CHM 211 (Organic Chemistry)
Summer 2010
 Dr. Ned H. Martin
 Office: Dobo 242E
 Telephone:

962-3453 (campus)
 Email: martinn@uncw.edu
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Texts
 Organic Chemistry, 7th edition, McMurry
 Optional


Study Guide and Solutions Manual for McMurry's Organic
Chemistry, 7th edition
Molecular model kit
 Course Website (Syllabus, Grading Policy):
http://www.uncw.edu/chem/Courses/Martinn/chm211martin/index.htm
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Grading Policy
 Four 40-minute tests, each worth 60 points.
 The final exam will consist of six sections. The first four
are like the four tests; the higher grade counts. Section 5
is new material (since the last test). Section 6 is
comprehensive. You may take (or not) as many of the first
four sections as you want. Everyone must take sections 5
and 6.
 There will be no make up exams.
 Each of the tests may include at least one problem from
the homework assignments. Tests 2- 4 may contain one
review question from the previous test.
 93%=A, 90%=A-, 87%=B+, 84%=B, 80%=B-, etc.
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Attendance & Homework
 Attendance is expected, but not officially monitored for
grading purposes.

Missing 1 day in the summer is like missing 1 week during a
regular semester!
 Homework problems are assigned, but not collected.
 Actively working the homework problems allows you to
test whether you understand the material and serves as a
review guide for the exams.
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Keys to Success in CHM 211
 Memorization alone is not sufficient.
 Reasoning alone is not sufficient.
 Study three times:
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
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Before the lecture
After the lecture
Before the test
 Actively do problems (Keep a notebook).
 Cooperate – form study groups.
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What is Organic Chemistry?
 The study of carbon-containing compounds
 Important because:
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Carbon forms 4 bonds, and can bond to itself in long chains
Carbon has three different geometries giving rise to a variety
of structures
Carbon bonds strongly to other common elements: O, N, Cl,
etc.
Organic compounds have many applications and uses: dyes,
medicines, fabric, plastics, food (protein, carbohydrates, fats,
oils), fuel, pesticides, paint, preservatives, hormones, etc.
 This PowerPoint covers:
Chapter 1. Structure and Bonding
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C (Carbon)
 Carbon’s atomic number = 6, therefore it has 6
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protons in its nucleus.
A neutral atom of 12C has 6 protons, 6 neutrons and
6 electrons; its amu = 12 ( = 6p + 6n)
A neutral atom of 13C has 6 protons, 7 neutrons and
6 electrons; its amu = 13 ( = 6p + 7n)
A neutral atom of 14C has ? protons, ? neutrons and
? electrons; its amu = ? ( = ?p + ?n)
Carbon’s atomic weight = 12.011; this is a weighted
average of the three isotopes: 12C, 13C, and 14C.
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Parts of an Atom
 Protons (+ charge) and neutrons (0 charge) are in the
center or nucleus of the atom
 Electrons (- charge) are considered to be a cloud of
charge around the nucleus. Orbitals describe where
the electrons are. Electrons have very little mass
compared to protons and neutrons.
 Electrons are found in s orbitals (spherical), p orbitals
(dumbbell), or d orbitals (various shapes)
 Electrons are grouped in different layers or shells.
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1.1 Atomic Structure
 Structure of an atom
Positively charged nucleus (very dense, protons
and neutrons) and small (10-15 m)
 Negatively charged electrons are in a cloud (10-10
m) around nucleus
 Diameter is about 2  10-10 m (200 picometers (pm))
[the unit Angstrom (Å) is 10-10 m = 100 pm]

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1.2 Atomic Structure: Orbitals
 Quantum mechanics: describes electron energies
and locations by a wave function, 
 A plot of  2 describes the region where electrons are
most likely to be
 An electron cloud has no specific boundary so we
represent its shape by the region of highest
probability of finding an electron.
 Solutions of the wave equation give rise to regions of
electron density on each atom of specific shapes
(atomic orbitals)
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Shapes of Atomic Orbitals
 Four different kinds of orbitals occupied by electrons
 Denoted s, p, d, and f (listed in increasing energy)
 s and p orbitals are most important in organic chemistry
 s orbitals: spherical, with the nucleus at center
 p orbitals: dumbbell-shaped, with the nucleus at the
center
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p-Orbitals
 There are three
perpendicular p
orbitals, px, py, and
pz, of equal energy
 Lobes of a p orbital
are separated by
region of zero
electron density,
called a node.
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1.3 Atomic Structure: e- Configuration
 The lowest energy electron configuration of an atom
of any element can be predicted by following three
rules:
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The aufbau principle: Electrons are filled into the lowest
energy orbitals first (1s, then 2s, then 2p, then 3s, then 3p,
then 4s, then 3d)
The Pauli exclusion principle: Only two electrons may
occupy an orbital; they must have opposite spin orientations.
Hund’s rule: If there are two or more equal energy
(degenerate) orbitals available, the electrons will spread out
among the orbitals with parallel spins, only pairing up after
the orbitals are half-filled.
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Examples of Electron Configuration
at. #
H
C
N
1s
2s
2px 2py 2pz
3s
3px 3py 3pz
1
6
O
7
8
F
9
Ne 10
Cl 17
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1.4 The Nature of the Chemical Bond
 Atoms form bonds because the compound that results is
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more stable than the separate atoms.
Ionic bonds in salts form as a result of electron transfers,
followed by electrostatic attraction between opposite
charges.
Organic compounds form covalent bonds by sharing
electrons (G. N. Lewis, 1916).
Lewis structures show valence electrons of an atom as
dots.
 Hydrogen has one dot, representing its 1s electron.
 Carbon has four dots (2s2 2p2).
Stable molecule results in a completed shell, an octet
(eight e-) for main-group atoms (two for hydrogen).
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Number of Covalent Bonds to an Atom
 Atoms with one, two, or three valence electrons form
one, two, or three bonds.
 Atoms with four or more valence electrons form as
many bonds as they need electrons to fill the s and p
levels of their valence shells to reach a stable octet.
 Carbon has four valence electrons (2s2 2p2),
therefore forms four bonds (CH4).
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Valence of Oxygen and Nitrogen
 Oxygen has six valence electrons (2s2 2p4), so it
forms two bonds (H2O).
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Valence of Nitrogen
 Nitrogen has five valence electrons (2s2 2p3), and it
forms three bonds (NH3).
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Non-bonding electrons
 Valence electrons not used in bonding are called
nonbonding electrons, or lone-pair electrons.

Consider the nitrogen atom in ammonia (NH3):
 N shares six valence electrons in three covalent bonds;
the remaining two valence electrons are a nonbonding
(lone) pair.
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1.5 Valence Bond Theory
 Covalent bond forms when two
atoms approach each other closely
so that a singly occupied orbital on
one atom overlaps a singly occupied
orbital on the other atom.
 Electrons are paired in the
overlapping orbitals and are
attracted to nuclei of both atoms.
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
The H–H bond results from the
overlap of two singly occupied
hydrogen 1s orbitals.
The H-H bond is cylindrically
symmetrical, sigma (s) bond.
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Bond Energy
 The reaction 2 H·  H2 releases 436 kJ/mol.
 The product has 436 kJ/mol less energy than two H
atoms: H–H has bond strength of 436 kJ/mol.
(1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ).
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Bond Length
 Distance between
nuclei that leads to
maximum stability.
 If too close, they
repel because both
nuclei are positively
charged.
 If nuclei are too far
apart, bonding is
weak.
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1.6 Hybridization: sp3 Orbitals and the
Structure of Methane
 Carbon has 4 valence electrons (2s2 2p2)
 In CH4, all C–H bonds are identical (tetrahedral)
 How can this be explained ??
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1.6 Hybridization: sp3 Orbitals and the
Structure of Methane
 sp3 hybrid orbitals: s orbital and three p orbitals
combine to form four equivalent, unsymmetrical,
tetrahedral orbitals (s+p+p+p = sp3), Pauling (1931)
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Tetrahedral Structure of Methane
 sp3 orbitals on C overlap with 1s orbitals on 4 H
atoms to form four identical C-H bonds
 Each C–H bond has a strength of 438 kJ/mol and
length of 110 pm
 Bond angle: each H–C–H is 109.5°, the tetrahedral
angle.
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1.7 Hybridization: sp3 Orbitals and the
Structure of Ethane
 Two C’s bond to each other by s overlap of an sp3 orbital from
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each C.
The other three sp3 orbitals on each C overlap with H 1s
orbitals to form six C–H bonds.
The C–H bond strength in ethane is 420 kJ/mol.
The C–C bond is 154 pm long and its strength is 376 kJ/mol.
All bond angles of ethane are tetrahedral.
H H
H C C H
H H
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1.8 Hybridization: sp2 Orbitals and the
Structure of Ethene (Ethylene)
 sp2 hybrid orbitals: A 2s orbital of C combines with
two 2p orbitals, giving 3 orbitals (s+p+p = sp2)
 sp2 orbitals are in a plane with 120° angles
 Remaining p orbital is perpendicular to the plane
90
H
120
H
C C
H
H
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Carbon-Carbon Bonds in Ethene
 Two sp2-hybridized orbitals overlap to form a s bond
 Two p orbitals overlap side-to-side to form a pi ()
bond
 sp2–sp2 s bond and 2p–2p  bond results in sharing
four electrons and formation of C=C double bond
 Electrons in the s bond are centered between nuclei
 Electrons in the  bond occupy regions on either side
of a line between nuclei, above and below the plane
of the atoms.
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Carbon-Hydrogen Bonds in Ethene
 Each of 4 H atoms form s bonds with four sp2 orbitals
 H–C–H and H–C–C bond angles are about 120°
 C=C double bond in ethene is shorter and stronger
than the C-C single bond in ethane
 The ethene C=C bond length is 133 pm (Recall that
the C–C bond length in ethane is 154 pm)
 The C+C bond strength is 611 kJ/mol, less than twice
the strength of a C-C (2 x 376 = 752).
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1.9 Hybridization: sp Orbitals and the
Structure of Acetylene H C C H
 The C
C in acetylene (ethyne) is a triple bond, with
the carbons sharing six electrons
 A carbon 2s orbital hybridizes with a single p orbital
giving two sp hybrids
 The other two p orbitals on each C remain
unchanged
 sp orbitals are linear, oriented 180° apart (on x-axis)
 The two p orbitals are perpendicular, on the y-axis
and the z-axis
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Orbitals of Acetylene
 Two sp hybrid orbitals from each C overlap to form an
sp–sp s bond.
 Two pz orbitals from each C form a pz–pz  bond by
sideways overlap; py orbitals overlap similarly to form
a second  bond.
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Bonding in Acetylene
 Sharing of six electrons forms a C
C.
 Two sp orbitals form s bonds with hydrogens.
 The C C bond strength is 835 kJ/mol, much less
than three times the strength of a C-C (3 x 376 =
1128). The C C bond length is 120 pm.
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1.10 Hybridization of Other Elements
 Elements other than C can
have hybridized orbitals.
 The H–N–H bond angle in
ammonia (NH3) is 107.3°, close
to the tetrahedral 109.5°.
 N’s orbitals (s+p+p+p) hybridize
to form four sp3 orbitals.
 One sp3 orbital holds two
nonbonding electrons, and
three sp3 orbitals have one
electron each, forming s bonds
to three Hs.
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Hybridization of Oxygen in Water
 The oxygen atom is sp3-hybridized.
 Oxygen has six valence-shell electrons but forms
only two covalent bonds, leaving two lone pairs.
 The H–O–H bond angle is 104.5°, slightly smaller
than the perfect tetrahedral angle (109.5º) because of
electron-electron repulsion between the lone pairs.
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1.11 Molecular Orbital Theory
 A molecular orbital (MO): where electrons are most likely
to be found (specific energy and general shape) in a
molecule.
 The two (or more) atomic orbitals combine to make two
(or more) molecular orbitals.
 Additive combination (bonding) MO is lower in energy.
 Subtractive combination (antibonding) MO is higher.
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Molecular Orbitals in Ethene
 The  bonding MO results from combining p orbital
lobes with the same algebraic sign.
 The  antibonding MO comes from combining lobes
with opposite signs.
 Only the bonding MO is occupied by electrons.
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Summary
 Organic chemistry – chemistry of carbon compounds
 Atom: positively charged nucleus surrounded by
negatively charged electrons
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Electrons occupy orbitals around the nucleus.
Different orbitals have different energy levels and different
shapes

s orbitals are spherical, p orbitals are dumbbellshaped
 Covalent bonds - electron pair is shared between
atoms
 Valence bond theory - electron sharing occurs by
overlap of two atomic orbitals
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Summary, cont’d
 Hybrid Atomic Orbital Theory - electron sharing
occurs by overlap of two orbitals formed by combining
(hybridizing) two or more atomic orbitals (sp, sp2, sp3)
 Molecular orbital (MO) theory - bonds result from
combination of atomic orbitals to give molecular orbitals,
which belong to the entire molecule
 Sigma (s) bonds - Circular in cross-section and are
formed by head-on interaction
 Pi () bonds – “dumbbell” shape, from sideways
interaction of p orbitals; located above and below the
s bond framework of the molecule
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Summary, cont’d.
 Carbon uses hybrid orbitals to form bonds in organic
molecules.
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In single bonds with tetrahedral geometry, carbon has four sp3
hybrid orbitals
In double bonds with planar geometry, carbon uses three
equivalent sp2 hybrid orbitals and one unhybridized p orbital
Carbon uses two equivalent sp hybrid orbitals to form a triple
bond with linear geometry, with two unhybridized p orbitals
 Atoms such as nitrogen and oxygen also hybridize to
form strong, oriented bonds

The nitrogen atom in ammonia and the oxygen atom in water
are sp3-hybridized
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Quick Review
 Carbon
 One s and three p orbitals hybridize to form four sp3
orbitals
 In methane and ethane, C is tetrahedral, with ~109.5°
bond angles
 In ethene, One s and two p orbitals hybridize to form
three sp2 orbitals. The bonds between the nuclei are
the s bonds from the overlapped sp2 orbitals. The
remaining p orbitals overlap side-to-side to form a 
bond. C-C  bonds are weaker than C-C s bonds.
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