_____University of Puget Sound
Experiment 7
Department of Chemistry
Chem 110
MODELS: MOLECULAR BONDING AND COVALENT COMPOUNDS
1. Understand how to use VSEPR to predict 3-dimensional properties of
molecules.
a. Write the Lewis Dot structures corresponding to given chemical formula.
b. Infer steric numbers and electron-pairs from Lewis Dot structures.
c. Write VSEPR formulas (AXmEn) corresponding to given Lewis structures.
d. Infer electron-pair geometry and molecular shape from VSEPR formulas.
e. State ideal angles associated with a given electron-pair geometry.
f. Name the molecular geometry of a given VSEPR formula using standard
nomenclature.
g. Predict deviations from ideal bond angles.
2. Develop proficiency in describing the structure of organic molecules.
a. Translate among molecular formulas, structural formulas, condensed
structural formulas, and line drawings of simple hydrocarbons.
b. State essential properties of simple hydrocarbons.
c. State what is meant by an isomer, and draw all isomers of given molecules.
INTRODUCTION
The shapes or structures of molecules profoundly affect the function of chemical
compounds whether in inorganic minerals, organic polymers, or biochemical compounds
in living systems. Thus, one of the most important topics in the study of molecules is how
they are put together—how their atoms are arranged relative to each other in space. A
theory of chemical bonding that explains molecular geometry—the three-dimensional
arrangement of atoms in a molecule or polyatomic ion—is a useful tool for the modern
chemist. A molecule's properties are greatly influenced by the strength and directional
character of covalent bonds. For example, the same atoms may be linked to one another in
many different ways to yield distinctly different compounds. Thus, although there is only
one molecule with the molecular formula C3H8, known as propane, there are three
compounds with the molecular formula C5H12, known as pentanes. The three compounds
have different physical properties and can be isolated from each other. Such compounds
are called isomers.
The arrangement of atoms in a molecule imparts a specific shape to it that is responsible
for some of the unique properties of the substance. For example, if water were linear
instead of bent, the chemistry of life as we know it would be quite different. Simply having
the molecular formula for a compound does not necessarily allow us to predict any
physical or chemical properties for that compound. Why is a molecule of boron trichloride
(BCl3) flat or planar, whereas ammonia (NH3) is shaped like a pyramid? Why is the carbon
dioxide (CO2) molecule linear and nonpolar, whereas sulfur dioxide (SO2) is bent and
polar? These questions can be answered only with a better understanding of the threedimensional structure of molecules.
The first step in the prediction of the three-dimensional geometry of a molecule or
polyatomic ion is to write a valid Lewis electronic structure for the species under
consideration. The Lewis structure is only a two-dimensional representation of the
chemical bonding; the application of VSEPR theory is used to visualize the threedimensional relationship of atoms in a molecule or polyatomic ion. Finally, an actual threeModels/Molecular Bonding and Covalent Compounds
1
dimensional molecular model can be constructed to complete the visualization process,
and to investigate the molecular polarity properties.
Drawing Lewis Dot Structures
To draw a Lewis structure you must first decide which atoms in the molecule are
attached to each other. The most symmetrical arrangement of atoms often has the greatest
chance of being correct. Once the
general arrangement of atoms in the
EXAMPLE
molecule is known, the Lewis dot
structure can usually be drawn using
Draw the Lewis dot structure for Water, H2O
the following procedure (see text):
THE FOLLOWING GUIDELINES WILL
HELP TO DETERMINE PLAUSIBLE LEWIS
DOT STRUCTURES FOR MOST
MOLECULES, EVEN WHEN THEY DON'T
OBEY THE OCTET RULE.
1.
2.
GUIDELINES FOR DRAWING LEWIS
STRUCTURES
Determine the valence electron
configuration for each atom. Sum the
valence electrons for all of the atoms,
plus or minus any electrons gained or
lost for ionic charge, to determine the
total number of valence electrons for
the molecule. This number of electrons
should appear in your final structure.
Step 1 ----- one outer shell e– for each H, and 6
e– for the outer shell of O.
Total = 8 e–
Step 2 ----- O is the central atom:
HO H
Step 3 ----- two bonds ∴ 2e–/bond ⇒ 4
electrons in bonds
8 – 4 = 4e– remaining
Step 4 ----- the distribution of the remaining
electrons gives
Determine the central atom and
connect the atoms (using a single bond
for each at this point), arranging them
about the central atom using the
Each atom has an octet in this structure
general guidelines found here. Most
all electrons have been used.
likely first choice Central atoms- C, N,
Si, Ge, P, As, B, O, S, Se, Cl in this
order. Typically the less electronegative atom is central (more complex molecules will have more
than one central atom).
H O H
and
3.
Deduct two electrons for each bond shown in the 'skeleton' structure, drawn in #2 above, from the
total number of valence electrons as determined in #1.
4.
Distribute the remaining electrons in pairs around the atoms until each has its octet (or two electrons
for H, He or Li when they form covalent molecules), leaving the central atom for last.
5.
If there are two electrons ‘too few’ to go around, there is probably a double bond in the structure. If
you are four electrons short, there will probably be 2 double bonds or one triple bond in the
molecule. The halogens (Cl, Br, I) rarely form double bonds and H and F never do. If you have too
many electrons, put the extra pair or two on the central atom (such an "expanded octet" is possible
only if the central atom is in the third row of the periodic table or beyond.)
6.
If your structure has one or more double bonds, determine if the double bond could be written in
another position. If so, draw one or more resonance structures.
Models/Molecular Bonding and Covalent Compounds
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2-
O
O
C
O
2-
O
O
C
O
2-
O
O
C
O
Figure 1. Resonance structures of Carbonate ion.
MORE EXAMPLES
Methanol (CH3OH). The number of valence electrons in methanol is 14, which must
appear in the Lewis structure. Assuming C and O are the central atoms and using a pair of
electrons to bond each of the atoms together gives
H
H C O H
H
In this drawing, only 10 of the 14 valence electrons have been used, so the remaining 4
are used to complete the octet of the O atom.
H
H C O H
H
Ethene (C2H4). The number of valence electrons in ethene is 12, which must be used in
the Lewis structure. With 2 C’s as the central atoms, single bond connections give
H H
H C C H
In this structure, only 10 of the 12 valence electrons have been used, so 2 more valence
electrons will complete the structure. If these electrons were placed as a lone pair on either
of the carbon atoms, the other carbon atom would not have its octet satisfied. However, by
placing a double bond between the C atoms, the octets of both C atoms are satisfied. The
finished structure is
Hydrogen cyanide (HCN). The number of valence electrons in hydrogen cyanide is 10,
which must again be used in the Lewis structure. Assuming C is central, single bonds give
H C N
In the structure for HCN, there remain 6 electrons that may be used to satisfy the octets
for the C and N atoms. Using a triple bond arrangement and one pair of nonbonding
electrons on the N atom gives the finished structure as
H C
N
The structure for CH3OH includes five single bonds, no double bonds, and two lone
(non-bonded) electron pairs. The structure for the second compound, C2H4, includes four
Models/Molecular Bonding and Covalent Compounds
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single bonds, one double bond, and no lone pairs. The third molecule, HCN, contains one
single bond, one triple bond, and one lone pair.
Usually, bonding electron pairs are designated by a line between atoms, as shown in
the following structures.
H
H
C
H
O
METHANOL
H
H
H
C
H
C
ETHENE
H
H
C
N
HYDROGEN CYANIDE
Figure 2. Lewis structures showing bonds (pairs of electrons) as lines.
In general, molecules can be classified into two general categories as to Lewis dot
structures.
1.
Molecules or polyatomic ions that obey the octet rule are satisfied with eight
electrons in the valence shell of each atom. Again, hydrogen requires only two
electrons in its valence shell–a duet.
2.
Molecules or polyatomic ions that “violate” the octet rule by either of two ways:
(a) The central atom has an excess of valence shell electrons. This is possible with
transition elements and period 3 or higher main group elements.
(b) The central atom has less than an octet around the central atom, as is often
found when elements of the beryllium and aluminum families are the central atoms.
Using VSEPR Theory to Determine 3D Shapes
Read in Gilbert, et. al., Section 9.2, especially pages 413 through 424-Central Atoms with
Lone Pairs and Bonding Pairs of Electrons.
Now that a process for determining Lewis dot structures has been developed, how does
one make the transition to three–dimensional representations? The simplest way to do this
is to use what is known as the Valence Shell Electron Pair Repulsion (VSEPR) theory.
The basic assumption of this theory is that the valence shell electrons are found in pairs,
and that these pairs repel one another because they are negatively charged. The best
geometry for a molecular species is one that minimizes repulsion between the electron
pairs in the valence shell of the central atom. This electronic geometry is determined by the
number of “electron groups” around the central atom. An electron group may be a single
bonding pair, a lone (nonbonded) pair, or a multiple bond, which counts as one "electron
group" since the electrons in a multiple bond are localized in the region between the two
atoms.
The geometry of covalent molecules is summarized in Table 1, below. When describing
the shapes in VSEPR, a generic molecule AXmEn is considered: the central atom is A, the
attached atom (or group of atoms) is X, any lone pair of electrons is E, and the number of
each is designated by the integers m and n.
Models/Molecular Bonding and Covalent Compounds
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Table 1. Electronic Pair Geometries and Molecular Shapes from VSEPR
#Lone e- Pair ➙
0
1
2
3
SN = 2
bond angles = 180°
SN = 3
bond
angles = 120°
SN = 4
bond angles = 109.5°
SN = 5
bond angles = 90°, 120°
SN = 6
bond angles = 90°
NOTE: For some 3-D representations of the molecular shapes can be more clearly drawn with all the
atoms in the same plane. For example an alternate 3-D sketch of the trigonal planar AX3 is
alternate 3-D sketch of the angular or bent AX2E is
.
and an
When there are lone pairs on the central atom, one must be careful to distinguish
between the electronic geometry (the arrangement of all of the electron groups) and the
molecular geometry (the arrangement of the bonded atoms around the central atom). The
AXE formula or designation is an aid to determining these geometries. Use the Steric
Number (SN), the total number of electron groups (m+n) in Table 1 to find the electron-pair
geometry and ideal bond angles; use the number of lone electron pairs to find the
molecular geometry derived from this electron-pair geometry (as E's replace X's).
Models/Molecular Bonding and Covalent Compounds
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Generally, when describing the shape of a molecule, the molecular geometry is used rather
than the electron-pair geometry. Experimentally, it is difficult to determine the positions of
the electrons with any certainty; the nuclei of the atoms are located and the electrons are
assumed to be in the appropriate regions around them.
EXPERIMENTAL PROCEDURE
WORK IN PAIRS EXCEPT FOR INDIVIDUAL UNKNOWNS
The models you will be using for this experiment are called “ball-and- stick” models.
Plastic balls are used to represent the atoms and plastic sticks are used to represent the
electrons (bonding and nonbonding). The balls have been painted various colors to allow
you to distinguish different atoms. The conventional color scheme and Steric Number (see
Table 1) of central atoms for your model kit are as follows. (the yellow, blue and black balls
all have 4 bonding spots available and can be used for any variation of an SN=4 species.)
Table 2. Color Scheme of Ball and Stick Models
CENTRAL ATOM
STERIC
DESCRIPTION
NUMBER, SN
Brown
Trigonal bipyramidal
Silver
Yellow
COLOR
COLOR
DESCRIPTION
2, 3, 5
Green
Halogen
Octahedral
2, 6
White
Hydrogen
Oxygen or Sulfur
4
Gray
Flex Bond (for multiple
bonds)
Blue
Nitrogen
4
Black
Carbon
4
Part 1.1
Gray
Rigid Bond (for single bonds)
Molecules and Ions
Write the Lewis formulas of the following molecules (or ions) and construct a model of
each using the molecular model kit. Examine the model to help complete the table in your
laboratory notebook. State the number of bonding electron groups and lone electron pairs
for the central atom. Determine the VSEPR formula for the central atom, and predict the
molecule’s molecular geometry. Compare the shape of your model with the predicted
geometry. Finally, looking at your model, make a 3-D sketch and predict the molecule’s
polarity. The central atom is italicized and underlined. Some of the molecules (or ions) do
not conform to the Lewis Octet Rule. Organize your data table in the following format
(vertical lines are acceptable in this table).
Try to form a firm connection between the two-dimensional Lewis structure and the
three-dimensional actual structure. Since models will not always be available, it is helpful
to be able to visualize in three dimensions how a given two-dimensional Lewis structure
will appear.
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Table 3. Sample Data and Results Table
Molecular
Formula
Lewis Dot
Structure
SN
Electron-pair
Geometry
4 tetrahedral
NH3
CH4
…
…
With Respect to the Central Atom
# of Bonded
atoms
# of Lone
e- pairs
VSEPR
Formula (e.g.,
AXE)
3
1
AX3E
…
…
Molecular
Shape
3-D “Wedge and
line” Sketch with
angles
Trigonal
pyramidal
…
…
MOLECULES OR IONS FOR MODELS (Central atom underlined.)
1. NH3
4. H2O
7. PO43–
10. BrF3
2. CH4
5. H3O+
8. SF6
11. NO33. CH3Cl
6. SO2
9. PCl2F3
12. XeF2
13. Unknown
Part 1.2
Determination of Unknown Molecule (or Ion)
Part 2
Organic Molecules: Alkanes
Obtain an unknown molecule(s) (or ion(s)) from the instructor and complete the same
information for the unknown(s) as above in your notebook.
NOTE: Please put your responses to the items preceded by a bullet (•) on ONE SEPARATE
PAGE(S) These require an action from the student by noting the observation in the lab
notebook. Answer all Questions (Q1, Q2, ...) in the notebook on ANOTHER SEPARATE
PAGE(S); these questions will usually require some explanation of your answer.
A hydrocarbon is a compound that contains only hydrogen and carbon. An alkane is a
hydrocarbon with only single bonds. Alkanes are named using a root that defines the
number of carbon atoms followed by the suffix -ane. Complete the following table for the
first five alkanes:
Table 4. Sample Data Table.
Name
Number of
carbon
atoms
methane
1
ethane
2
Molecular
formula
C2H6
Structural
Formula
H
H
H
C
C
H
H
H
Condensed
Structure
Formula
CH3CH3
propane
butane
pentane
Models/Molecular Bonding and Covalent Compounds
7
•Construct a 3-dimensional model of ethane.
Rotate the C-C bond in ethane. Look down the end of the C–C bond in ethane and
rotate the C-C bond until the hydrogens on the front carbon overlap the hydrogens on the
back carbon. This is called the eclipsed conformation. Now rotate the C-C bond so that the
hydrogens on the front are half way between the hydrogens on the back. This is called the
staggered conformation. Now rotate the C–C bond by 120°. You should have another
staggered conformation. These different forms of a molecule, that differ only in rotations
about bonds, are referred to as conformers (or conformational isomers). For most small
molecules at room temperature rotation around single bonds is rapid. Hence it is not
possible to isolate different conformers for these molecules.
•Build a model of propane.
•Construct a model of butane by taking a model of methyl group (methane minus one
hydrogen) and connecting it to the end carbon of your model of propane (minus one
hydrogen).
Now take another model of methane and connect it to the center carbon of another
model of propane.
•Draw the structural formula for this molecule next to the one above.
Q.1. Are these two butane molecules the same? Compounds that have the same
molecular formula but different structural formulas are called isomers. Isomers have
different structures and hence different properties (e.g., mp, bp).
•Write the molecular formula for pentane.
For any hydrocarbon the molecular formula can be represented as CnHm.
Q.2. Comparing the molecular formulas of methane, ethane, propane, butane, and
pentane, find the relation between m and n and give the general molecular formula for
any alkane, expressing m in terms of n.
•Build a model of pentane.
Q.3. How many isomers are there of pentane?
•Write the structural formulas for each of these isomers. [By the way, there are 366,319
isomers of the C20 alkane!]
By now you are probably getting tired of drawing structural formulas! For larger
molecules it becomes cumbersome to write out all the atoms in a molecule. For larger
organic molecules chemists often use line structures. In these representations an alkane
chain is represented by a zigzag line. The ends and each "bend" in the line represent a
carbon atom. Hydrogens are normally not shown; each carbon is understood to have
enough hydrogens to satisfy its tetravalency (i.e., 4 bonds). The line drawings for butane
and pentane are shown below.
•Make line drawings for each of your isomers of pentane.
•Build cyclohexane, C6H12 , and benzene, see below, and compare. Consider types of
bonds, bond angles, etc.
•Make line drawings for cyclohexane and benzene.
Models/Molecular Bonding and Covalent Compounds
8
Figure 3. The structural formulas of cyclohexane and benzene.
•Write the molecular formula of benzene.
This type of compound is called aromatic. Because of the presence of two important
resonance forms, it behaves differently from compounds with ordinary double bonds.
Q.4. How does this structure of benzene differ from cyclohexane (i.e., shape, bond
angles, etc.)?
•Use the molecular models to build your favorite molecule. What is the name of this
molecule?
Models/Molecular Bonding and Covalent Compounds
9
WHAT TO DO:
BRING to lab
1)
From Exp. 6-Natural Pigments as Acid-Base Indicators – Please hand in
the following items in a STAPLED packet:
• a page presenting the absorption spectra copied to a Word document with
an appropriate caption (The plot should only have the pH indicated for
each spectra. The curves on the plot should be identified so as to
distinguish between them, but do not annotate plot with the data. The
figure caption can be brief by referring to the Table 2 in the caption for
data of plot.). Be sure the page has the proper title format (see guidelines).
The title format should be used on any computer generated document
handed in.
• the notebook copy page with a completed Table 2 and
• the remaining copy pages from your lab notebook including the tables,
observations and sample calculations (essentially all remaining pages
from this experiment) at the next lab session (the week of Oct. 22-26).
2)
Print a copy of this experiment, read it and bring it to your laboratory class.
3)
Bring your Lab Notebook with the PreLab assignment completed. You will
not be allowed to do the experiment without the prelab assignment completed.
PreLab Assignment –
1. Read in Gilbert, et. al., Section 9.2, especially pages 413 through 424-Central Atoms
with Lone Pairs and Bonding Pairs of Electrons.
2. Set up your lab notebook appropriately for this experiment including a title bar.
Prepare data tables similar to Table 3 and Table 4 found in the experimental section
of this handout. Your tables must be much larger than the example tables in this
handout.
Please note that the data table similar to Table 3 must be drawn landscape in your
notebook using 3 pages: Molecules 1-4 on the first page, Molecules 5-8 on the
second page and Molecules 9-13 on the third page. Complete the first column
(Molecular Formula) of the data table for all twelve species (leave space for #13
the unknown as well!). Leave one additional page for notes, etc.
The data table similar to Table 4 should also be drawn landscape on one page in
your notebook. The rows must be much larger than the example table indicates! You
will need to have two to three pages after the table for observations and questions.
Table 3. Sample Data and Results Table for Parts 1.1 and 1.2 - Molecules and Ions.
With Respect to the Central Atom
Molecular
Formula
Lewis Dot
Structure
Electron-pair
Geometry
4 tetrahedral
NH3
CH4
SN
…
…
# of Bonded
atoms
# of Lone
e- pairs
VSEPR
Formula (e.g.,
AXE)
3
1
AX3E
…
…
Models/Molecular Bonding and Covalent Compounds
Molecular
Shape
3-D “Wedge and
line” Sketch with
angles
Trigonal
pyramidal
…
…
10
Table 4. Sample Data Table for Part 2-Organic Molecules.
Number of
Molecular
Structural
Name
carbon atoms
formula
Formula
methane
1
…
ethane
2
C2H6
Condensed
Structural Formula
…
…
CH3CH3
propane
butane
pentane
In Lab:
Carefully complete the Part 1 data table as you build the models.
For Part 2, complete the Part 2 data table as you record all comments.
Answer all bulleted “•” items on ONE SEPARATE PAGE(S) and questions
"Q" in the lab notebook on another SEPARATE PAGE(S).
Please check your “copy” pages to see that dot structures and 3-D Structures
are clear
To be Turned in the following week: SEE the MOODLE Page for DETAILED
INSTRUCTIONS. In addition to a worksheet (available in class) hand in all
the “copy” pages that contain your completed data tables and the answers to
the bulleted items and questions.
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