Molecular Geometry and VSEPR
We gratefully acknowledge Portland Community College for the use of this experiment.
Objectives



To construct molecular models for covalently bonded atoms in molecules and polyatomic
ions
To use valence bond theory to account for the bonding in covalently bonded systems
To apply the valence shell electron pair repulsion theory and valence bond theory to the
geometries and polarities of molecules
Discussion
An understanding of the structure of a m olecule is f undamental to an explanation of its ch emical and
physical properties. For example, water is a liquid at room temperature, dissolves innum erable salts and
sugars, is more dense than ice, and has a low vapor pressure, in part, because its molecules are bent rather
than linear.
Lewis Theory of Bonding
Because structure is so important, chemists have developed a number of theories to explain and predict
molecular geometries. In 1916, G. N. Lewis developed a theory that focused on the significance of
valence electrons (electrons in the outer shell) in chemical reactions and in bonding. He proposed the
"octet rule", in which atoms form bonds by losing, gaining, or sharing enough electrons to have the same
number of valence electrons (eight) as the nearest noble gas in the Periodic Table. The bond formed is
ionic or covalent depending on whether the electrons are transferred or shared between atoms.
The octet rule is valid for nearly all compounds formed between atoms of the second period and for a
large number of compounds formed between atoms of other representative elements. There are three
types of exceptions to the octet rule: molecules having an atom with less than eight electrons, molecules
having an atom with more than eight electrons and molecules having an odd number of electrons.
Elements of the second and third family may also form stable molecules in which they have less than
eight electrons surrounding the central element. For example the molecule BF3. Elements beyond the
third period may form stable molecules with more than eight electrons surrounding the central element.
For example PCl5. Radicals have an odd number of electrons therefore they have unpaired electrons, for
example the molecule NO.
More than 8 electrons
Less than 8 electrons
Odd number of Electrons- Radical
Lewis emphasized the importance of valence electrons using a Lewis symbol to represent
an atom of the element-the symbol for the element surrounded with its corresponding
number of valence electrons. For main group elements the number of valence electrons
is the same as their family number.
For example, Na is the Lewis symbol for the sodium atom, and Ca is the Lewis symbol for calcium.
The Lewis symbols for the atoms are used to account for the bonding in a compound; the resulting
representation is called the Lewis formula (or Lewis structure) for the compound. The Lewis formula for
water shows that by sharing valence electrons between the oxygen and hydrogen, all three atoms obtain
the same number of valence electrons as the nearest noble gas; that is, the hydrogen atoms are
isoelectronic with helium atoms and the oxygen atom is isoelectronic with the neon atom.
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. Sometimes in these structures a pair of bonding electrons is substituted
by a line, while non-bonding electrons or lone pairs are represented with dots. 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.
Double bonds, and triple bonds, occur in many molecules, usually between C, O, N, and/or S atoms.
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, C2H6O:
H
O
H
H
C
H
H
C
H
H
H
C
O
H
H
C
H
H
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 as in, C2H2Cl2:
H
Cl
C
Cl
C
C
Cl
H
Cl
Cl
C
C
H
H
Cl
H
C
H
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:
H
H
C
H
C
C
H
H
C
C
C
H
C
H
C
C
C
H
H
C
C
H
H
H
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 resonant structure.
Although a Lewis formula accounts for the bonding based on the valence electrons on each atom, it does
not explain how the valence electrons are shared, nor does it predict any three-dimensional structure for a
molecule. The valence bond (VB) theory and the valence shell electron pair repulsion (VSEPR) theory
provide some insight into the nature of the bonding between atoms and the three-dimensional structure of
the molecule.
Valence Bond (VB) Theory of Bonding
The basic postulate of VB theory is that a covalent bond forms when a pair of electrons is shared by
overlapping atomic orbitals between bonding atoms. When these overlapping orbitals point directly at one
another, creating a cylindrical symmetry of electron density along the internuclear axis, the bond is a
sigma () bond. For example, a  bond forms between a hydrogen atom and a fluorine atom making the
HF molecule when the 1s atomic orbital of the hydrogen atom (having a single electron) overlaps with a
2p atomic orbital of the fluorine atom (also having a single electron). The result is a pair of shared valence
electrons along the internuclear axis between the two atoms.
1s
2p
H
 bond
F
H  F
In 1931, Linus Pauling proposed that the orientation of orbitals involved in a  bonding determines the
three-dimensional structure of a polyatomic molecule or ion. A detailed look at the bonding and structure
of methane, CH4, supports this. The valence shell electron configuration for carbon is 2s22p2. Only two
"free" 2p electrons are available for bonding on the isolated carbon atom. For methane, however, all four
C-H bonds are equivalent, all H-C-H bond angles are 109.5° and the geometric structure is tetrahedral.
To account for the properties of CH4, VB theory states that the bonded carbon atom must not maintain the
use of the same 2s and 2p atomic orbitals as on the "isolated" atom. But rather, it must adjust its valence
shell orbitals to allow each of its four valence electrons to bond; in so doing, an independent region of
space for each electron is available for bonding.
2s
2p
As we began with four valence shell atomic orbitals, one s orbital and three p orbitals, we name the four
“new” orbitals in the bonded atom as a combination of these four, we say that each of the four “new”
orbitals is an sp3 orbital. The word hybrid is often tagged to the name of these “new” orbitals; the four sp3
hybrid orbitals are a “hybrid” of the one s and the three p orbitals from the “isolated” carbon atom; each
has a single electron available for sharing in the formation of a covalent bond.
2p
hybridization
2s
Hybridization of the valence shell orbitals on the carbon atom.
sp3
Each sp3 hybrid orbital bond overlaps with a 1s atomic orbital on the hydrogen atom to form four
equivalent C to H  bonds. The four C-H bonds are oriented in space as a tetrahedron to lessen
electrostatic interaction between the hydrogen nuclei.
s orbitals of hydrogen
sp3 hybrid orbitals of carbon
Four sp3 hybrid orbitals on a carbon in methane produce a tetrahedral structure.
For some molecules one or more pairs of nonbonding electrons in the valence shell of the central atom of
the molecule occupy hybrid orbitals. Because these hybrid orbitals are the same as the hybrid orbitals
forming the  bonds, they also contribute to the geometry of the molecule. Therefore, the orientation of
all hybrid valence shell orbitals (those forming the  bonds and those containing pairs of nonbonding
electrons) determines the geometry of the molecule.
Hybridization modes for valence shell orbitals in bonding atoms and their corresponding geometries are
summarized in Table 1.
Hybridization
Bonding
Electron
Pairs
Lone
Pairs
Molecular
Geometry
Bond
Angle
Example
sp
High Electron
Density Areas
Around Central
Atom
2
2
0
Linear
180
BeF2
sp2
3
3
0
Trigonal Planar
120
BF3
sp2
3
2
1
Bent / Angular
<120
GeF2
sp3
4
4
0
Tetrahedral
109.5
CH4
sp3
4
3
1
Trigonal
Pyramidal
<109.5
NH3
sp3
4
2
2
Bent / Angular
<109.5
H2O
sp3d
5
5
0
Trigonal
Bipyramidal
90,
120
PCl5
sp3d
5
4
1
SeeSaw
SF4
sp3d
5
3
2
T-Shape
sp3d
5
2
3
Linear
180,
120,
90
180,
90
180
XeF2
sp3d2
6
6
0
Octahedral
90
SF6
sp3d2
6
5
1
Square
Pyramidal
90,
180
IF5
sp3d2
6
4
2
Square Planar
90,
180
XeF4
Valence Shell Electron Pair Repulsion Theory of Bonding
PCl3
VSEPR theory proposes that the geometry of a molecule is determined by the repulsive interaction of
electron pairs in the valence shell of its central atom. The orientation is such that the distance between the
electron pairs is maximized so that electron pair-electron pair interactions are minimized. Construction of
the Lewis formula of a molecule provides the first link in predicting the geometry of the molecule.
Methane, CH4, has four bonding electron pairs in the valence shell of its carbon atom (the central atom in
the molecule). Repulsive interactions between these four electron pairs are minimized when the electron
pairs are positioned at the vertices of a tetrahedron. One can generalize that all molecules having four
electron pairs in the valence shell of their central atom have a tetrahedral arrangement of these electron
pairs. The nitrogen atom in ammonia, NH3, and the oxygen atom in water, H2O, also have four electron
pairs in their valence shell!
H
N
H
O
H
H
H
The arrangement of the bonding and nonbonding electron pairs around the central atom gives rise to the
corresponding electronic geometry, i.e., tetrahedral. The bonding electrons give rise to the molecular
geometry and actual shape of the molecule, trigonal pyramidal for NH3 and bent for H2O.
Various molecular shapes can be determined from Lewis formulas and VSEPR formulas which use the
following notations:
A
Xm
En
Refers to the central atom
Refers to "m" number of bonding pairs of electrons on A
Refers to "n" number of nonbonding pairs of electrons on A
If a molecule has the formula AXmEn, it means there are m + n electron pairs in the valence shell of A,
the central atom of the molecule; m are bonding and n are nonbonding electron pairs. For NH3 it would
AX3E and for H2O it would be : AX2E2
It should be noted here that electrons on the central atom contributing to a multiple bond do not affect the
geometry of a molecule. For example in SO2, the sulfur atom is sp2 hybridized, the VSEPR formula is
AX2E, and the molecule is V-shaped. Further applications of these theories are presented in more
advanced chemistry courses.
Table 1 summarizes the geometries of molecules or ions and the corresponding bond angle(s) predicted
by the VB and VSEPR theories.
Polarity of Molecules or Ions
Once the three-dimensional shape of a molecule is determined, its polarity can be qualitatively
understood. A molecule is polar if an unsymmetrical distribution of electrons exists in the molecule
resulting in a partial separation of charges. An unsymmetrical distribution of charge occurs when bonded
atoms have different electronegativities; the atoms having a higher electronegativity (electronegativity
increases as you move up and to the right in the periodic table) more strongly attract bonding electrons,
acquiring a greater electron density and a partial negative charge, -, relative to another portion of the
molecule, +. Thus all heteronuclear diatomic molecules are polar. The greater the electronegativity
differences of the atoms, the greater the distortion of the electron density, thus, the more polar is the
molecule.
The direction and magnitude of the polarity are represented by a vector, drawn in the direction of the
greater electron density; the + center of the molecule is indicated by a plus sign on the tail of the vector.
For example, the HF molecule is more polar than the HCl molecule because the fluorine atom is more
electronegative than the chlorine; both are more electronegative than the hydrogen atom. Therefore, the
+ is on the hydrogen; the vector is drawn in the direction of the fluorine and chlorine.
H-F
H-Cl
If more than one polar bond exists in a molecule, the entire molecule may be polar or nonpolar,
depending on the geometry of the molecule since the polarity from one bond may be cancelled by the
polarity of another. Consider BF3 and OF2. The BF3 molecule has a VSEPR formula of AX3, where each
B-F bond is polar. Table 1 shows that the BF3 molecule has a trigonal planar geometry (Figure 4). The
three more electronegative fluorine atoms attract the bonding electron pairs from the boron atom with
equal magnitude, that is all dipoles are cancelled.
Polarity of bonds in BF3
A geometric sum of the magnitude (all the same) and direction of the three vectors equals zero. Because
in the BF3 molecule there is no resultant vector that can be drawn to indicate a resultant + or - electron
density in the molecule, it is said to be a nonpolar molecule, even though each B-F bond is polar.
The OF2 molecule has the VSEPR formula AX2E2 and each O-F bond is polar. Table 1 shows that OF2
has a V-shaped geometry. The geometric sum of the magnitude (all the same) and direction of the two
vectors produces a resultant vector that is not equal to zero. A resultant vector can be drawn indicating a
+ and - electron density in the molecule. In such cases the molecule is polar; therefore, OF2 is a polar
molecule.
Experimentally, polar molecules are attracted to an electric field; nonpolar molecules are not. In an
experiment using an electrostatically charged rod, a polar liquid, such as water, flowing near an
electrically charged rod, is deflected slightly whereas a nonpolar liquid, such as benzene, is unaffected by
the electric charge.
Although the conclusions we have drawn 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 we have cited for octet rule 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 threedimensional 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
assemble models for a sizeable number of common chemical species and interpret them in the ways we
have discussed.
ALL INFORMATION FOR THESE MOLESCULES IS TO BE PALCED DIRECTLY INTO
YOUR LAB NOTEBOOKS
Using an appropriate set of molecular models, construct the following molecules/polyatomic ions, and
write their Lewis formulas. Determine the number of bonding orbitals (or electron pairs involved in
bonding), nonbonding orbitals (or number of nonbonding electron pairs), and degree of hybridization of
the valence shell orbitals on the central atom. Also determine the VSEPR formula, the geometry of the
molecule/polyatomic ion, and its polarity.
Molecule or
Polyatomic
Ion
1. CF3Cl
Lewis Structure
Data
Bonding orbitals or pairs
4
Nonbonding orbitals or
0
pairs
Hybridization
sp3
VSEPR formula
AX4
Geometry tetrahedral
Polar or nonpolar nonpolar
Resonance
no
Isomers
none
2. NH3
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
3. H2O
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
4. NO3-
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
5. C2H6
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Sketch of the 3-D
Molecule or
Polyatomic
Ion
6. C2H3O2Cl
Lewis Structure
Data
Bonding orbitals or pairs
7. PO43-
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
8. BF4-
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
9. C2H3OCl
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
10. PC12F2
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
11. SF6
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Sketch of the 3-D
Molecule or
Polyatomic
Ion
12. PCl2F3
Lewis Structure
Data
Bonding orbitals or pairs
13. BrF3
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
14. SF4
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
15. IF4+
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
16. SO32-
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Sketch of the 3-D
Molecule or
Polyatomic
Ion
17. XeF2
18. SF2
Lewis Structure
Data
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Bonding orbitals or pairs
Nonbonding orbitals or
pairs
Hybridization
VSEPR formula
Geometry
Polar or nonpolar
Resonance
Isomers
Sketch of the 3-D