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 CHM 211 (Organic Chemistry)Summer 2009 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:

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)

Carbon’s 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)

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: 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

1s 2s 2px 2py 2pz

H

C

N

O

F

at. #

1

6

7

8

9

3s 3px 3py 3pz

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 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 H–H bond results from the

overlap of two singly occupied hydrogen 1s orbitals.

The H-H bond is cylindrically symmetrical, sigma () 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 overlap of an sp3 orbital from 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.

C C

H

H

H

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

90 120

C CH

H

H

H

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Carbon-Carbon Bonds in Ethene

Two sp2-hybridized orbitals overlap to form a bond Two p orbitals overlap side-to-side to form a pi () bond sp2–sp2 bond and 2p–2p bond results in sharing four

electrons and formation of C=C double bond Electrons in the 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 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 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

C CH H

C C

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Orbitals of Acetylene

Two sp hybrid orbitals from each C overlap to form an sp–sp 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 . Two sp orbitals form 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.

C C

C C

C C

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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 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 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, 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 () 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 bond framework of the molecule

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Summary, cont’d.

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 bonds are weaker than C-C bonds.