4.3
DID YOU
KNOW
?
Confusing Mosquitoes
Mosquitoes can be dangerous when
they transmit diseases such as
malaria and West Nile fever.
Research has shown that molecules
with a spherical shape are better in
mosquito repellents than long, thin
molecules. It seems spherical molecules are better able to block the
sensory nerves in the mosquito’s
antennae. This makes it difficult for
mosquitoes to detect carbon
dioxide, moisture, and heat from
humans or animals.
VSEPR Theory
The shape of molecules has long been investigated through crystallography, using microscopes and polarimeters in the late 1800s and X-ray and other spectrographic techniques since the early 1900s. One of the most important applications of molecular shape
research is the study of enzymes. Enzymes are large proteins that are highly specific in
what they will react with. There are about three thousand enzymes in an average living
cell and each one carries out (catalyzes) a specific reaction. There is no room for error
without affecting the normal functioning of the cell; different molecular shapes help to
ensure that all processes occur properly. Despite extensive knowledge of existing enzymes,
the structure of these proteins is so complex that it is still effectively impossible to predict the shape an enzyme will take, even though the sequence of its constituent amino
acids is known. The study of molecular shapes, particularly of complex biological molecules, is still a dynamic field.
The Arrival of VSEPR
The valence bond theory created and popularized by Linus Pauling in the late 1930s
successfully explained many of the atomic orientations in molecules and ions, including
tetrahedral, trigonal planar, and linear orientations. Pauling’s main empirical work was
with the X-ray analysis of crystals. The valence bond theory of bonding, for which he is
primarily responsible, was created to explain what he “saw” in the laboratory. Pauling
extended the work of his friend and colleague, Gilbert Lewis, who is famous for creating electron dot structures.
When you studied the valence bond theory, including hybrid orbitals and sigma and
pi bonds, in Section 4.2, you became aware of the complexity of that approach. However,
it was not until 1957 that Australian Ronald Nyholm and Englishman Ron Gillespie
(Figure 1) created a much simpler theory for describing, explaining, and predicting the
stereochemistry of chemical elements and compounds. The theory that they created is
more effective for predicting the shape of molecules.
TRY THIS activity
Electrostatic Repulsion Model
The electrostatic repulsion of electron pairs around a central atom in a molecule can be
modelled using balloons.
Materials: safety glasses, 9 balloons, string
• Blow up the balloons and tie them off.
• Tie two balloons very close together and place them on a table.
(a) What is the orientation (e.g., angle) of two balloons?
• Repeat with three and then four balloons.
(b) What is the orientation of three balloons?
(c) What is the orientation of four balloons?
Figure 1
Dr. Ronald Gillespie co-created
VSEPR theory in 1957. He moved
from England to McMaster
University in Hamilton, Ontario, the
following year. His work in molecular
geometry and in chemistry education is renowned.
242 Chapter 4
• Reuse or recycle the balloons as directed.
(d) What are the pros and cons for using balloons for a physical model of the repulsion
of electron pairs about a central atom in a molecule?
NEL
Section 4.3
The name of the Nyholm-Gillespie theory is the valence-shell-electron-pair-repulsion theory, or VSEPR (pronounced “vesper”) theory. The theory is based on the electrical repulsion of bonded and unbonded electron pairs in a molecule or polyatomic
ion. The number of electron pairs can be counted by adding the number of bonded
atoms plus the number of lone pairs of electrons (Figure 3). Once the counting is done,
we can predict the 3-D distribution about the central atom by arranging all pairs of
electrons as far apart as possible.
VSEPR Theory
SUMMARY
VSEPR Valence Shell Electron Pair
Repulsion; pairs of electrons in the
valence shell of an atom stay as far
apart as possible to minimize the
repulsion of their negative charges
central atom the atom or atoms in
a molecule that has or have the
most bonding electrons; form the
most bonds
• Only the valence shell electrons of the central atom(s) are important for
molecular shape.
• Valence shell electrons are paired or will be paired in a molecule or polyatomic
ion.
lone
pairs
• Bonded pairs of electrons and lone pairs of electrons are treated approximately
equally.
H N H
H
• Valence shell electron pairs repel each other electrostatically.
• The molecular shape is determined by the positions of the electron pairs when
they are a maximum distance apart (with the lowest repulsion possible).
Using the VSEPR Theory
What is the shape of the hydrogen compounds of period 2: BeH2(s), BH3(g), CH4(g),
NH3(g), H2O(l), and HF(g)?
First, we draw Lewis structures of each of the molecules and then consider the arrangement of all pairs of electrons. The key idea is that all pairs of electrons repel each other
and try to get as far from each other as possible.
Lewis
structure
H Be H
Bond
pairs
Lone
pairs
Total
pairs
General
formula
2
0
2
AX2
Electron pair
arrangement
linear
O H
H
Figure 3
Both the ammonia molecule and the
water molecule have four pairs of
electrons surrounding the central
atom. Some of these are bonding
pairs and some are lone pairs.
Molecular
geometry
H
Be
H
linear
The Lewis structure indicates that BeH2 has two bonds and no lone pairs of electrons.
The total number of pairs of electrons around the central atom (Be) is two. These electron pairs repel each other. The farthest the electrons can get away from each other is
180°—a linear orientation.
Lewis
structure
Bond
pairs
Lone
pairs
Total
pairs
H
H B H
3
0
3
General
formula
AX3
Electron pair
arrangement
Molecular
geometry
trigonal planar
H
B
H
H
trigonal planar
BH3 has three bonds, which means three pairs of electrons around the central atom,
B. The three pairs of electrons repel one another to form a plane of bonds at 120° to
each other. This arrangement or geometry is called trigonal planar.
NEL
Chemical Bonding
243
LEARNING
TIP
Tetrahedral, trigonal planar, and
linear orientations of atoms are
common in molecules. To represent 3-D shapes on paper, a
solid lineis a bond in the
plane of the page; a dashed
line --- is a bond behind the
plane of the page; and a wedged
line is a bond out of the plane
of the page (toward the viewer).
X
tetrahedral
Lewis
structure
Bond
pairs
Lone
pairs
H
H C H
H
4
0
Total
pairs
4
General
formula
Electron pair
arrangement
Molecular
geometry
AX4
tetrahedral
H
C
H
tetrahedral
Lewis theory indicates that CH4 has four bonds or four pairs of electrons repelling each
other around the central atom, C. Experimental work and mathematics both agree that
a tetrahedral arrangement minimizes the repulsion. Tetrahedral bonds point toward the
corners of an equilateral pyramid at an angle of 109.5° to each other.
Lewis
structure
Bond
pairs
Lone
pairs
Total
pairs
General
formula
Electron pair
arrangement
H N H
H
3
1
4
AX3E
tetrahedral
Molecular
geometry
N
H
pyramidal
trigonal planar
X
linear
KNOW
H
H
X
DID YOU
H
H
?
Refinements of VSEPR Theory
A more detailed explanation of
molecular shape requires that
bonding pairs repel each other less
than lone pairs, and that electrons
of opposite spin repel each other
less than electrons with the same
spin.
The Lewis structure shows that NH3 has three bonding pairs (X in general formula)
and one lone pair (E in general formula) of electrons. The four groups of electrons
should repel each other to form a tetrahedral arrangement of the electron pairs just like
methane, CH4. The molecular geometry is always based on the atoms present and therefore, if we ignore the lone pair, the shape of the ammonia molecule is like a pyramid
(called pyramidal). We would expect the angle between the atoms, HNH to be 109.5°,
which is the angle for an ideal pyramid. However, in ammonia, the atoms form a pyramidal arrangement with an angle of 107.3°. This small difference is believed to occur
because there is slightly stronger repulsion between the lone pair of electrons and the
bonding pairs than between the bonding pairs. This causes the bonding pairs to be
pushed closer together.
Lewis
structure
O H
H
Bond
pairs
Lone
pairs
2
Total
pairs
2
General
formula
Electron pair
arrangement
AX2E2
4
Molecular
geometry
O
tetrahedral
H
H
V-shaped
According to the Lewis structure, the water molecule has two bonding pairs and two
lone pairs of electrons. The four pairs of electrons repel each other to produce a tetrahedral orientation. The geometry of the water molecule is called V-shaped with an angle
of 104.5°. Notice that this angle is again less than the ideal angle of 109.5° for a tetrahedral arrangement of electron pairs. The slightly stronger repulsion between the lone
pairs of electrons and the lone pairs and the bonding pairs is thought to force the bonding
electron pairs closer together.
244 Chapter 4
Lewis
structure
Bond
pairs
Lone
pairs
Total
pairs
General
formula
H F
1
3
4
AXE3
Electron pair
arrangement
Molecular
geometry
tetrahedral
H
F
linear
NEL
Section 4.3
Based upon the Lewis theory of bonding, the hydrogen fluoride molecule has one
bonding pair and three lone pairs of electrons. The four electron pairs repel to create a
tetrahedral arrangement for the electrons. This has little effect on the geometry of the
hydrogen fluoride molecule, which is linear at 180°—as all diatomic molecules are.
Shapes of Molecules
SUMMARY
VSEPR theory explains and predicts the geometry of molecules by counting pairs of
electrons that repel each other to minimize repulsion. The process for predicting the
shape of a molecule is summarized below.
Step 1 Draw the Lewis structure for the molecule, including the electron pairs
around the central atom.
Step 2 Count the total number of bonding pairs (bonded atoms) and lone pairs
of electrons around the central atom.
DID YOU
KNOW
?
Molecular Shape-Shifting
Knowing a molecule’s shape is
useful, but knowing how the
shape changes during a chemical
reaction is invaluable. For example,
the process by which HIV latches
onto its cellular host is believed to
depend on a molecular shape
change. A new technique of ultrafast X-ray diffraction now allows
scientists to observe how a molecule changes shape as it reacts.
This technique uses very short
X-ray and laser pulses to determine shapes on a time scale of
tens of picoseconds.
Step 3 Refer to Table 1 or Appendix C3 and use the number of pairs of electrons to predict the shape of the molecule.
Table 1 Using VSEPR Theory to Predict Molecular Shape
General Bond
formula* pairs
2
AX2
AX3
3
Lone
pairs
0
Total
pairs
2
0
3
Molecular shape
Geometry**
Shape diagram
linear (linear)
X
A
X
X
trigonal planar
(trigonal planar)
Examples
CO2, CS2
BF3, BH3
A
AX4
4
0
4
LEARNING
X
X
X
tetrahedral
(tetrahedral)
CH4, SiH4
TIP
For more molecular shapes see
Appendix C3.
A
X
X
X
AX3E
3
1
4
trigonal
pyramidal
(tetrahedral)
NH3, PCl3
A
X
X
X
AX2E2
AXE3
2
1
2
3
4
4
V-shaped
(tetrahedral)
linear (tetrahedral)
A
H2O, OCl2
X
X
A
X
HCl, BrF
*A is the central atom; X is another atom; E is a lone pair of electrons.
**The electron pair arrangement is in parenthesis.
Shape of a Polyatomic Ion
SAMPLE problem
Use the Lewis structure and VSEPR theory to predict the shape of a
sulfate ion, SO42-.
O
Determining the shape of a polyatomic ion is no different than determining
S O
O
the shape of a molecule. Again, you first obtain the Lewis structure of the
ion, as shown in Section 4.1. For the sulfate ion, the central sulfur atom is
O
surrounded by four oxygen atoms. There is a total of 32e. An acceptable
6 + 4(6) + 2 32e
Lewis structure is shown.
NEL
Chemical Bonding
245
Notice that you have four pairs of electrons around the central sulfur atom. This
corresponds to the AX4 category and therefore, the ion has a tetrahedral shape.
2–
O
S
O
O
O
SAMPLE problem
Shapes of Molecules with Two Central Atoms
Use the Lewis structure and VSEPR theory to predict the geometry of the B2F4 molecule. Provide your reasoning.
If a molecule has more than one central atom, such as two
F
F
boron atoms in this example, consider the shape around each
atom first, using the same procedure as molecules with one
F B
B F
central atom. Then combine these individual shapes to describe
or draw the overall geometry of the molecule. As in previous
examples, draw the Lewis structure first to determine the
number of pairs of electrons.
Three pairs of electrons around a boron atom means an AX3
F
F
general case, and hence a trigonal planar arrangement around
each central boron atom. This arrangement of electrons creates
B
B
the minimum repulsion of electron pairs.
Diboron tetrafluoride is composed of molecules with trigonal
F
F
planar shape around each central boron atom, producing an
overall planar molecule as shown to the right.
Practice
Understanding Concepts
1. Explain how the words that the VSEPR acronym represents communicate the main
ideas of this theory.
2. Use VSEPR theory to predict the geometry of a molecule of each of the following
substances. Draw a diagram showing the shape of each molecule.
(c) H2S(g)
(e) SiBr4(l)
(a) BeI2(s)
(b) PF3(g)
(d) BBr3(g)
(f) HCl(g)
3. Use VSEPR theory to determine the shape of each of the following polyatomic ions:
(a) PO43
(b) IO3
4. Cubane is a hydrocarbon with the formula, C8H8. It has a cubic shape, as its name
implies, with a carbon atom at each corner of the cube. This molecule is very
unstable and some researchers have been seriously injured when crystals of the
compound exploded while being scooped out of a bottle. Not surprisingly, it has
some uses as an explosive.
(a) According to VSEPR theory, what should be the shape around each carbon
atom? Why?
(b) If we assume an ideal cubic shape, what would be the bond angles around
the carbon?
(c) Explain how your answers to (a) and (b) suggest why this molecule is so
unstable.
246 Chapter 4
NEL
Section 4.3
Applying Inquiry Skills
5. Where did the evidence come from that led to the creation of VSEPR theory?
6. Locate two or more VSEPR Web sites and compare them. Which do you prefer?
List two or more criteria for evaluating the sites and indicate how each site did
based upon each criterion.
GO
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Making Connections
7. Enzymes make up the largest and most highly specialized class of protein mole-
cules. Describe briefly how their three-dimensional structure influences their function. How does the “lock-and-key” analogy relate to molecular shapes and the
highly specific nature of enzyme reactions?
GO
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8. What are optical isomers? Describe the role that molecular shape plays in classi-
fying optical isomers.
GO
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Extension
9. The VSEPR theory can be extended to five and six electron pairs to explain several
other shapes of molecules, as shown in Table 2.
Several other shapes are possible if one or more of the total number of electron
pairs is a lone pair.
(a) Draw Lewis structures for PCl5 and SF6.
(b) Draw the Lewis structure for ClF3. If the two lone pairs are in the trigonal
plane, predict the molecular shape.
(c) Draw the Lewis structure for ICl4. If the two lone pairs are above and below
the plane of the atoms, predict the molecular shape.
Table 2 Expanded VSEPR Theory to Predict Molecular Shape
General
formula
AX5
Bond
pairs
5
Lone
pairs
0
Total
pairs
5
Molecular shape
Geometry
Shape diagram
trigonal
X
bipyramidal
(trigonal
bipyramidal)
90˚
X
PCl5
X
A
120˚
Example
X
X
AX6
6
0
6
octahedral
(octahedral)
SF6
X
90˚
X
90˚
X
A
X
X
X
NEL
Chemical Bonding
247
Multiple Bonding in VSEPR Models
Evidence of multiple bonding can be obtained from, for example, the reaction rate of
hydrocarbons; e.g., the fast reaction of alkenes and alkynes with bromine or potassium
permanganate, compared to the slower reaction rate with alkanes. Further evidence
indicates that the multiple bonds are shorter and stronger than single bonds between the
same kind of atoms. Evidence from crystallography (e.g., the X-ray analysis of crystals)
indicates that these multiple bonds can be treated like single bonds for describing,
explaining, and predicting the shape of a molecule. This has implications for using
VSEPR theory for molecules containing multiple bonds. Let’s look at some examples.
SAMPLE problem
VSEPR and Double Covalent Bonds
Ethene (ethylene, C2H4(g)) is the simplest hydrocarbon with multiple bonding.
Crystallography indicates that the orientation around the central carbon atoms is
trigonal planar. Is VSEPR theory able to explain the empirically determined shape
of this molecule?
DID YOU
KNOW
?
Smelling Molecular Shapes
Limonene, C10H16, is a naturally
occurring compound that can exist
in two different isomeric forms.
One form is present in orange
rinds and has a pleasant orange
odour. The other isomer is identical
in composition but different in
shape. This isomer smells like pine
or turpentine and is found in
lemons and pine needles.
The first step in testing the ability of VSEPR theory to explain
H H
the shape of ethene is to draw a Lewis structure of the molH C C H
ecule.
The second step is to count the number of “pairs” of
electrons around the central atoms (the carbon atoms). In
H
H
the case of multiple bonding such as a double bond, the
double bond contains two pairs of electrons. The crystalloC
C
graphic evidence indicates a trigonal planar arrangement.
H
H
The only way that VSEPR theory can accommodate this evidence is to count a multiple bond as a single group of
electrons. In other words, you are counting the number of bonded atoms. There are three
bonded atoms or sets of bonding electrons around each of the central carbon atoms
(AX3). These electron groups repel each other to obtain minimum repulsion and, thus, a
minimum energy state. The result, according to VSEPR theory, is a trigonal planar orientation—three atoms on a plane at 120°.
VSEPR theory passes the test by being able to explain the trigonal planar shape of
ethene. Now let’s see if VSEPR theory can pass another test by predicting the stereochemistry of the ethyne molecule.
Example 1
LEARNING
TIP
Treat double and triple covalent
bonds as one group of electrons
when using VSEPR theory to
predict shapes of molecules
containing these bonds.
Predict the shape and draw the diagram of the ethyne (acetylene, C2H2(g)) molecule.
Solution
H C
C H
H
C
H
C
The shape of the ethyne molecule is linear.
Example 2
Predict the shape and draw the diagram for a nitrite ion, NO2.
ACTIVITY 4.3.1
Shapes of Molecules (p. 277)
Molecular models are important
tools for understanding molecular
shapes.
248 Chapter 4
Solution
O N O
−
N
O
O
The shape of the nitrite ion is trigonal planar.
NEL
Section 4.3
Practice
Understanding Concepts
10. In order to make the rules of VSEPR theory work, how must multiple (double and
triple) bonds be treated?
11. Use Lewis structures and VSEPR theory to predict the shapes of the following
molecules:
(a) CO2(g), carbon dioxide (dry ice)
(b) HCN(g), hydrogen cyanide (odour of bitter almonds)
(c) C3H6(g), propene (monomer for polypropylene)
(d) C3H4(g), propyne
(e) H2CO(g), methanal (formaldehyde)
(f) CO(g), carbon monoxide (deadly gas)
Applying Inquiry Skills
12. Is VSEPR a successful scientific theory? Defend your answer.
Making Connections
13. Astronomers have detected an amazing variety of molecules in interstellar space.
(a) One interesting molecule is cyanodiacetylene, HC5N. Draw a structural diagram for this molecule and predict its shape.
(b) How do astronomers detect molecules in space?
DID YOU
KNOW
?
Theories
Just as scientists have special definitions for words that apply in the
context of science (e.g., the definition of “work” in physics), philosophers who study the nature of
chemistry have special definitions
for terms such as “theory.” To them
a theory is not a hypothesis. To
philosophers of chemistry a theory
uses the unobservable (such as
electrons and bonds) to explain
observables (such as chemical
and physical properties).
As a student you have probably
found words that are used in different contexts in and out of the
science classroom, such as “salt”
and “decomposition.” What are
some other words that you use differently in other contexts?
Molecular Geometry Research
A Canadian researcher doing important work in molecular geometry is Dr. Richard
Bader (Figure 4), Professor Emeritus at McMaster University. Dr. Bader’s work includes
the theoretical determination of electron density maps for small molecules, giving a visible interpretation of the molecular shapes and bonding within molecules (Figure 5). Note
that the shape shown is consistent with the type of structure you have previously used
to represent ethene. Viewed this way, single and double bonds are not really single or
double structures—they are just different concentrations of electron density. Similarly,
the unique carboncarbon bonds in benzene are part of a total molecular electron
orbital structure, resulting in a particular electron density in the region around the ring.
Dr. Bader’s work builds on previous bonding theories that have a critical limitation.
“We have in chemistry an understanding based on a classification scheme that is both
powerful and at the same time, because of its empirical nature, limited.” Dr. Bader applies
quantum mechanics theory to determine the atomic structures of molecules and crystals. To the extent that this theory is supported by empirical evidence, it allows development
of new theory, and may eventually lead to the ability to use computer models to accurately explain and predict forces, structures, and properties that currently can only be
observed and measured.
Figure 4
Dr. Richard Bader
Figure 5
A representation of the electron
density for an ethene molecule.
NEL
Chemical Bonding
249
EXPLORE an issue
Take a Stand: Linus Pauling and the
Vitamin C Controversy
Linus Pauling became interested in chemistry at a young age
because a friend had a chemistry set. Pauling graduated from
Oregon State University in 1922 and obtained his Ph.D. from the
California Institute of Technology in 1925. After a year in Europe
studying with Sommerfeld, he became a chemistry professor at
the California Institute of Technology in 1927 and remained
there throughout his academic career.
Pauling’s scientific fame came from his theory of chemical
bonding, including the ideas of a shared pair of electrons, polar
covalent bonds, electronegativity, and resonance structures.
These ideas revolutionized thinking about molecular structure.
For this work Pauling was awarded the Nobel Prize in chemistry
in 1954. He continued his study of molecular structure and was
one of the first to suggest helical structures of proteins and the
relationship between disease and abnormal molecular structure.
After the Second World War, Pauling used his fame as a
Nobel Prize winner to vigorously fight the nuclear arms race of
the United States and the Soviet Union. For his outspoken leadership against nuclear testing he was awarded the Nobel Peace
Prize in 1962, becoming one of a very few people who have won
two Nobel Prizes.
Pauling’s fame as a scientist and as a social activist meant
that he could easily command media attention whenever he
spoke. When he announced in 1970 that large doses (megadoses) of vitamin C could prevent the common cold, and other
illnesses as well, many people paid close attention. However,
not everyone agreed that vitamin C is as useful in megadoses
as Pauling claimed. There is still a huge interest in Pauling’s
suggestion, in spite of no clear scientific evidence supporting
his claim and some scientific studies that dispute it.
Decision-Making Skills
Define the Issue
Analyze the Issue
Identify Alternatives
Defend the Position
Research
Evaluate
(a) Briefly describe some claims being made today for the
beneficial use of large doses of vitamin C.
(b) Briefly describe some objections and criticisms of the
claimed benefits.
In small groups, discuss the following questions and obtain a
consensus within the group. Report on your conclusions.
(c) Pauling and other proponents of the benefits of megadoses of vitamin C, and the doctors and scientists
opposed to this view have all claimed that science is on
their side. Anyone can have an opinion or belief, but science requires more. What are the requirements for a
claim to be scientifically valid? List some criteria.
(d) To what extent do you think Pauling’s fame influenced
public and scientific opinion about the benefits of
vitamin C? Suppose someone with no scientific training
and unknown to the public made this claim, would
anyone notice or consider it seriously? If a Nobel Prize
winner makes a claim disputed by other lesser-known
scientists, how do we decide what to believe?
(e) The vitamin C controversy is not the first time a famous
scientist has made a claim that is disputed by most of
the scientific community. What are the repercussions for
a scientist who goes against the rest of the scientific
community? Who usually “wins”? Is the practice and
work of science completely objective?
GO
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Section 4.3 Questions
Understanding Concepts
1. Use Lewis structures and VSEPR theory to predict the
molecular shape of the following molecules. Include a 3-D
representation of each molecule.
(a) H2S(g), hydrogen sulfide (poisonous gas)
(b) BBr3(l), boron tribromide (density of 2.7 g/mL)
(c) PCl3(l), phosphorus trichloride
(d) SiBr4(l), silicon tetrabromide
(e) BeI2(s), beryllium iodide (soluble in CS2(l))
2. Use Lewis structures and VSEPR theory to predict the molec-
ular shape around the central atom(s) of each of the following
molecules. Provide a 3-D representation of each molecule.
(a) CS2(l), carbon disulfide (solvent)
(b) HCOOH(g), acetic acid (vinegar)
(c) N2H4(l), hydrazine (toxic; explosive)
(d) H2O2(l), hydrogen peroxide (disinfectant)
(e) CH3CCCH3(l), 2-butyne (reacts rapidly with bromine)
250 Chapter 4
3. Draw the Lewis structure and describe the shape of each
of the following ions:
(a) IO4
(b) SO32
(c) ClO2
Making Connections
4. Briefly describe Dr. Bader’s contribution to our under-
standing of molecules.
5. Search the Internet for information on the current workplace
and position of Dr. Ronald Gillespie, the co-creator of VSEPR
theory. What degrees does he hold? What are some of the
awards that he has won? What is his major topic of research?
6. Some scientists argue that taste has developed as a pro-
tective mechanism. Many poisonous molecules taste bitter
and ones that are useful to us have a more pleasant, often
sweet, taste. Write a brief summary about the relation of
taste to molecular structure.
GO
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NEL