Section 7.3 Molecular Geometry and Lewis Dot Structures
7.3 Molecular Geometry and Lewis Dot Structures
Molecular
formulas and
bonding
Everyone knows that the formula for water is H2O. This
tells us that two hydrogens and one oxygen bond
together to form a molecule of water. However, it
doesn’t tell us how they are bonded together. Are the
hydrogen atoms attached to each other? You could
imagine both possibilities shown in the diagram on the
right. One is correct and one is not. Lewis structures are
a convenient tool for figuring out the correct way to
connect atoms into molecules.
Lewis dot
diagrams
If you use the Lewis dot diagrams, you can predict that both hydrogen atoms bond to a
central oxygen. You can also predict the formula of H2O. To assemble a molecule from
individual atoms, you share one unpaired electron from one atom with an unpaired
electron from another atom. Now both electrons “belong” to each atom, and a covalent
bond is formed. Let’s see how this would be done to form a water molecule.
Solving the
molecule
puzzle
You can think of the Lewis structures for individual atoms like puzzle pieces that can be
put together to assemble molecules. The goal is to end up with each atom having no
unpaired electrons and to be surrounded by eight valence electrons (unless the atom is
hydrogen, in which case two paired electrons should be next to each hydrogen atom).
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Isomers
Same
molecular
formula but
different
properties
Sometimes there are two or more different but correct ways to fit the atomic puzzle
pieces together. This gives you different molecules with the same chemical formula. For
example, C2H6O can form ethanol, the alcohol that is found in beer, wine, and liquor.
The same C2H6O can also be dimethyl ether, often used in spray cans as a propellant.
These two molecules have the same chemical formula but are very different substances
with different properties. When you have multiple molecules that can be represented by
the same chemical formula, the molecules are called isomers. Isomers are formed by
bonding the same group of atoms together in different ways. The diagram shows two
isomers of C2H6O.
Give three isomers for the formula C3H8O. Show the Lewis dot diagram and
the structural formula for each molecule.
Asked:
The Lewis dot diagrams and structural formulas for the three molecules represented by the formula C3H8O
Given:
Carbon has four unpaired electrons, hydrogen has one, and oxygen
has two. Three carbons, eight hydrogens and one oxygen form each
molecule.
Relationships: The atoms will bond together such that all unpaired electrons will be
paired up with electrons from other atoms.
Solve:
isomer: a specific structure of a molecule, only used when a chemical formula could
represent more than one molecule.
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Section 7.3 Molecular Geometry and Lewis Dot Structures
Double and triple bonds
Double and
triple bonds
Sharing a pair of electrons (one from each
atom) is called a single bond. It is possible
for atoms to share two or even three pairs
of electrons. Sharing two pairs of electrons
is a double bond and sharing three is a
triple bond. Carbon, nitrogen, and oxygen
commonly form double and triple bonds.
Order of bond
formation
The process for writing Lewis dot structures for molecules with double and/or triple
bonds is very similar to what you have done already. First you single bond all the atoms
together. Once you do this you will find that there are still some unpaired electrons left
over. If those unpaired electrons are on two atoms that are already sharing electrons, then
you can bring those two unpaired electrons into the same region where the first pair of
shared electrons is already drawn. This can only be done with unpaired electrons on two
atoms that are already sharing electrons. If, after single bonding all of the atoms together
your unpaired electrons are not on two atoms that are already bonded, then you need to
change how you originally put the atoms together to form single bonds.
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Molecular geometry
Lewis dot
structures are
2D but real
molecules are
3D
Lewis dot structures can all be drawn on a flat twodimensional piece of paper. However, molecules are
three-dimensional (3D) objects. In fact, the precise
3D shape of a molecule is very important in
determining its physical and chemical properties,
such as boiling point and viscosity. For complex
molecules like those typically found in biological
systems, shape can make the difference between a
drug that treats cancer effectively or one that does
nothing at all.
VSEPR
Using Lewis dot structures, we can predict the shapes of
simple molecules and gain insights into shapes of larger ones.
Using the Valence Shell Electron Pair Repulsion (VSEPR)
theory, you can make predictions about the angles that
attached atoms will form. Since the first three letters of
VSEPR stand for “Valence Shell Electron,” we are talking
about the valence electrons, the ones represented by Lewis
dots. The last two letters stand for “Pair Repulsion.” Paired electrons are not shared in
chemical bonds, but they do affect the shape of a molecule. Paired electrons repel each
other and also repel shared electrons that are involved in a chemical bond.
Bonds and
electron
density
Because similar charges repel each other, electrons that are being shared in bonds and
unshared pairs of electrons will push away from each other. This repulsion is part of what
causes a molecule to form a particular shape. Each bond or unshared pair of electrons
represents a region of electron density around an atom. Depending on how many
regions of electron density exist, different shapes will be formed.
VSEPR (Valence Shell Electron Pair Repulsion): a theory that states that the
shapes of molecules are dictated, in part, by the repulsion of the shared electrons and
the unshared pairs of electrons.
region of electron density: an area represented by shared or unshared electrons
around an atom.
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Section 7.3 Molecular Geometry and Lewis Dot Structures
Two areas of electron density
Linear bond
shape:
two points
form a line
If you rub a balloon against your hair it will grab some
electrons from you, making the balloon negatively
charged. If you had two negatively charged balloons and
brought them close together, they would repel each other
because they are both have the same sign of charge. They
push away from each other in opposite directions until
they are as far away as possible from each other.
Repulsion and
bond shape
The same thing happens when there are two regions of electron density around an atom.
The electrons in both regions push away to be as far apart as possible. For example,
consider acetylene (C2H2). Around each carbon there are only two areas where electrons
are located—the ones being shared in the triple bond between the carbons and the ones
being shared in the single bond between the carbon and hydrogen atoms. These electrons
repel each other until they are as far away as possible from each other. This causes the
molecule to be linear around each of the carbon atoms.
There are two isomers for the formula C3H4. Show the Lewis dot diagram for
each molecule, and indicate which atoms are at the center of a linear part of
the molecules.
Asked:
The linear parts of each isomer of C3H4
Given:
There are two different isomers. Part of each molecule will be linear.
The molecules are made from three carbons and four hydrogens.
Relationships: Each atom that has two regions of electron density around it will
form a linear part of the molecule.
Solve:
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Three areas of electron density
Three points
form a plane
Now imagine that you have three negatively
charged balloons. Again they will repel each other
and move as far away from each other as they
possibly can. In this case the balloons will point to
the corners of a triangle, forming 120° angles
between the balloons.
Trigonal planar
When there are three regions of
electron density around an atom,
the electrons in all three regions
push away until they are as far
apart as possible. A good example
is ethene (C2H4). There are three
areas of electron density around
each carbon atom. One area is
being shared in the double bond
between carbons. The other two are
shared in the single bonds between the carbon and each of two hydrogen atoms. When
three regions of electron density repel each other, a trigonal planar shape is formed.
Acetic acid when mixed with water is commonly known as vinegar and has the
formula C2H4O2. The correct isomer has both oxygens bonded to the same
carbon. Draw the Lewis dot structure for this isomer and indicate where the
molecule will be trigonal planar.
Asked:
Given:
The trigonal planar parts of acetic acid
The formula for acetic acid is C2H4O2 and both oxygens are bonded
to the same carbon.
Relationships: Each atom that has three regions of electron density around it will
form a trigonal planar part of the molecule.
Solve:
trigonal planar: a flat triangular geometry typical of molecules with three regions of
electron density surrounding a central atom. The angle between the atoms is 120
degrees.
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Section 7.3 Molecular Geometry and Lewis Dot Structures
Four areas of electron density
Tetrahedral
geometry: 3D
You might think that the shape formed by four
charged balloons would look something like an
“×” or a “+” with all four balloons as far apart
from each other as they can get in a plane.
However, there is a way to orient the balloons so
that they are even farther apart: Have the
balloons point to the corners of a tetrahedron, a
pyramid shape formed from four triangles. If the
four balloons lie in a plane then the angle
between each balloon is 90°, but pointing the
balloons into the four corners of a tetrahedron
increases the angle to 109.5°, moving them even
farther apart.
Lone pairs
In all of the Lewis dot structures we have seen so far, all of the electrons have been
shared between atoms in covalent bonds. When looking at the 3D version of the
molecule, each bond represented one or more pairs of shared electrons. However, atoms
such as nitrogen and oxygen have some electrons that are already paired, even before any
bonding has occurred. These electrons, known as lone pairs, are “invisible” in the balland-stick view of a molecule, but their presence affects the geometry of the molecule just
as much (if not more) than the electrons being shared in bonds.
lone pairs: electrons that are paired up in a Lewis dot diagram but are not shared
between atoms in a covalent bond.
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Multiple shapes due to four regions of electron density
Different
geometric
shapes
There are three shapes that are all based on having electron density in four places around
an atom: tetrahedral, trigonal pyramidal, and bent. If there are four single bonds to a
central atom, then the four bonded atoms mimic the balloons and form a tetrahedral
shape. If there are three atoms bonded to a central atom that also has one unshared pair of
electrons, then the unshared pair acts just like a shared pair, repelling the other electrons
that are being shared in the three bonds. However, because one corner of the tetrahedron
is “missing,” we call this trigonal pyramidal. If there are two atoms bonded to a
central atom that has two pairs of unshared electrons, then they all repel as in previous
examples, except that the three atoms form a bent shape.
What shapes are formed within the isomer of C4H5NO, which has a triple bond
connecting nitrogen?
Solve:
tetrahedral: a shape formed when a central atom is bonded to four other atoms that
all point into the corners of a tetrahedron.
trigonal pyramidal: a three-dimensional shape formed when a central atom has
one unshared pair of electrons and is bonded to three other atoms whose central
plane does not include the central atom.
bent: a shape formed when a central atom has two unshared pairs of electrons and is
bonded to two other atoms.
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Trans fats and your health
You have probably heard how some fats are good and some are bad for you. Terms such as saturated,
unsaturated, trans fats, and partially hydrogenated can all be found on nutritional labels. These labels are
really describing the different structures of fat molecules. Fats and oils are made by both plants and animals
as a means for storing chemical energy. Each molecule is composed of a glycerol backbone containing
three carbons, and each carbon is attached to a long-chain fatty acid. Chemists often refer to fats and oils as
triglycerides, where tri- means three and glyceride means it has a glycerol backbone. The triglyceride
resembles a three-pronged fork, with each prong representing a long-chain fatty acid.
The term “saturated” means each carbon in a fatty acid chain is bonded to the largest number of hydrogen
atoms possible. Saturated fats such as butter are solid at room temperature, because their long straightchain fatty acids contain only single carbon-to-carbon bonds. This allows them to stack very neatly
together, like the logs in the walls of a log cabin.
The fatty acids in unsaturated fats contain some carbon-carbon double
bonds. These double bonds create kinks along the chain that interfere
with the stacking of these fat molecules. They are unable to stack in an
organized way. Unsaturated fats, such as olive oil, are liquid at room
temperature.
Saturated fats can have a negative impact on cardiovascular health.
Unsaturated fats, although still high in calories, are beneficial, helping
to reduce the bad cholesterol and increase the good type of cholesterol.
Originally, scientists thought the story of good fat versus bad fat ended
with saturated versus unsaturated. However, “trans fats,” unnatural fats
made during the partial hydrogenation process, are as bad as or even
worse than saturated fats. The difference in this example is the way the
carbons are bonded together around the double bond.
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If you walk down the grocery store aisles, you will see items
being advertised as containing zero grams of “trans fat.” Some
food labels also say that they are “heart healthy.” Companies
are making sure consumers notice their efforts to reduce or
eliminate the harmful trans fats. So what is a trans fat, and why
are they considered bad for us?
There are two very interesting facts about naturally occurring
unsaturated fats. Naturally occurring unsaturated fats form cis
double bonds. “Cis” in Latin literally means “on the same
side.” This refers to the fact that both hydrogen atoms are on
the same side of the double bond. Second, cis fats react more
readily with oxygen in the air, resulting in a shorter shelf life
for a food product. When cis fats react with oxygen, they break
apart, forming smaller chains that taste and smell rancid. Manufacturers discovered that when they used the
healthier naturally occurring fats to make their food products, it resulted in a much shorter shelf life. So
manufacturers searched for a way to increase the shelf life of items made with unsaturated fat. The method
of partial hydrogenation was developed in the early 1900s, and this process allowed for a longer shelf life
for food products. CriscoTM was the first product of this new technology. Proctor and Gamble first marketed
Crisco in 1911.
“Trans” in Latin means “across,” which indicates the hydrogen atoms are on opposite sides of the double
bond. When a naturally occurring cis fat is heated at high pressure in the presence of hydrogen gas (H2) and
a metal catalyst, such as nickel, one hydrogen atom is added to each carbon in a double bond. This results
in the carbon becoming single bonded as it is in a saturated fat. During the partial hydrogenation process
only some of the double bonds are transformed to single bonds. The remaining double bonds are converted
from cis to trans. This is how trans fat is made. Trans fats are less likely to react with oxygen than the
naturally occurring cis fats, and this makes them last longer in food products, increasing shelf life.
Originally, it was thought that these fats were healthier for us than saturated fats. Margarine was advertised
as a heart-healthy alternative to butter! Doctors recommended margarine for patients with heart disease or
high cholesterol up until the late 1980s. Current research has shown that trans fats cause an increase in
blood levels of low density lipoproteins (LDLs). These LDLs are known as the “bad” cholesterol, because
as they travel in the blood they can deposit some cholesterol in the arteries. It is also important to note that
just because a label says “zero grams of trans fats” does not mean it is healthy to consume large quantities.
Try to avoid eating these unhealthy fats. Instead, get as much as possible of your daily fat calories from
“good oils” such as olive oil.
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Chapter 7 Review
Vocabulary
Match each word to the sentence where it best
fits.
Section 7.1
electrostatic force
polarize
nonpolar covalent
bond
polar covalent bond
ionic bond
1. The _____________ is what holds atoms together
in bonds.
2. When the electrons are shared unevenly between
two atoms they cause the molecule to
_____________ in that area.
3. A/An _____________ forms when electrons are
shared unevenly between atoms.
4. If one or more electrons are transferred between
atoms causing them to become positively and
negatively charged, then a/an _____________
forms.
5. If two atoms have the same or very similar
electronegativities, then the bond between them
will be a/an _____________.
Section 7.2
free radical
antioxidant
octet rule
covalent bonds
unpaired electrons
paired electrons
6. A/An _____________ is highly reactive as a result
of having one unpaired electron and can do
damage to other necessary molecules in your cells.
7. One reason dark-colored fruits and vegetables are
considered healthy is their high content of
_____________ molecules.
8. The _____________ states that elements transfer
or share electrons in chemical bonds to reach a
stable configuration of eight electrons.
9. In _____________ electrons are shared between
atoms instead of transferred.
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10. _____________ form bonds and _____________
do not form bonds.
Section 7.3
isomer
lone pair
VSEPR
tetrahedral
region of electron
density
trigonal pyramidal
bent
trigonal planar
11. If an atom has a/an _____________ then some
valence electrons will not be shared in covalent
bonds.
12. You can predict the shapes of simple molecules
using _____________ theory.
13. Each bond and each pair of electrons not shared in
a bond around an atom is considered to be a/an
_____________.
14. If you can make more then one molecule from the
same set of atoms, then each molecule made from
that set is called a/an _____________.
15. The shape typically formed in a molecule around
an oxygen atom is _____________.
16. There are two shapes that involve four atoms. The
one that has all the atoms lying flat is
_____________, whereas the one where they can’t
be all placed in the same flat surface is
_____________.
17. If you had four things that are trying to repel as far
away from each other as possible, they will point
into the corners of a/an _____________ shape.
Conceptual Questions
Section 7.1
18. Why don’t all atoms repel each other if the outer
part of the atom consists of negative electrons?
19. Describe why sharing electrons causes atoms to
bond together.
20. In a ball-and-stick view of a molecule, what do the
sticks represent?
A NATURAL APPROACH TO CHEMISTRY
21. Describe the main difference between covalent
and ionic bonds.
22. The surface view of a molecule makes it look like
the molecule is solid. What does this surface really
represent?
23. Describe what the red, white, and blue colors
applied to some surface views of a molecule tell
you about the charge of the molecule near that part
of its surface.
24. If a covalent bond is formed when electrons are
shared between two atoms, then what must be true
of the electronegativity of both atoms?
a. The electronegativities must be very different
for each atom.
b. The electronegativities must be exactly the
same for each atom.
c. The electronegativities must be similar, but not
necessarily the same.
25. Explain the reasoning behind your answer to the
previous question. Be sure to talk about the effect
an atom’s electronegativity has on the electrons.
26. Describe the difference in how electrons are
shared in a nonpolar covalent bond versus a polar
covalent bond.
27. Which of the molecules below is bonded together
with a polar covalent bond, and how do you
know?
28. Which part of molecule “B” in the previous
question is partially negative?
31. All of the atoms are the same in a chunk of metal,
so they all have the same ionization energy and
electronegativity; they are both low. Why don’t we
call the type of bonds formed between the atoms
in a metal “covalent bonds”?
32. Most molecules are made from more than two
atoms bonded together, so there is more than just a
single bond in a typical molecule. If one of the
bonds is a polar covalent bond, but the rest of the
bonds are nonpolar covalent bonds, then what can
you conclude?
a. The molecule is considered to be nonpolar.
b. The molecule is considered to be polar.
c. The molecule has a polar part or may be
entirely polar.
Section 7.2
33. Why do atoms of elements in the same group on
the periodic table tend to form the same charge if
they become an ion, or form the same number of
covalent bonds when bonded in this way to other
atoms?
34. Using the properties of ionization energy and
electronegativity, explain why atoms of elements
from the left side of the periodic table tend to form
positive ions but atoms from elements on the right
side of the table form negative ions.
35. Which valence electrons are the ones involved in
ionization and covalent bonding?
a. the unpaired electrons from the Lewis dot
diagram of the neutral unbonded atom
b. the paired electrons from the Lewis dot
diagram of the neutral unbonded atom
36. How many different isomers are shown below?
29. Why does having one atom with low ionization
energy and low electronegativity paired with an
atom that has high ionization energy and high
electronegativity result in the formation of ions
and an ionic bond?
30. Why do most atomic level images of ionic
substances show many atoms forming a crystal,
but images of molecular substances usually just
show a specific set of bonded atoms?
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225
Chapter 7 Review
37. Are these two molecules isomers of each other or
not? Explain.
44. What are the three different possible shapes
formed around an atom that is surrounded by four
regions of electron density?
Quantitative Problems
Section 7.1
38. How many valence electrons surround a stable
atom in a Lewis dot configuration?
39. Formaldehyde has the formula
CH2O and has one double bond
to allow all atoms to form stable
valence structures. Explain why
this initial attempt to come up
with the correct Lewis dot
structure will not work.
Section 7.2
40. Describe in your own words what Valence Shell
Electron Repulsion Theory is and what it is used
for.
41. A region of electron density is:
a. the electrons shared in a bond
b. the electrons that form the inner core of the
atom
c. the valence electrons that are not shared and
are called lone pairs
d. both A and C
42. Which number of regions of electron density will
give you which molecular shapes? Match the
correct number with the correct shape.
Regions of
electron density
2
3
4
Shape
tetrahedral
linear
trigonal planar
43. Describe the difference between trigonal planar
and trigonal pyramidal shapes.
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45. Polar covalent bonds form when electrons are
shared unevenly. Because electrons are not fully
transferred, this means they form partially
negative and partially positive charges on the
atoms. What color indicates a partial positive and
what color indicates a partial negative charge on
parts of a molecule?
46. Calculate the type of bond that will form between
the following pairs of atoms:
a. carbon and hydrogen
b. oxygen and hydrogen
c. lithium and bromine
d. carbon and oxygen
e. fluorine and fluorine
47. Bonds formed between atoms with an
electronegativity difference of less than 0.5 are
considered to be:
a. nonpolar covalent
b. polar covalent
c. ionic
48. Bonds formed between atoms with an
electronegativity difference between 0.5 and 2.1
are considered to be:
a. nonpolar covalent
b. polar covalent
c. ionic
49. Bonds formed between atoms with an
electronegativity difference greater than 2.1 are
considered to be:
a. nonpolar covalent
b. polar covalent
c. ionic
A NATURAL APPROACH TO CHEMISTRY
50. Most of the molecules in gasoline have the
formula C8H18. Is gasoline a polar or nonpolar
covalent molecule?
51. Some amino acids have polar side chains and
some do not. The tyrosine side chain is shown in
the green box in the drawing.
Section 7.3
56. Write the Lewis dot structure and structural
formula for PH3.
57. Write the Lewis dot structure and structural
formula for C2H6.
58. Write the Lewis dot structure and structural
formula for CBr4.
59. Write the Lewis dot structure and structural
formula for CH5N.
60. Write the Lewis dot structure and structural
formula for P2H2.
61. Write the Lewis dot structure and structural
formula for SiH3N.
a. Is the side chain for tyrosine polar or not?
b. If it is polar, are all of the bonds polar or just
some of them?
c. If just some of them are polar which bonds are
polar and which ones are not?
52. Describe the relationship between the electron
configuration of the noble gases and the charge
formed by specific ions. Give at least two
examples and use electron configurations to
support your answer.
62. Write the Lewis dot structure and structural
formula for HCP.
63. Write the Lewis dot structure and structural
formula for N2.
64. Write as many isomers of the Lewis dot structure
and structural formula for CH3ON as you can.
65. What is the Lewis dot structure and structural
formula for erythrose, shown here?
53. Describe the relationship between the electron
configuration of the noble gases and the number of
covalent bonds formed by specific atoms. Give at
least two examples and use electron
configurations to support your answer.
54. What is the most stable number of valence
electrons to have and how do you know?
55. What ratio of ions will be formed in the following
substances, so that the positively charged ions will
cancel out the negatively charged ions?
a. sodium sulfide (Na and S)
b. magnesium chloride (Mg and Cl)
c. aluminum oxide (Al and O)
66. For each of the previous problems (58, 59, 62, and
64) describe how many regions of electron density
exist for the carbon atoms in those molecules.
67. If part of a molecule is linear, what is the angle
formed between adjacent bonds?
68. If part of a molecule is trigonal planar, what is the
angle formed between adjacent bonds?
69. If part of a molecule is tetrahedral, what is the
angle formed between adjacent bonds?
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