CHEM 403
Exp 4
Experiment 4:
THE GEOMETRICAL STRUCTURE OF MOLECULES*
Many years ago it was observed that, in many of its compounds, the carbon atom formed
four chemical linkages to other atoms. As early as 1870, graphic formulas of carbon compounds
were drawn as shown:
Although such drawings as these would imply that the atom-atom linkages, indicated by
flat lines, lie in a plane, chemical evidence, particularly the existence of only one substance with
the graphic formula
requires that the linkages be directed in three dimensions toward the corners of a tetrahedron, at
the center of which is the carbon atom.
The physical significance of the chemical linkages between atoms, expressed by the lines
in molecular structure diagrams, became evident soon after the discovery of the electron. In 1916
in a classic paper, G. N. Lewis suggested, based on chemical evidence, that the single bonds in
graphic formulas involve two electrons and that an atom tends to hold eight electrons in its
outermost or valence shell.
Lewis’ proposal that atoms generally have eight electrons in their outer shells proved to be
extremely useful and has come to be known as the octet rule. It can be applied to many atoms, but
is particularly important in the treatment of covalent compounds of atoms in the second row of the
Periodic Table. For atoms such as carbon, oxygen, nitrogen, and fluorine, the eight valence
electrons occur in pairs that occupy tetrahedral positions around the central atom core. Some of
the electron pairs do not participate directly in chemical bonding and are called unshared,
nonbonding, or lone pairs; however, the structures of compounds containing such unshared pairs
reflect the tetrahedral arrangement of the four pairs of valence shell electrons. In the H2O molecule,
which obeys the octet rule, the four pairs of electrons around the central oxygen atom occupy
essentially tetrahedral positions; there are two unshared nonbonding pairs and two bonding pairs
that are shared by the O atom and the two H atoms. The H-O-H bond angle is nearly but not exactly
tetrahedral since the properties of shared and unshared pairs of electrons are not exactly alike.
*Adapted from Slowinski, E. J., Wolsey, W. C. Chemical Principles in the Laboratory 9th ed.
CHEM 403
Exp 4
Most molecules obey the octet rule. Essentially, all organic molecules obey the rule, and
so do most inorganic molecules and ions. For species that obey the octet rule it is possible to draw
electron dot, or Lewis, structures. The previous drawing of the H2O molecule is an example of a
Lewis structure. Here are several others:
In each of these structures, there are eight electrons around each atom (except for H atoms,
which always have two electrons). There are two electrons in each bond. When counting electrons
in these structures, one considers the electrons in a bond between two atoms as belonging to the
atom under consideration. In the CH2Cl2 molecule just shown, for example, the Cl atoms each
have eight electrons, including the two in the single bond to the C atom. The C atom also has eight
electrons, two from each of the four bonds to that atom. The bonding and nonbonding electrons in
Lewis structures are all from the outermost shells of the atoms involved, and are the so-called
valence electrons of those atoms. For the main group elements, the number of valence electrons in
an atom is equal to the last digit in the group number of the element in the Periodic Table. Carbon,
in Group 4, has four valence electrons in its atoms; hydrogen, in Group 1, has one; chlorine, in
Group 17, has seven valence electrons. In an octet rule structure the valence electrons from all the
atoms are arranged in such a way that each atom, except hydrogen, has eight electrons.
Often it is quite easy to construct an octet rule structure for a molecule. Given that an
oxygen atom has six valence electrons (Group 6) and a hydrogen atom has one, it is clear that one
O and two H atoms have a total of eight valence electrons; the octet rule structure for H2O, which
was discussed earlier, follows by inspection. Structures like that of H2O, involving only single
bonds and nonbonding electron pairs, are common. Sometimes, however, there is a “shortage” of
electrons; that is, it is not possible to construct an octet rule structure in which all the electron pairs
are either in single bonds or are nonbonding. C2H4 is a typical example of such a species. In such
cases, octet rule structures can often be made in which two atoms are bonded by two pairs, rather
than one pair, of electrons. The two pairs of electrons form a double bond. In the C2H4 molecule,
shown above, the C atoms each get four of their electrons from the double bond. The assumption
that electrons behave this way is supported by the fact that the C=C double bond is both shorter
and stronger than the C-C single bond in the C2H6 molecule, also pictured above. Double bonds,
and triple bonds, occur in many molecules, usually between C, O, N, and/or S atoms.
Lewis structures can be used to predict molecular and ionic geometries. All that is needed
is to assume that the four pairs of electrons around each atom are arranged tetrahedrally. You have
seen how that assumption leads to the correct geometry for H2O. Applying the same principle to
the species whose Lewis structures were listed earlier would predict, correctly, that the CH2Cl2
molecule would be tetrahedral (roughly anyway), that NH3 would be pyramidal (with the
nonbonding electron pair sticking up from the pyramid made from the atoms), that the bond angles
in C2H6 are all tetrahedral, and that the C2H4 molecule is planar (the two bonding pairs in the
double bond are in a sort of banana bonding arrangement above and below the plane of the
molecule). In describing molecular geometry, only the positions of the atomic nuclei, not the
electrons, are included. The NH3 molecule is pyramidal, not tetrahedral.
CHEM 403
Exp 4
It is also possible to predict polarity from Lewis structures. Polar molecules have their
center of positive charge at a different point than their center of negative charge. This separation
of charges produces a dipole moment in the molecule. All heteronuclear diatomic molecules are
polar, but covalent bonds between different kinds of atoms are polar. In some molecules the
polarity from one bond may be canceled by that from others. Carbon dioxide, CO2 which is linear,
is an overall nonpolar molecule. Methane, CH4, which is tetrahedral, is also nonpolar. Among the
species whose Lewis structures have been listed, H2O, CH2Cl2, NH3, and -OH are polar. C2H6 and
C2H4 are nonpolar.
For some molecules with a given molecular formula, it is possible to satisfy the octet rule
with different atomic arrangements. A simple example would be
The two molecules are called isomers of each other, and the phenomenon is called isomerism.
Although the molecular formulas of both substances are the same, C2H6O, their properties differ
markedly because of their different atomic arrangements.
Isomerism is very common, particularly in organic chemistry, and when double bonds are
present, isomerism can occur in very small molecules:
The first two isomers result from the fact that there is no rotation around a double bond, although
such rotation can occur around single bonds. The third isomeric structure cannot be converted to
either of the first two without breaking bonds.
With certain molecules, given a fixed atomic geometry, it is possible to satisfy the octet
rule with more than one bonding arrangement. The classic example is benzene, whose molecular
formula is C6H6:
These two structures are called resonance structures, and molecules such as benzene, which
have two or more resonance structures, are said to exhibit resonance. The actual bonding in such
molecules is thought to be an average of the bonding present in the resonance structures. The
stability of molecules exhibiting resonance is found to be higher than that anticipated for any single
resonance structure.
CHEM 403
Exp 4
Although the conclusions regarding molecular geometry and polarity can be obtained from
Lewis structures, it is much easier to draw such conclusions from models of molecules and ions.
The rules for assembling Lewis structures transfer readily to models. In many ways, the models
are easier to construct than are the drawings of Lewis structures on paper. In addition, the models
are three-dimensional and hence much more representative of the actual species. Using the models,
it is relatively easy to see both geometry and polarity, as well as to deduce Lewis structures. In this
experiment, you will have the opportunity to use models to help visualize a number of common
chemical species and interpret them in the ways discussed.
Procedure
In this experiment, you will practice drawing Lewis structures for compounds where all
atoms (with the exception of hydrogen) obey the octet rule; such atoms have four electron pairs
around the central core.
In writing and interpreting Lewis structures, it is possible, indeed desirable, to proceed in
a systematic manner. The recommended steps for completing a Lewis structure and determining
its properties will be illustrated below with a molecule whose formula is CH2O.
1. Determine the total number of valence electrons in the species. This is easily done once you
realize that the number of valence electrons on an atom is equal to the number of the main group
to which the atom belongs in the Periodic Table. For CH2O,
C → Group 4; H → Group 1; O → Group 6
Therefore, each carbon atom in a molecule or ion contributes four electrons, each hydrogen
atom one electron, and each oxygen atom six electrons. The total number of valence electrons
equals the sum of the valence electrons on all the atoms in the species. For CH2O this total
would be 4 + (2 x 1) + 6, or 12 valence electrons. If you are working with an ion, add one
electron for each negative charge or subtract one for each positive charge on the ion.
2. On paper or using a model kit, assemble a skeleton structure for the molecule, joining all atoms
by single bonds. In some cases this can only be done in one way. Usually, however, there are
various possibilities, some of which are more reasonable than others. Typically, the central
molecule in a structure is written first. For many small molecules, most of the other atoms are
attached directly to the central atom. In CH2O, the skeleton structure can be assembled by
connecting the C atom to both the two H atoms and to the O atom.
3. Count the number of electrons used, and the number left over, from assembling the skeleton
structure. Each bond represents two electrons, so the skeleton structure has used (3 x 2), or 6 of
the 12 valence electrons; there are 6 electrons remaining to be used to complete the Lewis
structure.
4. Distribute the remaining electrons to satisfy the octet rule for each atom. Each H atom has one
bond, or two electrons, and does not need any more electrons. The C atom has three bonds, or
six electrons, and needs two more electrons to satisfy the octet rule. The O atom has one bond,
or two electrons, and needs six more electrons to satisfy the octet rule. If you had eight electrons
CHEM 403
Exp 4
remaining, you could simply place them in pairs (called lone-pairs) around the C and O atoms
as needed to complete the Lewis structure. However, in this example, you do not have enough
electrons to fully satisfy all octets using single bonds and lone pairs and must take an additional
step to complete the structure.
5. If you have too few electrons to satisfy all atoms’ octets, you will need to form additional bonds
between atoms. By adding another bond between the C and O atoms, two of your remaining
electrons can be shared between the atoms and help to satisfy both atoms’ octets. By drawing a
double bond between the C and O atoms, the C atom no longer needs any more electrons, and
has a full octet. The remaining four electrons can be added around the O atom as two lone-pairs
to complete its octet, as seen in Figure 1. If you assembled this molecule using a model kit, it
would look similar to Figure 2, with the bonds and lone pairs represented by sticks.
Figure 1: Lewis Structure of CH2O
Figure 2: Ball and Stick Model of CH2O
6. To describe the molecular geometry of the molecule CH2O, first determine the electron pair
geometry of the central atom. In this example, the C atom has three sets of electron pairs coming
from it; the double bond counts as a single “attachment” from the central C atom. Referencing
a table of geometries, this electron pair geometry is trigonal planar. To determine the molecular
geometry, count the number of lone pairs on the central C atom. In this example, there are no
lone pairs, so the molecular geometry is also trigonal planar.
7. To describe the polarity of the molecule CH2O, compare the atoms or electron groups attached
to the central atom. In this example, since oxygen is a more polar atom than the hydrogens, it
will pull electron density unevenly from the central carbon atom, creating a polar molecule.
8. Finally, investigate the possibility of the existence of isomers or resonance structures. It turns
out that in the case of CH2O, one can construct an isomeric form that obeys the octet rule, in
which the central atom is oxygen rather than carbon. However, as a general rule, carbon atoms
CHEM 403
Exp 4
almost always form a total of four bonds; nonbonding electron pairs on carbon atoms are very
rare.
Resonance structures can be common in molecules with double bonds, and usually arise when
there is a choice as to which atom is bonded with a double bond and which atom takes a lone
pair. For CH2O, however, there are no resonance structures, as there was only one option for
placing the double bond, as hydrogen atoms only form one bond.
A. Using the steps outlined above, construct Lewis structures of the molecules and ions listed
below. In the Data & Calculations section, draw the complete Lewis structure for each molecule,
showing nonbonding as well as bonding electrons. Describe the geometry of the molecule or
ion, and state whether the species is polar. Finally, draw the Lewis structures of any likely
isomers or resonance forms.
CH4
CH2Cl2
CH4O
H2O
H3O+
HF
NH3
H2O2
N2
P4
C2H4
C2H2Br2
C2H2
SO2
SO42CO2
SCNNO3HNO3
C2H4Cl2
B. If stability requires that each atom obey the octet rule, predict the stability of the following
species:
PCl3
H3O
CH2
CO
CHEM 403 – Exp 4
Name:______________________________ Section: _________
Data & Calculations
A.
Species
CH4
CH2Cl2
CH4O
H2O
H3O+
HF
NH3
H2O2
N2
P4
C2H4
Lewis Structure
Molecular
Geometry
Polar?
Isomers or
resonance structures
Species
Lewis Structure
Molecular
Geometry
Polar?
Isomers or
resonance structures
C2H2Br2
C2H2
SO2
SO42-
CO2
SCN-
NO3-
C2H4Cl2
B.
Stability Predicted for PCl3 _________, H3O _________, CH2 _________, and CO _________
CHEM 403 – Exp 4
Name:______________________________ Section: _________
Prelab Assignment
You are asked by your instructor to draw the Lewis structure of the NH2Cl molecule. Being of a
conservative nature, you proceed as directed in the Procedure section.
1. First, find the number of valence electrons in NH2Cl. For counting purposes with Lewis
structures, the number of valence electrons in an atom of a main group element is equal to the
last digit in the group number of that element in the Periodic Table.
N is in Group ____________; H is in Group ____________; Cl is in Group ____________
In NH2Cl there is a total of ____________ valence electrons.
2. Assemble a skeleton structure for the molecule, connecting the N, H’s, and Cl to make one unit.
Use the rule that N atoms typically form three bonds, whereas Cl and H atoms usually form
only one. The N atom, listed first in the formula, will be the central atom.
3. How many electrons did you need to make the skeleton structure? ____________ How many
electrons are left over? ____________
4. If your model is to obey the octet rule, each atom must have four electron pairs (except for
hydrogen atoms, which need and can only have one electron pair). Complete your Lewis
structure below with the remaining electrons, which can be represented as nonbonding electron
pairs around the atoms.
5. Describe the molecular geometry of NH2Cl (Refer to a table of molecular geometries for this
question). ___________________
6. Is the NH2Cl molecule polar? ____________ Why?
7. Would you expect NH2Cl to have any isomeric forms? ____________ Explain your reasoning.
Would NH2Cl have any resonance structures? ____________ If so, draw them below.
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