ammonia, NH₃, is a little more difficult.

 

 

 

 

The Lewis structure for ammonia shows that there are four electron groups (3 single bonds and 1 lone pair of electrons) therefore the electron group geometry is also tetrahedral.  It should be noted however that CH₄ has 4 atoms bonded to the central atom, while NH₃ only has 3 atoms bonded to the central atom.  Ammonia therefore they will not have the same shape as CH₄.  Molecular shape describes the arrangement of atoms about the central atom.  When determining the molecular shape, you must consider the electron group geometry and the number of atoms bonded to the central atom.  (Lone pairs are ignored at this point.)  The possible combinations of electron groups and bonded atoms are summarized below.

 

Table 2.  Electron Group Geometries and Molecular Shapes

 

# of Electron Groups	# of Bonded Atoms	Electron Group Geometry	Molecular Shape
2	2	linear	linear
3	2	trigonal planar	bent  (120°)
3	3	trigonal planar	trigonal planar
4	2	tetrahedral	bent   (109.5°)
4	3	tetrahedral	trigonal pyramidal
4	4	tetrahedral	tetrahedral

 

Using Table 2, we can predict that CH₄ has a tetrahedral molecular shape while NH₃ has a trigonal pyramidal molecular shape.

 

After the geometries have been assigned to a molecule, we decide if there is more than one correct structure for it.  These correct structures are called resonance structures.  Lastly, we can use the molecular shape to determine if electron density is evenly distributed across the molecule.  If electron density is unevenly distributed across the molecule, the molecule is said to be polar.  A molecule with a uniform charge distribution is nonpolar.  But first you must learn how to draw Lewis dot structures…

 

Procedure

1. Drawing Lewis structures. This procedure will be illustrated using SO₂ as an example.
2. Determine the total number of valence electrons in the molecule. The number of valence electrons from an atom can be calculated by its location in the periodic table.  So, in this case, S and O are both in group VIA, so each atom contributes 6 electrons.  Hence the total number of valence electrons in SO₂ is 18 (3 atoms ´ 6 valence electrons).

For ions, it is necessary to add or subtract electrons depending on the charge of the ion.  For anions, the magnitude of the charge should be added as additional valence electrons.  For example, for OH^– , the total number of valence electrons is eight:  six from oxygen, one from hydrogen, and 1 for the negative charge  (6 + 1 + 1 = 8).  For cations, the magnitude of the charge should be subtracted from the number of valence electrons.  For NO⁺, the total number of electrons is 10 (5 for nitrogen plus 6 for oxygen, and subtract one for the charge).

2. Determine which atom is the central atom and place a pair of electrons between it and the other atoms. Generally, look for the atom that there is only one of in the formula.  In SO₂, there is only 1 sulfur atom (and 2 oxygen atoms) therefore sulfur is the central atom. Knowing this, we can construct the following crude sketch:

 

 

3. Subtract the number of electrons used to connect the atoms from the total number of valence electrons. Remember each single bond is composed of 2 electrons.  For SO₂, 18 electrons (total) – 4 electrons (from the two bonds in step 1) = 14 electrons left over.  Therefore we have 14 electrons left to place around the molecule.
4. Add the appropriate number of electrons around each atom. Hydrogen requires 2 electrons, boron requires 6 electrons, and all other elements require 8 electrons.  Start by placing electrons on the outer atoms to give them a compete octet.  If more electrons are available, place them on the central atom.  If the central atom lacks an octet, form multiple bonds with outer atoms (see below).

In our example, S and O both require 8 electrons.  So first we put 6 electrons around one oxygen (which gives it 8 including the two in the bond) and another 6 electrons around the other oxygen.  At this point our structure will look like this:

 

 

Now we have only two electrons left to place around the S atom.  The question is, do we have enough electrons?  If we place the two electrons around the S atom, sulfur will have only 6 electrons (as shown in the following structure), and we need 8.

 

 

We need to use the electrons more efficiently by making one of the lone pairs on an O atom a double bond.  If we move a lone pair to make a double bond, we get the following structure.

 

 

This is the completed Lewis structure for SO₂ because all of the atoms are surrounded by eight electrons (octet rule!).  Remember an octet for hydrogen (H) is only two.

1. Procedure to Determine the Electron Group and Molecular Shapes of a Molecule. The electron group geometry can be determined by counting the number of groups of electrons (atoms + lone pairs) around the central atom and then looking up the appropriate geometry in Table 1.  In the case of SO₂, we count three groups from the Lewis structure (2 atoms + 1 lone pair).  From Table 1, the electron group geometry is trigonal planar.

The molecular shape can be determined by counting the number of atoms bonded to the central atom, and using the number of electron groups determined above to select the appropriate geometry from Table 2.  SO₂ has two bonded atoms and three electron pairs, so Table 2 indicates that the molecular shape is 120° bent.

 

1. Resonance Structures. Some molecules have more than one correct Lewis structure.  These are called resonance structures.  In order for a molecule to have resonance structures, it must have at least one multiple bond.  Molecules with only single bonds cannot have resonance structures.

In the case of SO₂, the molecule has been drawn above with the double bond to the oxygen to the right of the sulfur.  However, it could have also been drawn between the sulfur and the oxygen on the left, as shown below.  These are the resonance structures for SO₂.

 

 

 

1. Polarity of a Molecule. The last piece of information to be obtained about a molecule concerns the distribution of electron density and charges around the molecule.  A molecule with a uniform distribution of electron density is nonpolar; and one with an asymmetrical distribution is polar.  A molecule is nonpolar only if it has no lone pair electrons about the central atom and all groups attached to the central atom are identical (both conditions must be met to be nonpolar).  Another way to state this is if the electron group and molecular shapes are the same and the atoms attached to the central atom are identical, then the molecule is nonpolar.

In the case of SO₂, the Lewis structure shows us that the molecule is polar because the sulfur atom has a lone pair.

 

 

Procedure

Determine the Lewis structure, electron group geometry, molecular shape, presence or absence of resonance structures, and the polarity for a series of molecules given on the work­sheet.  Once you have filled in the worksheet, build a model of each compound using your model kit.  Compare the model you built with the responses you provided in the lab.  Carbon tetrachloride is worked out for you as an example.

 

 

 

Names:                                                                      Date:                                                             

 

	CCl4	BF3	SO3	CO2	ClO2–
Crude Sketch
Calculations (# of valence electrons, # of bonds, etc.	1(4) + 4(7)=32   32-4(2) = 24
Lewis Structure
# electron groups,electron group geometry	4tetrahedral
# of bonded atoms, molecular shape	4tetrahedral
Resonance structures (if any)	none
Polar or nonpolar	nonpolar

 

 

 

	H2O	SO42-	NO2+	PO43-	NO3–
Crude Sketch
Calculations (# of valence electrons, # of bonds, etc.
Lewis Structure
# electron groups,electron group geometry
# of bonded atoms, molecular shape
Resonance structures (if any)
Polar or nonpolar

 

 

 

 

	CO32-	SO2	NO2–	PF3	SiI4
Crude Sketch
Calculations (# of valence electrons, # of bonds, etc.
Lewis Structure
# electron groups,electron group geometry
# of bonded atoms, molecular shape
Resonance structures (if any)
Polar or nonpolar

 

 

 

	NH3	H3O+	NH4+	SO32-	CHCl3
Crude Sketch
Calculations (# of valence electrons, # of bonds, etc.
Lewis Structure
# electron groups,electron group geometry
# of bonded atoms, molecular shape
Resonance structures (if any)
Polar or nonpolar

 

 

 

 

 

LAB EXERCISE 1                             Name____________________________

 

Formulas and Names

 

1. Complete the table below by providing the formula and name of the compound formed in the rectangle where the cation (positive ion) and anion (negative ion) intersect.

 

Cl^–                                                 O ^2-                              P ^3-

Li+	LiCl Lithium Chloride
Ca 2+
Cr3+
Mn 4+

 

 

LAB EXERCISE 1                                   Name______________________________

 

2. Complete the table below by providing the formula and name of the compound formed in the rectangle where the cation (positive ion) and anion (negative ion) intersect.

 

NO₃^–                                 CO₃ ^2-                                     PO₄ ^3-

Na+	NaNO3 sodium nitrate
Pb2+
Fe 3+
Ti4+

 

 

LAB EXERCISE 1 (continued)               Name______________________________

 

3. Give the formulas for the following ionic compounds.

 

a- manganese (III) chloride          _________________________________

 

b- magnesium sulfide                   _________________________________

 

c- barium bromide                        _________________________________

 

d- potassium phosphide                _________________________________

 

e- cadmium acetate                       _________________________________

 

f- calcium permanganate              _________________________________

 

g- cesium bromate                        _________________________________

 

h- beryllium dichromate               _________________________________

 

i- ammonium peroxide                 _________________________________

 

j- iron (III) chloride                      _________________________________

 

k- uranium (V) sulfide                  _________________________________

 

l- nickel oxide                               _________________________________

 

m- tin (IV) nitride                         _________________________________

 

n- bismuth (III) hypochlorite        _________________________________

 

o- mercury (I) thiocyanate            _________________________________

 

LAB EXERCISE 1 (continued)               Name______________________________

4. Name the following ionic compounds.

 

a- BaI₂                               _________________________________

 

b- K₃PO₄                           _________________________________

 

c- AlI₃                               _________________________________

 

d- Ag₃N                            _________________________________

 

e- Li₂Cr₂O₇                       _________________________________

 

f- MnCO₃                          _________________________________

 

g- CsHCO₃                       _________________________________

 

h- Sr(BrO₂)₂                      _________________________________

 

i- Li₂CrO₄                         _________________________________

 

j- CrF₆                               _________________________________

 

k- TiO₂                              _________________________________

 

l- Co₃P₂                             _________________________________

 

m- MnS₂                           _________________________________

 

n- Zn(OH)₂                       _________________________________

 

o- Ga(ClO)₃                      _________________________________

 

LAB EXERCISE 1 (continued)               Name______________________________

5. Name the following molecular compounds.

 

a- NF₃                               _________________________________

 

b- P₂S₅                              _________________________________

 

c- SBr₂                              _________________________________

 

d- CO                                _________________________________

 

e- P₄S₃                               _________________________________

 

f- SiI₄                                   _________________________________

 

g- SCl₆                              _________________________________

 

h- OCl₄                              _________________________________

 

i- SeO₂                              _________________________________

 

j- N₂O₃                              _________________________________

 

k- B₂H₆                             _________________________________

 

l- AsI₃                                 _________________________________

 

LAB EXERCISE 1 (continued)         Name______________________________

6. Give the formulas for the following molecular compounds.

 

a- nitrogen disulfide                     _________________________________

 

b- phosphorus pentabromide        _________________________________

 

c- carbon tetrachloride                  _________________________________

 

d- sulfur hexabromide                   _________________________________

 

e- sulfur trioxide                           _________________________________

 

f- silicon disulfide                        _________________________________

 

g- diarsenic pentasulfide               _________________________________

 

h- boron trichloride                       _________________________________

 

i- dinitrogen monoxide                 _________________________________

 

j- carbon tetrahydride                   _________________________________

 

LAB EXERCISE 1 (continued)         Name______________________________

7. For each acid, name the anion bonded to hydrogen, then name the acid.

 

Anion name                                                     Acid name

 

a- HI                      _______________________                          _______________________

 

b- HCl                   _______________________                          _______________________

 

c- H₂S                   _______________________                          _______________________

 

d- H₃PO₄               _______________________                          _______________________

 

e- H₃PO₃               _______________________                          _______________________

 

f- H₂SO₄               _______________________                          _______________________

 

g- H₂SO₃               _______________________                          _______________________

 

h- HClO₄               _______________________                          _______________________

 

i- HClO                 _______________________                          _______________________

 

 

 

 

 

 

 

 

 

 

 

LAB EXERCISE 1 (continued)          Name______________________________

 

8. For each acid, give the formula of the anion bonded to hydrogen, then give the formula of the acid.

 

Anion formula                                     Acid formula

 

a- hydrobromic acid    ________________________            ________________________

 

b- hydrofluoric acid    ________________________            ________________________

 

c- hydrocyanic acid    ________________________            ________________________

 

d- nitric acid                ________________________            ________________________

 

e- sulfuric acid                        ________________________            ________________________

 

f- periodic acid            ________________________            ________________________

 

g- acetic acid               ________________________            ________________________

 

h- sulfurous acid         ________________________            ________________________

i- hypobromous acid   ________________________            ________________________