1
CHEMISTRY 313
Tuesday, Thursday 9:00 – 10:15 a.m.
Exploratory Hall L004
Spring 2016
Prof. S. W. Slayden
333 Planetary Hall
sslayden@gmu.edu
703-993-1071
http://mason.gmu.edu/~sslayden
Office Hours: 10:30 – 11:30 a.m.
and other times by appointment
Week of
Text Chapter
Klein Chapter
Jan. 18
1
Introduction
1, 2
1, 2
2
2, 4
8
3; (11.2, 11.6)
3
15
4; 2.4, 10.3, 10.3A,B, 10.5A
5
22
4
6
29
5
7
25
Feb. 1
Mar. 7
(Add deadline 1/26)
Feb. 16
(Drop deadline 2/19)
SPRING BREAK
14
5
21
3.1, 6
8.1-8.5
28
6
9, 12.3
Apr. 4
6
10.7-10.9
11
7
10.1-10.6
18
8
25
11
11 (except 11.5,
11.10), 8.6-8.7
12.1, 12.2, 12.4,
12.8, 12.10
May 2
Exam / Note
Mar. 22
Apr. 12
May 5, 8:30 a.m.
Note the nonstandard time!
2
Course Materials
Text
"Organic Chemistry, 11th Ed.", Solomons and Fryhle
Study Guide/Solutions Manual
To accompany the text
Lecture Supplement
"Chemistry 313 Supplement", available on BlackBoard and
in the Bookstore (by request)
Molecular Models
Two brands available in the bookstore and many more at
amazon.com – choose one set
Workbook
"Organic Chemistry as a Second Language", D. R. Klein
(any edition)
Scantrons
four; Form No. 882-E or comparable
All course materials can be purchased at the GMU campus bookstore. None of these materials is
required. However, you will probably be more successful in the course if you read and study the
textbook, attend class and take notes on the Lecture Supplement, and take advantage of the other
suggested resources as necessary to help you understand organic chemistry.
Resources
This syllabus and all other information pertaining to the course (such as text sections to be
read/omitted, text errors, suggested problems) appear on the course web page (accessed through
http://mason.gmu.edu/~sslayden). The on-line syllabus is the official syllabus for the course.
Occasionally, announcements must be made to the entire class via e-mail. The e-mail will be sent
to your GMU e-mail account.
Copies of exams and the answer keys from a previous semester are available on Blackboard. In
the library are textbooks from which you can supplement your reading and problems. Other
chemistry resources are available on the internet, some of which can be accessed through my
Web site.
You can record the lecture if you wish, but please don’t put the recording device on my desk or
where I can see it.
In the Classroom
You are expected to be on time for class. If you must enter class late, do so quickly and quietly.
You are expected to remain in class for the entire class period. Walking in and out of class is
distracting to the rest of us. It is a courtesy, if you do have to miss class or leave early, to contact
the professor beforehand and to leave as unobtrusively as possible. If you miss a class, you are
expected to get notes, etc. from other students.
3
Focus on class material during class time. Sleeping, talking to others, doing work for another
class, reading a newspaper or using an electronic device are unacceptable and disruptive. Cell
phones, pagers, and other electronic communication devices are not allowed in this class. Please
keep them stowed away and out of sight. Do not eat or drink in the classroom. It is difficult, if
not impossible, to take notes or interact effectively while eating. Do not close your books, rustle
your papers, or walk out of the classroom before class is dismissed. These actions are a
distraction and communicate a lack of respect.
Enthusiasm is contagious. Professors respond and teach better to an alert, attentive, and
interested class.
Examinations and Grades
Course Grade: 4 Exams --three exams given during the semester; the fourth exam is given
during final exam week. Each exam is graded as the percentage of questions you answered
correctly. The highest grade possible is thus 100%. The four exam grades are counted equally
and the average of the exams is your final numerical course grade.
Unexcused absence from an exam will result in an assigned grade of zero. If you are unable to
attend an exam, you are expected to talk to me before the exam or as soon thereafter as possible.
(E-mail and voice mail do not satisfy this requirement, although you may leave a message to
alert me.) Documentation for your absence may be required before any consideration is given for
taking a make-up exam, usually during the same week as the originally scheduled exam. Excuses
for “last minute” occurrences that interfered with studying before the exam or that interfered
with attending the exam will not be accepted. Allow enough time for commuting on exam days –
no extra time will be allowed for tardiness, unless there was some catastrophe.
The course web site contains a page of Instructions and Frequently Asked Questions for
taking exams in this course. Please read it.
If classes are canceled by the University on a scheduled exam day, the exam will be given at the
next scheduled class period after classes resume. Hour exams will not be rescheduled (unless
classes have been canceled for an extended period and I notify you). If I am absent for an exam,
another Chemistry Department instructor will take my place. University Information is 703-9931000.
Make-up Periods for Final Exams
If the university is unexpectedly closed on an exam day, make-up times will be announced as
soon as they are determined and will be posted on the University-wide Class Cancellations page.
The last scheduled day of the final exam period is the scheduled make-up day. The Friday,
Saturday, and Sunday during the final exam period are also potential make-up days. Students and
faculty must be available for the make-up day(s).
4
Final letter grades are determined after all numerical grades for all students are calculated
at the end of the semester. Letter grades are not assigned for individual exams. You may
tentatively assume that your exam grade corresponds to a C if it is within approximately +7
points of the number that I define as mid-C for the exam (see the BlackBoard announcement
after each exam). Generally, grades below ~40% are failing grades (F) and grades above 90% are
excellent (A). When each exam is returned, the mid-C and the high and low grades will be
announced on BlackBoard so that you can estimate your standing in the class. The mid-C at the
end of the semester will be the average of the mid-C grades on the exams so you can estimate
your final grade based on your on-going average. The class website link shows a typical grade
distribution chart.
There are no extra credit assignments available for the course. Please do not ask about “curving
grades” if by that you mean some arbitrary number of points might be added to your grade in
order to raise it.
Disability Services
If you are a student with a disability and you need academic accommodations, please see me
after contacting the Disability Resource Center (DRC) at 703-993-2474. All arrangements for
academic accommodations must be initiated through that office.
Honor Code
It is the responsibility of all students to be familiar with the GMU Honor Code. All examinations
are closed book, and the use of notes or other written material is not permitted. A periodic table
will be provided with the exam. You may use, but not share, molecular models during any exam.
You may not use or have in your immediate presence any electronic device.
Academic Policies
The University Catalog is the source of information about all academic policies:
http://catalog.gmu.edu/
5
STEPS YOU CAN TAKE TO MAXIMIZE YOUR PERFORMANCE
IN ORGANIC CHEMISTRY
The textbook authors, Solomons, Fryhle and Snyder, in their foreword to
the 11th edition "To the Student" on p. xxxvi, present some excellent
suggestions for studying organic chemistry. If you follow those
suggestions and develop your study habits accordingly, you are likely to
be more successful in the course. In short, these are their suggestions:
1. Keep up with your work from day to day - never let yourself get
behind.
2. Study material in small units, and be sure that you understand each
new section before you go on to the next.
3. Work all of the in-chapter and assigned problems.
4. Write when you study. (This is one of the most comprehensive and
most important of all the suggestions.)
5. Learn by teaching and explaining.
6. Use the answers to the problems in the Study Guide in the proper
way.
7. Use molecular models when you study.
6
INTRODUCTION
H
heat
H
N
H
N
C
O
H
O
N
C
H
N
H
H
H
ammonium cyanate
urea
“I can produce urea, without having the need for kidneys or an animal at all, be it
human being or a dog. The cyanogen acidic ammonia is urea.”
Letter from F. Wöhler to J. Berzelius, 1828
ISOMERS Isomers have identical molecular formulas but have different arrangements of atoms
in their molecules, that is, different structures.
Constitutional Isomers (Structural Isomers) differ in their "connectivity", that is, in the order
of attachment of the atoms in the molecule. These isomers have different physical
and chemical properties.
7
CH3
Structure
CH3 O
CH3
CH3 CH2 O
H
CH3 CH2 CH2 CH3
CH3 C
CH3
H
Name
dimethyl ether
ethanol
n-butane
isobutane
b.pt. (oC)
−24.9
78.5
−0.5
−12
density (g/ml)
0.661
0.789
0.601
0.603
The properties of a substance depend upon its structure.
structure  physical & chemical properties
Structure
Chemical
transformations
Bonding
Electrons
“Sugar, salicin, and morphium will be produced artificially. Of course, we do not know the
way yet by which the end result may be reached since the prerequisite links are unknown to
us from which these materials will develop—however, we will get to know them.”
F. Wöhler and J. von Liebig
8
ELECTRONIC STRUCTURE OF ATOMS
Atom consists of a dense inner core -- the nucleus (protons, neutrons) -- and electrons that
surround the nucleus.
Particle
Relative Mass
Charge
proton (p)
neutron (n)
electron (e–)
1
+1
0
–1
1
1/1821
The number of protons in the nucleus (atomic number, Z) uniquely determines the atom's
identity.
In a neutral atom, #p = #e; #n is variable (atomic isotopes).
{ 12C6 13C6 14C6 }
isotopes of carbon
{1H1 2H1  D (deuterium) }
isotopes of hydrogen
Electrons have particle-like properties (mass) and wave-like properties (diffraction).
Solve
--------------->
WAVE EQUATION
WAVE FUNCTIONS (and quantum numbers)
Wave Functions express the energies and the positions of electrons in an atom.
Quantum Numbers
n is the designation for the principle, or main, energy level (or energy shell) in the atom that may
be populated with electrons.
The closer an electron is to the nucleus, the less energy it has and the more stable it is.
l is an energy sublevel (or subshell) within a given main energy level, n. [ s p d f sublevels].
There are restrictions on the number of sublevels within a given main level.
ml is an orbital within a given sublevel. An orbital is a region of space in which an electron is
most likely to be found. [ one s three p five d seven f ] Orbitals have unique "shapes".
Maximum of two electrons per orbital.
ms designates the quantum number for the electron spin. The opposite nature of their spins
differentiates between the two electrons that may occupy a given orbital (thus having the
same n, l, and ml). Spins may be designated as +1/2, –1/2; clockwise, counterclockwise; 
9
Pauli Exclusion Principle Each electron in an atom has a unique set of quantum numbers
(n, l, ml, ms)
Main level
(n)
Sublevel
(l)
Orbital
( ml )
Spin
( ms )
Max. #
electrons per
sublevel
1
s
s
± 1/2
2
Max. # electrons
per main level
2
2
s
s
± 1/2
2
p
pxpypz
± 1/2
6
8
3
s
p
d
s
pxpypz
five d's
± 1/2
± 1/2
± 1/2
2
6
10
18
4
s
s
± 1/2
2
p
d
f
pxpypz
five d's
seven f's
± 1/2
± 1/2
± 1/2
6
10
14
32
5
6
7
*
*
*
* Although each main energy level contains a number of sublevels
equal to n (verify for n = 1-4, above), main levels 5-7 do not
require more than 4 sublevels to accommodate the electrons in all
known atoms (due to sublevel energy overlap). The sublevels,
orbitals, and electron-occupancy for n = 5-7 will be the same as for
n = 4.
Electron Configuration is the description of the distribution of electrons within an atom. In the
ground state, electrons will occupy the atomic orbitals of lowest energy.
n l#e– or n ml #e–
H1 = 1s1
Li3 = 1s22s1
B5 = 1s22s22p1
10
11
Hund's Rule In filling degenerate orbitals with electrons, each orbital is occupied singly with
electrons of the same spin; subsequently, electrons of opposite spin are added.
ENERGY LEVEL DIAGRAM
2px
2py
2pz
2s
Inc.
Energy
1s
SUGGESTED PROBLEMS:
1. Write electron configurations for all atoms Z = 1 - 18.
2. Select a few atoms from problem 1. and show the application of Hund’s Rule by drawing
energy diagrams as shown above.
Valence Electrons are those electrons that occupy the outermost (highest) main energy level (n)
of an atom. They are the most loosely held and engage in chemical reactions.
Electron Dot Symbols designate the valence electrons (dots) surrounding the inner core
electrons and nucleus (atomic symbol) of an atom.
MAIN BLOCK ATOMS
I
II
III
IV
V
VI
VII
VIII
A
A
A
A
A
A
A
A
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Rule of Electronic Stability
Electronic stability is greatest when atoms have an n(s2p6) configuration in their valence shell
and resemble the closest Noble Gas configuration. The exceptions are elements that can attain
the 1s2 configuration more easily than the s2p6 configuration, such as H and Li.
Atoms gain, lose, or share electrons in order to achieve this "octet" of electrons.
Mg = 1s22s22p63s2

Cl = 1s22s22p63s23p5

+2
Mg = 1s22s22p6
Cl −1 = 1s22s22p63s23p6
Electronegativity is the tendency of an atom to attract electrons toward itself (to the positively
charged nucleus). E.N. is a periodic property that increases up a family and
increases across a row in the periodic table.
Electronegativity
B < H < C
Electronegativity increases across a row
 The positive charge on the nucleus increases as protons are added to the nucleus.
 Negatively charged valence electrons are added to the same main energy level.
 The effective positive charge of the nucleus increases for all the electrons that are in the
same valence shell and so the attraction between the electrons and the nucleus increases
across the row.
Electronegativity decreases down a column
 The positive charge on the nucleus increases as protons are added to the nucleus.
 Negatively charged valence electrons are added to main energy levels further from the
nucleus.
 Although there is an increase in the number of positive protons in the nucleus, the
increased number of inner electrons screens the effective positive charge on the nucleus
from attracting the outer valence electrons.
13
CHEMICAL BONDING
Chemical bond -- the force of attraction that holds atoms together in a molecule.
The force is the attraction between the negatively charged electrons and the
positively charged nuclei. The drive for bond formation is to achieve a filled
valence shell.
Ionic bond – Results from the electrostatic attraction between ions of opposite
charge.
K
K
+
e-
F
+
e-
F
Covalent Bond -- results from sharing of two electrons between two atoms. The
shared electrons are electrostatically attracted to both nuclei.
There are several ways to represent covalent bonding between atoms in a molecule.
Before going into detail about how to represent bonding using the atomic orbitals
of atoms, we will construct structural formulas using the same electron dots as for
individual atoms.
WRITING ELECTRON DOT STRUCTURES
(1) From the molecular formula, calculate the number of valence electrons
contributed by all the atoms in the molecule. Adjust the total number of
electrons for net charge, if any.
(2) Arrange the atomic symbols to correspond with the known connectivity.
Generally, more electronegative atoms and H's tend to be terminal atoms in a
molecule. Ignore molecular geometry for now.
(3) Distribute the valence electrons in pairs so that each atom is bonded to at least
one other atom. Distribute the remaining valence electrons among the atoms,
giving them multiple bonds and/or non-bonded electrons. Try to give each atom
an octet (except H).
14
(4) Check your drawing to make sure that all atoms have no more electrons than
their valence level can accommodate: a first row element can have no more than
two; elements in the second row of the periodic chart can have no more than
eight electrons; elements in the third and higher rows may have more than eight
electrons because of the availability of the unfilled d sublevel.
P and S are notable examples of atoms that may have more than eight electrons in
their valence shells – they are in the 3rd main energy level.
Example: C2H4O2
Common bonding for atoms without formal charges:
I
II
III
IV
V
VI
VII
A
A
A
A
A
A
A
A
A
A
A
A
A
SUGGESTED PROBLEM: There are 7 structural isomers with the molecular formula C2H4O2 that
contain one double bond. Draw as many of them as you can. Include all non-bonding electrons in
your drawings. You can find the answers at chemspider.com.
15
Formal Charge on an atom is a method of electron bookkeeping that indicates the difference in
electron "ownership" by an atom in its atomic vs. its bonded molecular state. It does not indicate
the actual charge distributions on the atoms in a molecule.
(1) Find the atom's valence Group number (I-VIII).
(2) Subtract from (1) the number of valence electrons the atom “owns” in the molecule (all of its
non-bonding electrons and half of its shared bonding electrons).
(3) Total: zero, there is no formal charge on that atom; nonzero, there is a negative (excess
electrons) or positive (deficiency of electrons) formal charge on the atom.
The sum of the formal charges on the atoms in a molecule equals the net charge of the molecule
or ion.
H
There is a net ionic charge on the polyatomic ion.
O
H
There is no formal charge on hydrogen.
H
There is a formal +1 charge on oxygen.
N
C
O
Resonance is a term that indicates there is more than one way to reasonably assign electrons in a
molecule without moving or rearranging the nuclei or changing the connectivity.
The molecule's real electronic structure (the resonance hybrid) is a blend of the individual
resonance contributors. The resonance contributors do not necessarily contribute equally to the
overall hybrid structure; rather the major contributor(s) will be those that are the most plausible
Lewis structures. (See next page for guidelines on “better” resonance structures.)
H
O
O
O
C
C
C
H
H
H
H
H
16
Some electronic arrangements are more stable than others.
 For neutral molecules, structures in which there are no formal charges on atoms are usually
"better" than those where formal charges are required.
 Where formal charges are necessary, these are generally of as small a magnitude as possible.
 Among alternative electronic structures, the most stable is that in which negative formal
charges are placed on the more electronegative atoms.
 A structure in which like formal charges (both positive or both negative) reside on adjacent
atoms is less stable.
17
VSEPR THEORY
VSEPR (Valence Shell Electron Pair Repulsion) Theory accounts for the geometries of many
molecules by considering the repulsive influences of the electrons around a given atom. There is
a difference between molecular geometry (determined by electrons) and molecular shape
(considering atoms only). We consider only geometry in this course.
Predicting Geometries and Bond Angles around a central atom
(1) Count each pair of bonding and non-bonding electrons on one atom. Count each multiple
bond as only one pair of bonding electrons.
(2) 2 pair: linear (180°); 3 pair: trigonal planar (120°); 4 pair: tetrahedral (109.5°); 5 pair:
bipyramidal; 6 pair: octahedral (the last two are not considered in this course).
(3) Non-bonding electron-pair repulsion > bonding electron-pair repulsion.
A
A
A
A
A
A
SUGGESTED PROBLEMS
A
A
1. Draw Lewis structures for and predict geometries of: BeH2, BH3, CH4, water, ammonia.
2. Look at a set of resonance structures for a molecule and verify that each contributor has the
same predicted geometry. Remember, to be resonance contributors, structures differ only in the
placement of electrons, not nuclei.
Review the formaldehyde example shown earlier by comparing the predicted geometries of the
two important resonance structures.
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SUGGESTED PROBLEMS
1. a) In the Lewis structure for acetic acid drawn during class, it was explicitly stated that we
would consider only one resonance structure. Draw another resonance structure in which all
atoms have an octet. Include formal charges.
b) Draw resonance contributors (include formal charges) for urea that resemble the one for acetic
acid above. Remember, none of these atoms can have more than 8 electrons.
2. There are three connectivities (structures) for the anion composed of carbon, nitrogen and
oxygen: CNO, NOC, NCO. Some resonance contributors for the NCO structure are shown
below. Confirm that each contributor accounts for all the valence electrons and that all of
them show the same number of valence electrons. Confirm the assignment of formal charges.
2
N
C
O
N
C
O
N
C
O
N
C
O
Explain why the one “best” resonance representation (this is not the hybrid) for this anion is
the one shown below, taking into account the guidelines for “better” resonance structures.
N
C
O
Try drawing resonance contributors for the other two connectivities of these atoms, CNO and
NOC. Explain why the most stable connectivity is NCO (shown above).
3. Peroxynitrous acid ( connectivity ONOOH) is useful for fast oxidations; however, the
compound rapidly isomerizes to nitric acid. Draw electron-dot structures for peroxynitrous
acid and nitric acid.
Explain why these two species are isomers.
Draw all resonance contributors for peroxynitrous acid, stating which contributors are
equivalent and which are not. Assign formal charges to all atoms in the drawings. Repeat for
nitric acid.
4. The molecule I below has an important resonance contributing structure II. Explain why this
might be an important contributor and explain why III is not a valid resonance contributor.
[Hint: count electrons.]
19
H
H
H
N C
N C
I
II
N C
III
5. Draw two resonance contributors with formal charges for 1-aza-3-bora-allene, H2NCBH2.
Which do you think is the major contributor? Why? [Hint: think octet.]
Some selected answers to above problems
1a)
H O
H C C O H
H
b) Notice the similar relationship of the C=O and O: in acetic acid and the C=O and N: in urea.
Also, you can draw two equivalent resonance structures for urea.
3.
O N O O H
O N O O H
O N O O H
None of these are equivalent resonance contributors. The middle structure
probably contributes the least because N does not have an octet.
Notice that although
O in the third structure has a formal positive charge, O has an octet. Having an octet is always
more important than
20
6. Draw resonance contributors for the polyatomic ions carbonate and nitrate. Include formal
charges. Are the contributors equivalent or non-equivalent? The answers are in your text. These
are excellent first problems to solve.
Another question in your text asks you to write “a” structure for sulfate, (SO4)2–. However, this
problem is similar to the ones above, that is, there are several resonance contributors, no one of
which is adequate to show the true structure. A notable difference between the central atoms N,
C, and S in these three problems is that S is in the 3rd row of the Periodic Table and so can have
>8 electrons in its valence shell. You can try to complete the complete set of resonance structures
below. There are 32 valence electrons in sulfate.
O
1-
1-
2+
O
S
O
O
O
1-
1-
O
O
1+
S
O
O
1-
1-
O
1+
S
O
S
O
1-
O
(The double bond can be
between S and any of the O's)
O
O
1-
1-
1-
1-
O
1-
(The two double bonds can be
between S and any of the O's.
Draw the other structures
with two double bonds in the
space below. There should be
5 more.)
Here is one of them:
1-
O
1-
O
O
1+
S
O
1-
O
O
O
O
O
S
O
O
1-
1-
1+
S
1-
O
1-
1-
You can repeat this process using three double bonds (total 4) and four double bonds (1, the
answer shown in your textbook).
The resonance structures in each column are equivalent to each other. The sulfate ion has a
tetrahedral shape, with 4 equivalent S–O bonds.
O
O
S
O
O
21
SYMBOLIC REPRESENTATIONS OF COVALENT BOND FORMATION IN H2
1. Lewis Electron Dot Structures
H
H
H
H
H
H
2. AO/MO Pictures
energy released
-435 kJ/mol
 molecular orbital
1 s atomic orbitals
3. Energy Level Diagrams
Number of Atomic Orbitals = Number of Molecular Orbitals
1s AO
[Conservation of Orbitals]
*
(antibonding MO)

(bonding MO)
1s AO
Unoccupied molecular orbital (antibonding, *)
Atomic orbitals
Occupied molecular orbital (bonding, )
22
DRAWING ATOMIC AND MOLECULAR ORBITAL PICTURES
Sigma () covalent bonds result from a "head to head" overlap of atomic orbitals
along the internuclear axis.
The table below shows how a sigma covalent bond can be formed from
combinations of s+s orbitals or s+p orbitals or p+p orbitals. The identity of the
orbital is irrelevant – the bond is formed head-to-head along the internuclear axis.
ATOMIC ORBITALS
from two participating atoms
BONDING MOLECULAR
ORBITALS
two atoms are bonded together
Smoothed
outline
drawing
showing
nuclei
s
s



s
p



p
p



Any single bond is necessarily a sigma bond.

Each of the two atomic orbitals that form the bond can be occupied by one electron or
two electrons can occupy one orbital and the other orbital is empty.
The resulting bonding molecular orbital is occupied by the total of two electrons.

Later, we discuss molecules of the type C=C or C≡C, containing a “double bond” or a “triple
bond”. We will see that in a set of multiple bonds (double or triple bond), only one can be a
sigma bond; the others are pi () covalent bonds.
23
Examples
Li2 (dilithium)
Each lithium is 1s22s1 and thus has one valence electron. The
unfilled valence s atomic orbitals on each lithium can combine to form a sigma molecular orbital.
Li · + Li·  Li—Li
The AO/MO picture on the previous page for s + s →  should be drawn for this bond-forming
reaction. The energy level diagram for the reaction is identical to the one previously for H2,
except the reacting AO's are labeled as 2s.
H2O (dihydrogen oxide)
Each hydrogen atom is 1s1; the oxygen atom is
1s2(2s22px2py1pz1) (confirming Hund's Rule). The unfilled s orbitals of the hydrogens, and the
unfilled p orbitals of the oxygen, can combine to form two sigma molecular orbitals. Only the
valence bonding electrons are shown below – add the non-bonding 2s2 and 2px2 valence
electrons as dots on oxygen.
H· + H· + ·O·  H—O—H
The AO/MO picture shown previously for s + p →  should be drawn for this bond-forming
reaction. The energy level diagram for the bond formation is a bit different from those shown on
a previous page because the H s atomic orbital and the O p atomic orbitals are not of equal
energy and because there is involvement of lower energy orbitals. So it is not included here. We
will look at the molecular orbitals in H2O from a different point of view in the section on
"Hybridization".
SUGGESTED PROBLEMS
1. For these bonding problems, each atom that participates in bond formation will have one
unpaired electron in the reacting orbital. For the atoms H and F, write the electron
configuration and specify the valence orbital that contains the single electron.
2. Draw a dot structure equation for formation of H−F and F−F from their constituent atoms
similar to the one shown for H2 on the previous page and the examples above.
3. Draw AO/MO pictures for the formation of molecular HF and F2 from their atoms using the
pictures on a previous page as a guide. Show only the relevant atomic and molecular orbitals
involved in the bond formation, ignoring the others.