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 CHM 211 (Organic Chemistry)Summer 2009

• Dr. Ned H. Martin
• Office Dobo 242E
• Telephone
• 962-3453 (campus)
• Email martinn_at_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/Mart
  inn/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.
• 93A, 90A-, 87B, 84B, 80B-, 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
• 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
• 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)

• Carbons atomic number 6, therefore it has 6
  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)
• Carbons 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
• 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.
• Hunds 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
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1.4 The Nature of the Chemical Bond

• Atoms form bonds because the compound that
  results is 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.
• The HH 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 HH 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 CH 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 (sppp
  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 CH bond has a strength of 438 kJ/mol and
  length of 110 pm
• Bond angle each HCH is 109.5, the tetrahedral
  angle.

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1.7 Hybridization sp3 Orbitals and the Structure
of Ethane

• Two Cs bond to each other by s overlap of an sp3
  orbital from each C.
• The other three sp3 orbitals on each C overlap
  with H 1s orbitals to form six CH bonds.
• The CH bond strength in ethane is 420 kJ/mol.
• The CC bond is 154 pm long and its strength is
  376 kJ/mol.
• All bond angles of ethane are tetrahedral.

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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 (spp
  sp2)
• sp2 orbitals are in a plane with 120 angles
• Remaining p orbital is perpendicular to the plane

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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
• sp2sp2 s bond and 2p2p ? bond results in
  sharing four electrons and formation of CC
  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
• HCH and HCC bond angles are about 120
• CC double bond in ethene is shorter and stronger
  than the C-C single bond in ethane
• The ethene CC bond length is 133 pm (Recall that
  the CC bond length in ethane is 154 pm)
• The CC 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

• The 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 spsp s bond.
• Two pz orbitals from each C form a pzpz ? 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 .
• Two sp orbitals form s bonds with hydrogens.
• The bond strength is 835 kJ/mol, much
  less than three times the strength of a C-C (3 x
  376 1128). The bond length is 120 pm.

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1.10 Hybridization of Other Elements

• Elements other than C can have hybridized
  orbitals.
• The HNH bond angle in ammonia (NH3) is 107.3,
  close to the tetrahedral 109.5.
• Ns orbitals (sppp) 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 HOH 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
• Electrons occupy orbitals around the nucleus.
• Different orbitals have different energy levels
  and different shapes
• s orbitals are spherical, p orbitals are
  dumbbell-shaped
• 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, contd

• 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, contd.

• Carbon uses hybrid orbitals to form bonds in
  organic molecules.
• 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 ? bonds from the overlapped sp2
  orbitals. The remaining p orbitals overlap
  side-to-side to form a ? bond. C-C p bonds are
  weaker than C-C s bonds.