Solution Manual For Organic Chemistry: A Guided Inquiry, 2nd Edition

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Stay on top of your textbook work with Solution Manual For Organic Chemistry: A Guided Inquiry, 2nd Edition, a guide offering complete solutions for every exercise.

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ChemActivity 1: Bond Angles and Shape
(What are the bond angles and shape of CH 4 ?)
Model 1: Planetary Model of an Atom
In a planetary model of an atom, negatively charged electrons (–1 each) are arranged around a
positively charged nucleus (+Z = nuclear charge) in a series of shells that look like orbits.
+1
= electron
+6 +7 +8 +9 +10
H C N O F Ne
shell #1
shell #2
Figure 1.1: Valence Shell Representations of Hydrogen, Carbon, Nitrogen, Oxygen, Fluorine, and Neon
core electrons = electrons in any inner shell(s) (don’t participate in bonding)
core atom = the nucleus (made up of protons and neutrons) plus the core electrons
valence electrons = electrons in the outermost shell (participate in bonding)
valence shell = outermost shell, where valence electrons are found
Electrons DO NOT “orbit” the nucleus like the planets orbit the sun. In ChemActivity 3 we will study a more complex model
in which electrons are described as inhabiting 3-dimensional regions of space called “orbitals” (1s, 2s, 2px, 2py, 2pz, 3s, etc.).
Critical Thinking Questions
1. (E) What does the number (+Z) at the center of each atom in Figure 1.1 represent, and what
number would you expect at the center of a representation of a bromine atom (Br)?
2. (E) How many total electrons does an oxygen atom have, and how could you find the answer to
this using a periodic table?
3. (E) How many valence electrons does each atom in Figure 1.1 have, and what number on a
periodic table gives you these answers?
4. What is the maximum number of electrons that can fit in…
a. (E) shell No. 1?
b. (E) shell No. 2 (Neon has a full Shell No. 2)?
c. Describe how the answers to a) and b) are contained in the structure of the periodic table.

ChemActivity 1: Bond Angles and Shape 9
Model 2: Bonding and Non-bonding Electron Domains
Bonding electron domain = shared valence electrons (2, 4, or 6e—) localized between two core atoms
3 types Æ Single Bond (1 pair, 2 electrons); Double Bond (2 pairs, 4 e—); Triple Bond (3 pairs, 6 e—)
Non-bonding electron domain (“lone pair”) = pair of valence electrons (2 e—) not involved in a bond
One way to think of a bond: two positively charged core atoms mutually attracted to the negatively
charged electrons that are localized between them.
+8valence shell
representation*
dot shorthand
electron domain
bond-line
ball & stick
space filling
H H O O N N
+1 +1 +8 +7+7
O O N NH H
H H O O N N
Figure 1.2: Example of a Single, Double, and Triple Bond
Critical Thinking Questions
5. (E) How many electrons are in a triple bond?
6. (E) Identify each lone pair shown in the first four rows of Figure 1.2.
7. (E) Each molecule in Figure 1.2 has exactly one bonding electron domain. Identify it and…
a. label what type of bonding electron domain it is.
b. report the number of electrons in each bonding electron domain.
8. You hear a student from a nearby group say that “Electron domains repel one another.” Cite
evidence from Figure 1.2 to support or refute this statement.

10 ChemActivity 1: Bond Angles and Shape
Model 3: Bond Angles
bond angle = angle defined by any three atoms in a molecule. (e.g., ∠HBH in BH3, below)
According to the Valence Shell Electron Pair Repulsion model (VSEPR) electron domains spread out
as far as possible from one another (repel one another).
Even for a molecule with different-sized electron domains (e.g., H2CO), bond angles remain very close
to those you would expect if all the electron domains were the same size.
Table 1.1: Bond Angles of Selected Molecules
Bond-line Structure Approximate Bond Angle
H—Be—H N C H O C O 180
B
H
H H
C
H H
C
O
H H
C
H
H H
120
Critical Thinking Questions
9. Use VSEPR to explain why the ∠HBH bond angle of BH3 is 120o. (Hint: What is one-third of
360o?)
10. Both the ∠HCH and ∠HCO bond angles of H2CO (formaldehyde) are very close to 120o, but one
is slightly smaller than the other. Predict which is smaller, and explain your reasoning.
11. Use VSEPR to assign a value of “close to 180o” or “close to 120o” to each bond angle marked
with a dotted line. (These angles are drawn as either 90o or 180o, but may be another value.)
H C C C
NH H
C C CH
H H
H N C O
H
H C H
O H
C C C
H
H
H H
H
12. Consider the following flat drawing of methane (CH4).
a. What is ∠HCH bond angle implied by this drawing if you assume it is flat?
b. Are the electron domains of this flat CH4 spread out as much as possible?
C
H
H
HH
c. Use model materials to make a model of CH4 (methane). If you assembled it correctly, the
four bonds (bonding electron domains) of your model will be 109.5o apart.
d. In which representation, the drawing above or the model in your hand (circle one) are the
H’s of CH4 more spread out around the central carbon?

ChemActivity 1: Bond Angles and Shape 11
e. Confirm that your model looks like the following drawing. The wedge
bond represents a bond coming out of the page, and the dash bond
represents a bond going into the page.
H
C
H H
H
f. You will often see methane drawn as if it were flat (like on the previous page). Why is this
misleading, and what is left to the viewer’s imagination when looking at such a drawing?
13. Use VSEPR to assign a value of (close to) 109.5, 180 or 120 to each marked bond angle.
C C C
H
H
H
H N
H
C
H
C
H
H
O
O H H N C
H
O C
H
H
H
H
H
H O C H
H
H
H
14. A student draws the picture of ammonia (NH3) in the box below, left, and predicts it will be a flat
molecule with ∠HNH bond angles of exactly 120o. Unfortunately, the student left something out.
N
H
H H
120 o
N
H H
H
lone pair in non-bonding domain
Student's incorrect drawing
Accurate Representation of Ammonia
a. What did the student omit from his drawing?
b. What is the actual ∠HNH bond angle of ammonia (based on the drawing above, right)?
c. Explain why water, ammonia, and methane (shown below) all have about the same bond
angles (close to 109.5o) even though they have different numbers of bonds.
H
O
H
N
H H
H
H
C
H H
H
water ammonia methane
Memorization Task 1.1: Correlation between #No. of Electron Domains and Bond Angle
Electron Domains (bonding + non-bonding) Bond Angle Close To… Examples
4 109.5o CH4, NH3, H2O, H4N+
3 120o BH3, H3C+, CH2, H2CO
2 180o BeH2, O2N+, CO2, HCN

12 ChemActivity 1: Bond Angles and Shape
Model 4: Shape
A central atom = an atom bonded to two or more other atoms. (e.g., Oxygen in H—O—H)
Each central atom has a shape determined by the arrangement of the atoms attached to it.
Memorization Task 1.2: The five molecular shapes we will encounter in this course
H
C
H
H
H
tetrahedral
N
H
H
H
C atom in
the center
is not a
corner
pyramidal
O C O
linear
H
O
H
bent
C
HO
H
trigonal
planar
dotted line (---) shows
edges that are NOT bonds
Critical Thinking Questions
15. (E) Explain why the molecule H—F is not associated with an official shape as defined in Model 4.
16. How many central atoms does the molecule H2NCH3 have, and what is the shape about each?
N
H
H C
H
H
H
17. Indicate the bond angle and shape about each central atom.
C
H
H F O
H
H C CH H O O HH N
H
H
HHH
H
H
C
H
H H
18. Explain how there can be two kinds of bent: “bent-109.5o” and “bent-120o,” and give an example
of each from the previous question. (Note that “bent-109.5o” is more common than “bent-120o.”)
19. A student makes the following statement: “The shape of water is tetrahedral because the four
electron pairs about oxygen are approximately 109.5o apart and point to the corners of a
tetrahedron.” What misconception does this statement convey?
20. A student who missed this class needs to know how to predict the bond angles and shape of a
molecule from looking at its bond-line representation. Write a concise but complete explanation
for this student.

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ChemActivity 1: Bond Angles and Shape 13
Exercises
1. Draw a valence shell representation of a
a. Helium atom.
b. Sulfur atom.
2. In the box, draw a bond-line representation of the molecule shown on the left. Be sure to include
only valence electrons (either as line bonds or lone pairs).
+8
+1+6
+7
+1
+1
+1
line-bond rep.
+1
+1
3. Consider the incomplete valence shell representation below.
a. Assume the atom is neutral, and write the correct nuclear charge at the center of the atom.
b. What is the identity of this neutral atom?
4. How many valence electrons does a neutral
a. K atom have?
b. C atom? N atom? O atom?
5. Consider the molecules AlCl 3 (aluminum chloride) and CF4 (carbon tetrafluoride).
a. Draw the valence shell representation of each.
b. Predict the value of the XYX bond angle, and explain your reasoning.
6. Draw an example of a bent molecule with a bond angle of near 109.5 o; then draw a different bent
molecule with a bond angle of about 120 o.

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14 ChemActivity 1: Bond Angles and Shape
7. Label each atom marked with an arrow with the appropriate shape name, and estimate the bond
angles around it as being close to one of 180o, 120o, 109.5o or 90o. (circled charges indicate the
charge on the molecule or fragment)
C
H
H F O
H
H H NH H O O HH N
H
H
HH
8. Make a model of each of the following molecules:
H O Cl Cl C Cl
Cl
H
H O Cl H
H
Cl C Cl
Cl
H
H C O
H
H
H
HN
C
O
OH H
C CH H
H H
C
C C
C
C
H
H
H
H
HH
H
H
H
H
C C C O
H H
H
H
H
H
H
O
a. Based on your model, draw a bond-line representation with as many atoms as possible in the
plane of the paper. Use wedge and dash bonds to represent any atoms that do not lie in the
plane of the paper.
b. Indicate each unique bond angle and the shape of each unique central atom.
9. Read the assigned pages in the text, and do the assigned problems.
Using “The Big Picture” & “Common Points of Confusion” Sections
It can be fun to “discover” your own answers as you are asked to do in this workbook, but…
How do you know if your “discovered” understanding is valid?
The answer is: Even practicing scientists and professors never know for sure that they are correct. In
some ways, deciding if you are right is the hardest part of being a scientist. A practicing scientist cannot
“check the key” to see if her new theory is correct! In this course and in real life you must constantly
test and improve your current understanding by applying it to problems and discussing it with peers.
One goal of this course is to develop the ability to recognize the signs when you are correct; and equally
importantly, recognize the signs when you are missing something critical.
After completing a ChemActivity, one way to start checking your understanding is to read the Big
Picture and Common Points of Confusion sections (examples on the next page). If the homework or
these sections do not make sense to you then you are missing something important. You need to go
back and study the activity, do more problems, read the textbook more closely, or seek help from a
peer, teaching assistant or your instructor. (More advice about how to know if you are “learning the
right thing” can be found in the “To the Student” section the precedes the Table of Contents for this
book, and the Frequently Asked Questions section of the IntroActivity that precedes this chapter.)

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ChemActivity 1: Bond Angles and Shape 15
The Big Picture
After this week you will rarely be asked to report a bond angle or shape. Yet it is critical that you be
able to do both. Doing well in organic chemistry largely depends on your ability to see molecules as
three-dimensional objects. The electrons of most every central atom you will encounter are arranged
109.5 o or 120 o apart (180o arrangements are quite rare), but it gets complex when you are expected to
see a molecule with multiple central atoms in 3D. The purpose of this activity is to get you started
thinking about tetrahedral and trigonal planar geometries. If you do not already have a model set,
borrow or purchase one. You will need it for the first half of this course while you are “programming”
your brain to see the two-dimensional drawings on the page as 3D objects.
Common Points of Confusion
At the end of each chapter you will find a brief explanation of common student misconceptions. This
section may be useful if a homework problem does not make sense or as final preparation for a quiz.
• When asked the question, “What is the shape of water?” students sometimes answer “tetrahedral”
because they know that the four electron domains of water spread out into a tetrahedral-type
pattern. However, the answer is “bent” because shape is determined by the location of the atoms.
Similarly, it makes no sense to ask what is the bond angle between the two lone pairs of water—
there should be approximately 109.5o between these lone pairs, but this is difficult to measure.
• Some students mistakenly assume they are expected to predict the EXACT bond angle of a given
molecule. Generally, you are expected to predict only if the bond angle is closest to 180, 120, or
109.5 degrees.

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ChemActivity 2: Lewis Structures
(How do I draw a legitimate Lewis structure?)
Model 1: G. N. Lewis’ Octet Rule
In the early part of the last century, a chemist at the University of California at Berkeley named Gilbert
N. Lewis devised a system for diagramming atoms and molecules. Though simple, the system is still
used today because predictions made from these diagrams often match those based on experiment.
Lewis proposed the following representations for the first ten elements with their valence electrons.
H
Li Be B C N O F Ne
He
Figure 2.1: Electron Dot Representations of Elements
Only He and Ne are found in nature as shown above. All the other elements are found either as a
charged species (ion) or as part of a molecule that can be represented as a legitimate Lewis structure.
CHECKLIST: a Legitimate Lewis Structure is a dot or line bond representation in which…
I. The correct TOTAL number of valence electrons is shown.
II. The sum of the valence electrons around each hydrogen atom is two.
III. The sum of the valence electrons (bonding pairs + lone pairs) around each carbon, nitrogen,
oxygen, or fluorine atom is eight–an octet. (this is the “octet rule”)
Note that Lewis’ rules apply to H, C, N, O and F. We will find that atoms in the next row of the periodic table (e.g., silicon,
phosphorus, and sulfur) and beyond commonly violate the octet rule.
O H
H
O H
H
O H =
"lone pair"
"bonding
pair"
we will usually use
line-bond structuresoxygen atom +
O =
oxygen atom + 2 electrons
O
-2
water
oxygen -2 anion
H
F F =
2 fluorine atoms
F
F 2 (fluorine gas)
F F F
-1 -1
2 hydrogen
atoms
Figure 2.2: Examples of combinations that form legitimate Lewis structures
Read this page once, and begin answering the Critical Thinking Questions on the next page.

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ChemActivity 2: Lewis Structures 17
+8 +1+1+7 +1+1
+1
+9 +1
N
H
HH O HH F
Shell
Lewis H
Figure 2.3: Valence Shell and Lewis Representations of Selected Compounds
Critical Thinking Questions
1. (E) Confirm that each molecule or ion in Figures 2.2 and 2.3 is a legitimate Lewis structure.
2. The valence shell of an atom in a legitimate Lewis structure (see Figure 2.3) has what in common
with the valence shell of a noble gas? (Noble gases are stable elements found in the last column of
the periodic table, e.g., He, Ne, Ar, etc.)
3. Draw a shell representation and Lewis structure for
the ion of fluorine that you predict is most likely to
be stable, and explain your reasoning.
4. Draw a Lewis structure of a neutral molecule that
you expect to be a stable and naturally occurring
combination of one carbon atom and some number
of fluorine atoms.
5. The following structure is NOT a legitimate Lewis structure of a neutral O2 molecule.
a. Explain why it is not legitimate.
O O
b. Which item on the legitimate Lewis structure CHECKLIST in Model 1 is violated?
6. It is impossible to draw a legitimate Lewis structure of a neutral NH4 molecule. Hypothetically,
how many valence electrons would such a neutral NH4 molecule have if it could exist?
a. The +1 cation, NH4+, does exist. How many valence electrons does one NH4+ ion have?
b. Draw the Lewis structure for NH4+
7. Describe how to calculate the total number of valence electrons in a +1 ion, in a -1 ion.

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18 ChemActivity 2: Lewis Structures
Model 2: Two Lewis Structures for CO2
O C O O C O
I II
Experiments indicate that both carbon-oxygen bonds of
carbon dioxide (CO2 ) are identical.
Critical Thinking Questions
8. (E) Are both structures of carbon dioxide (CO2) in Model 2 legitimate Lewis structures?
9. (E) Which Lewis structure best fits experiments indicating that both C to O bonds are identical?
Model 3: Formal Charge
One of the Lewis structures of CO2 in Model 2 is less favored because it has an imbalance of charge. To
find the “hot spots” of + and – charge in a structure we must calculate the formal charge of each atom.
Memorization Task 2.1: Formal Charge = (Group Number) – (no. lines) – (no. dots)
• Group Number = Column number on the periodic table (or number of dots on atom in Fig. 2.1)
• No. lines = Number of line bonds to the atom in the structure
• No. dots = Number of non-bonded electrons on an atom in the structure
Critical Thinking Questions
10. (E) According to the periodic table at the end of this book, what is the Group Number of nitrogen?
a. (Check your work.) Does this match the number of dots on N in Fig. 2.1?
b. (E) How many line bonds are attached to N on the structure of NH3?
c. (E) How many nonbonded electrons are drawn on N in NH3?
d. Calculate the formal charge of each atom in NH3? N HH
H
11. (Check your work.) Most atoms in organic
molecules (including all atoms of NH3) have
a zero formal charge. Confirm that each
atom at right has a zero formal charge.
C
H
H H
H OH
H
BrH
H
C
H
O
C NH
12. In this course we will often encounter +1 and -1 formal charges, though rarely will we see formal
charges of +2, -2, +3, -3, etc., because they are generally unfavorable. By convention, only nonzero
formal charges are shown on a structure. Plus 1 and minus 1 formal charges are shown as a + or
–or as a circled + or – ( / ). Confirm the formal charge assignments below.
NC
H
H
H
H
O HC
H
H C
H
CH
C C C
H
C
H
H
H
H
H
Cl

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ChemActivity 2: Lewis Structures 19
13. A complete Lewis structure must show all nonzero formal charges. Complete each of the
following Lewis structures by adding any missing formal charges.
H C
O
O
S H
H
O O
H
H
N H
H
P C
H
H
C CC
H
OC
O
CN
N C O H3 C C O
C
H N N
H
HH
H
H
H
H
H
H
H
H
H
H
H
C
H
H
H
H
H
H H
14. Net Charge = total charge on a molecule. (Check your work.) Structures in the top row of the
previous question have a net charge of +1, structures in the middle row have a net charge of zero,
and structures in the bottom row have a net charge of -1.
15. T or F: The sum of the formal charges on a Lewis structure is equal to the net charge on the
molecule or ion. (If false, give an example from CTQ 13 that demonstrates this is false.)
16. T or F: If the net charge on a molecule is zero, the formal charge on every atom in the molecule
must equal zero. (If false, give an example from CTQ 13 that demonstrates this is false.)
17. Identify the one Lewis structure in CTQ 13 that is NOT legitimate, and explain what attribute of a
legitimate Lewis structure it is missing.
(Check your work.) The top-center Lewis structure in CTQ 13 is a key exception to the octet rule
called a carbocation. We will study carbocations extensively in the course. For reasons we will discuss
later, a carbocation carbon rarely is involved in a double or triple bond. That is, a carbocation almost
always has three single bonds, as shown on the next page.
Model 4: Condensed Structures and Using R, X, & Z as Placeholders
R, X, and Z are not elements but placeholders for atoms or groups of atoms. For example, ethanol…
C
H
O C
H
H H
H
H
can be
written
as...
ROH where... R = C
H
C
H
H
H
H
HOCH2 CH3 HOR -CH2 CH3
ethanol an "ethyl" groupgeneral structure for an alcohol
Condensed Structures
give just enough
information to draw the
structure

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20 ChemActivity 2: Lewis Structures
Memorization Task 2.2: Conventions for the use of R, X, and Z as placeholders
• R is used to represent H or an alkyl group. An alkyl group is a straight or branched chain
made from C and H atoms with formula CnHm e.g., -CH3, -CH2CH3, -C(CH3)3, etc.
• X is used to represent F, Cl, Br, or I (the common “halogens”)
• Z will be used to represent any atom or group of atoms
If the identity of R, X, or Z is not specified, assume a wide range of legal identities are possible.
For example: in the table below, the formal charges shown hold true regardless of the identities of Z.
Model 5: Recognizing Formal Charges for C, N, O, and X
+1 0 −1
C
Z
ZZ
Note: The two other ways to draw a
carbocation are very uncommon.
C
Z
Z
Z
Z
Z C Z
Z
C ZZ
Z C Z C
Z
Z
Z
Z C
Z
C
ZC
N
Z N Z Z N Z
four ways (draw the two that are missing)
draw three ways draw two ways
O
draw three ways
draw two ways draw one way
Z X Z X Z
two ways (less common)
draw one way draw one way (an anionic atom)
X
Critical Thinking Questions
18. Complete the box in Model 5 for N+1, by drawing the other two ways an N can carry a +1 formal
charge. (Hint: These two structures should have molecular formulas +NZ4, and +NZ3, respectively.)
19. Complete the rest of the table for N, O or X by drawing the number of Lewis structures specified.
20. For a legitimate Lewis structure…
a. What is the formal charge of any nitrogen with four bonds? … three bonds? … two bonds?
b. What is the formal charge of any oxygen with three bonds? … two bonds? … one bond?
c. What is the formal charge of any halogen (X) with two bonds? … one bond? … zero bonds?
21. Parts a-c in the previous CTQ mean you can quickly recognize the formal charge on an N, O and X
without considering non-bonded electrons (dots). Explain why a similar statement equating formal
charge and number of bonds DOES NOT WORK for carbon (i.e., you have to look for the dots).

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ChemActivity 2: Lewis Structures 21
Memorization in Organic Chemistry:
Memorization is a small but important part of learning organic chemistry. Memorization Tasks such as
the one below will be clearly marked throughout this book. This is done to encourage you to memorize
the critical bits of information in these boxes AND help you realize that you can derive everything
outside of these boxes from the key concepts in the ChemActivity.
Memorization Task 2.3: Corrleation between No. of bonds and formal charge for C, N, O, X
Before the next class: Study the patterns in Model 5, and do practice problems until you can QUICKLY
recognize the formal charge (+1, 0, or -1) of any C, N, O or X in a structure without counting.
For example, an N with four bonds should look “wrong” without a +1 formal charge; and an N with two
bonds should look “wrong” without a -1 formal charge. (Write a similar rule for oxygen!)
Note that for carbon you must count the number of bonds AND check whether there is a lone pair on C.
That is, a C with three bonds and no lone pair should look “wrong” without +1 formal charge; and a C
with three bonds and one lone pair should look “wrong” without a -1 formal charge.
Exercises
1. Make a checklist that can be used to determine if a Lewis structure is correct and that it is the best
Lewis structure.
2. Turn back to Model 2, and add any missing formal charges to each Lewis structure of CO2.
a. Based on the concept of formal charge, which is the better Lewis structure for CO 2 (in
Model 2), Lewis structure I or Lewis structure II? Circle one, and explain your reasoning.
b. Is your choice consistent with the experimental data?
3. Shown below are two possible Lewis structures for the amino acid called glycine.
H N
H
C C
H
H O
O H
Structure I
H N
H
C C
H
H O
O H
Structure II
a. Predict the ∠COH bond angle based on the Lewis structure on the left.
b. Predict the ∠COH bond angle based on the Lewis structure on the right.
c. Which prediction do you expect to be more accurate? Explain your reasoning.
4. Draw the Lewis structure of a neutral molecule that is a naturally occurring combination of
hydrogen atoms and one sulfur atom. What is the shape of this molecule?

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22 ChemActivity 2: Lewis Structures
5. Draw legitimate Lewis structures of the following species, and predict the geometry about the
central atom (shape).
+
b. NO c. Na. NH3 2 2O (try with N or O
as the central atom)2−
d. CCl e. CO4 3
f. N2 (Note: based on the definition, a molecule with only two atoms does not have a shape.)
6. For each element, predict (and draw a Lewis structure of) the most commonly occurring ion
(some of these have a charge greater than +/- 1)
a. sulfur c. magnesium
b. iodine d. oxygen
7. Predict which of the following species is least likely to exist.
+ −
NOCH HO2
2
−
8. The molecules BH3 and SF6 and the ion SO4 exist and are stable. Draw a Lewis structure of each,
and comment on whether they are exceptions to Lewis’ octet rule.
9. These are NOT
legitimate Lewis
structures (and are
missing formal
charges). Show (as in
the example) how a
pair of electrons can
be moved to make the
Lewis structure
legitimate.
N
O
O
O
H
(curved arrow shows
where the electron pair
was moved from and to)
legitimate Lewis structure
H N C H
H O
N N C H
HH
H
F O O H
H N C H
H O
10. Fill in missing formal charges where needed (all lone pairs are shown).
a) b) c)
NO O
O
C
H
HH
d) e) f)
N
H
H
HH CH
H
H
C
O
O H N N N
H
CS S
CS 2 CH3–
NO 3–
NH4+ CH 3COOH HN3

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