Organic Chemistry Structure and Function Fifth Edition W' H' Freeman

1
Organic ChemistryStructure and FunctionFifth
EditionW. H. Freeman CompanyNew York

• K. Peter C. Vollhardt
• Neil E. Schore

2
CHAPTER 1Structure and Bonding in Organic
Molecules
3
The Scope of Organic Chemistry An Overview
1-1

• Functional groups determine the reactivity of
  organic molecules
• Alkanes No functional groups, only carbon and
  hydrogen. (Chapter 2)
• Alkane Reactions Alkane bond strengths and
  reactions. (Chapter 3)
• Cyclic Alkanes New properties and changes in
  reactivity (Chapter 4)
• Stereoisomerism Same connectivity different
  relative positioning of substituents in space
  (Chapter 5)
• Haloalkanes Substitution Reactions and
  Elimination Reactions (Chapters 6 and 7)
• Alkynes C-C triple bonds (Chapter 13)
• Aldehydes and Ketones Carbonyl Compounds CO
  double bonds. (Chapters 16 and 17)

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The Scope of Organic Chemistry An Overview
1-1

• Amines Nitrogen containing functional group
  (Chapter 21)
• Tools For Identification Spectroscopy (Chapters
  10, 11, 14 and 20)
• Carbohydrates and Amino Acids Multiple
  Functional Groups (Chapters 24 and 26)

5
The Scope of Organic Chemistry An Overview
1-1

• Synthesis is the making of new molecules.
• Wöhlers Synthesis of Urea
• Synthesis Construct complex organic chemicals
  from simpler, more readily available ones
  (Chapter 8).
• Reactions are the vocabulary, and mechanisms are
  the grammar of organic chemistry.
• Reactants (Substrates) Starting compounds
• Products
• Reaction Mechanism Underlying details of a
  reaction
• Reaction Intermediate Chemical species formed
  and then destroyed on the pathway between
  reactants and products.

6
Coulomb Forces A Simplified View of Bonding
1-2
Bonds are made by simultaneous coulombic
attraction and electron exchange. When two atoms
approach, the electrons of one are attracted by
the protons of the other and vice-versa.
Energy is released as the two atoms approach each
other. When the atoms get too close together,
the energy begins to rise again due to repulsions
between the two nuclei and the two sets of
electrons.
7
This minimum energy is called the bond strength,
and the distance between the two nuclei at this
point is called the bond length.
8
Ionic and Covalent Bonds The Octet Rule
1-3

• Covalent Bonds are based on the sharing of
  electrons.
• If the electrons are not shared equally, a polar
  covalent (partially ionic) bond is formed,
  otherwise a pure covalent bond is formed.

• Ionic Bonds are based on the transfer of one or
  more electrons from one atom to another. The
  resulting cation and anion are electrostatically
  attracted to each other.

9
Ionic and Covalent Bonds The Octet Rule
1-3
The periodic table underlies the octet
rule. Electrons in atoms occupy levels or shells
of fixed capacity. The first has room for 2, the
second 8, and the third 16. Noble gases have 8
valence electrons (Helium 2) and are
particularily stable. Other elements lack octets
in their outer electron shells and tend to form
molecules in such a way as to create a stable
octet arrangement.
10
In pure ionic bonds, electron octets are formed
by transfer of electrons. Alkali metals react
with halogens by the transfer of one electron
from the alkali metal to the halogen. Both atoms
achieve a noble gas configuration The alkali
metal of the preceding inert gas, the halogen of
the following inert gas.
IPNa 119 kcal mol-1 EACl -83 kcal
mol-1 -LE -120 kcal mol-1 ?E -84 kcal
mol-1
11
Valence electrons are conveniently indicated by
placing dots around the symbol for an element.
The letters represent the nucleus and the core
electrons, and the dots represent the valence
electrons
Hydrogen can either lose an electron to form an
H ion, or gain an electron to form a H-, or
hydride, ion
12

• In covalent bonds, electrons are shared to
  achieve octet configurations.
• Ionic bonds between identical atoms of the same
  element do not form.
• The high ionization potential of hydrogen
  prevents it from forming ionic bonds with
  halogens and other non-metallic elements.
• Ionic bonds are also unfeasible for carbon since
  it would require the loss of 4 electrons to
  achieve the octet of the preceding inert gas, or
  the gain of 4 electrons to achieve the octet of
  the following inert gas.

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• In these and similar cases, covalent bonding
  occurs. Atoms share electrons to achieve a noble
  gas configuration.

• In certain cases, one atoms supplies both of the
  electrons in the bond

• Often 4 electron (double) and 6 electron (triple)
  bonds are formed

14
In most organic bonds, the electrons are not
shared equally polar covalent bonds.

• Pure covalent bonds (perfect sharing of
  electrons) and ionic bonds (complete transfer of
  electrons) are two extreme types of bonding.
• Most bonds lie somewhere between these extremes
  and are called polar covalent bonds.
• Each element can be assigned an electronegativity
  value which represents its electron-accepting
  ability when participating in a chemical bond.
• The larger the difference in electronegativety
  between two atoms participating in a chemical
  bond, the more ionic is the bond.
• Bonds between atoms of different
  electronegativity are said to be polar bonds. A
  partial negative charge is found on the atom of
  higher electronegativity and an equal but
  positive charge on the other atom.

15

• As a rule of thumb, electronegativity differences
  less than 0.3 represent pure covalent bonds, from
  0.3 to 2.0 polar covalent bonds, and greater than
  2.0 ionic bonds.
• The separation of opposite charges in polar
  covalent molecules results in the formation of
  dipoles

16

• In symmetrical molecules such as CO2 and CCl4,
  the individual dipoles will cancel and the
  molecule is left with a zero dipole moment.

17
Electron repulsion controls the shapes of
molecules.

• The shapes of molecules can be predicted using
  the VSEPR method.
• Bonding and non-bonding electron pairs on the
  same atom will arrange themselves in three
  dimensions to be as far apart as possible.
• In the case of 2 electron pairs, as in BeCl2, a
  linear arrangement results. For 3 electrons
  pairs, as is in BCl3, a trigonal arrangement
  results, and in the case of 4 electron pairs, a
  tetrahedral arrangement occurs

18
Electron-Dot Model of Bonding Lewis Structures
1-4
Lewis structures are drawn by following simple
rules.

• Draw the molecular skeleton.
• Count the number of available valence electrons.
• Add one electron for each negative charge, if an
  anion.
• Subtract one electron for each positive charge,
  if a cation.
• Depict all covalent bonds by two shared
  electrons, giving as many atoms as possible a
  surrounding electron octet, except for H, which
  requires a duet.
• Elements at the right of the periodic table
  (non-metals) may contain lone pairs of electrons.
• Correct Lewis Structure Incorrect
  Lewis Structures

19
It is often necessary to use double or triple
bonds to satisfy the octet rule
20

• Assign charges to atoms in the molecule.
• Charge ( valence electrons in free, neutral
  atom)
• ( unshared electrons on the atom)
• ½( bonding electrons surrounding the atom)

In molecules such as nitric acid, charges occur
on individual atoms, even though the molecule
itself is neutral.
21
The octet rule does not always hold.

• The molecule or ion has an odd number of
  electrons.
• NO, CH3, NO2
• The central atom has a deficiency of electrons.
• CH3, BeCl2, BH3

• Past row 2 of the periodic table, the central
  atom may be surrounded by more than 8 electrons
  (expanded octet).

22
Covalent bonds can be depicted by straight lines.
Bonding pairs of electrons are most often
represented as straight lines single bonds as a
single line, double bonds as two parallel lines,
and triple bonds as three parallel lines. Lone
pairs of electrons are either shown as dots or
are omitted. Structures of this type are called
Kekulé structures.
23
Resonance Forms
1-5
The carbonate ion has several correct Lewis
structures.
Three equivalent structures must be drawn to
accurately represent the carbonate ion. The only
difference between these structures is the
placement of electrons.
24
But what is its true structure? The true
structure can be thought of as the average of all
three structures which is called a resonance
hybrid.
The 2 negative charges are delocalized over all
three oxygen atoms.
25
Other examples of resonance
26
Not all resonance forms are equal.

• Structures with a maximum of octets are most
  important.

• Charges should be preferentially located on atoms
  with compatible electronegativity. If this
  conflicts with rule 1, then rule 1 takes
  precedence.
• Structures with less separation of opposite
  charges are more important resonance contributors
  than those with more charge separation.

27
In some cases charge separation is necessary and
guideline 1 takes precedence over guideline 2
If there are two or more charge separated
resonance structures which comply with the octet
rule, the most favorable one places the charges
on atoms of compatible electronegativity
28
Atomic Orbitals A Quantum Mechanical Description
of Electrons around the Nucleus
1-6
The electron is described by wave equations. An
electron within an atom can have only certain
definite energies called energy states. Moving
particles such as electrons exhibit a wavelength
determined by the de Broglie relation
Where h is Planks constant, m is the mass of the
electron in kg, and v is the velocity of the
electron in m/s.
29
The electron waves contain nodes, where the
amplitude of the wave changes sign, and can
interact with each other, producing either
constructive or destructive inteference
30
The wave theory of electron motion is called
quantum mechanics. The quantum mechanical
equations describing the motion of the electrons
are called wave equations. The solutions of these
equations are called wave functions and are
represented by the Greek letter, ?. The square of
the wave function, evaluated at a point in space,
(x,y,z) represents the probability of finding the
electron at that point at any given time. Each
wave function corresponds to a specific discrete
energy and the system is said to be quantized.
31
Atomic orbitals have characteristic shapes.
In an artists rendition of an atomic orbital, a
surface is drawn which contains most of the
probability of finding the electron at a given
time. Nodes in a function become points or
planes of 0 probability of finding the
electron. Higher energy wave functions have more
nodes than do lower- energy wave functions. The
simplest atomic orbitals are spherical in shape
and are called s orbitals. The lowest energy s
orbital is the 1s orbital.
32
The next highest energy orbital is the 2s
orbital. Note that it contains a spherical nodal
surface.
Of slightly higher energy than the 2s orbital are
3 degenerate 2p orbitals. These orbitals are
shaped like a figure 8 and point along the 3
cartesian axes.
33
Following the 1s, 2s, and 2p orbitals are the 3s,
3p, 4s, 3d, etc. orbitals. Organic chemistry
deals primarily with the lower s and p orbitals.
34
The Aufbau principle assigns electrons to
orbitals.

• Lower energy orbitals are filled before those
  with higher energy.
• No orbital may be occupied by more than two
  electrons (Pauli Exclusion Principle). If two
  electrons occupy a single orbital, they must have
  opposite spins. Electrons of opposite spins in
  the same orbital are called paired electrons.
• Degenerate orbitals must each receive a single
  electron of the same spin before pairing of
  electrons occurs (Hunds rule).

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Atoms having a completely filled set of atomic
orbitals are said to have a closed shell
configuration. Atoms with a completely filled set
are said to have an open shell configuration.
The process of filling up the energy level
diagram one electron at a time is called the
Aufbau process. The d orbitals on atoms of row 3
and higher are involved in the formation of
expanded octets (10 and 12 electrons about a
central atom).
37
Molecular Orbitals and Covalent Bonding
1-7
The bond in the hydrogen molecule is formed by
the overlap of 1s atomic orbitals.
Atomic orbitals on different atoms may overlap.
The overlap of electron waves represented by the
atomic orbitals may result in constructive (in
phase) or destructive (out of phase)
inteference. In-phase overlap between two 1s
orbitals results in a new orbital having lower
energy than either of the s orbitals. This new
orbital concentrates the electron probability
between the two nuclei. Out-of-phase overlap
between two 1s orbitals results in a new orbital
having higher energy than either of the s
orbitals. This new orbital places most of the
electron probability to the left and right of the
two nuclei.
38
An energy level diagram can now be made of the
two overlapping orbitals, and the Aufbau process
used to determine the electronic configurations
of H2 and He2
39
The overlap of atomic orbitals gives rise to
sigma and pi bonds.
When n atomic orbitals overlap, n new molecular
orbitals are formed. When n is 2, one bonding
orbital and one antibonding molecular orbital are
formed. The energy lowering of the bonding
orbital and energy raising of the antibonding
molecular orbital with respect to the atomic
orbitals is called the energy splitting. The
energy splitting indicates the strength of the
bond formed. Atomic orbitals of the same size and
energy overlap to form the strongest
bonds. Geometrical factors also affect the degree
of overlap. Orbitals exhibiting directionality in
space (p orbitals) can overlap to form sigma (?)
or pi (?)bonds. All carbon-carbon single bonds
contain one sigma bond. Double and triple bonds
contain extra pi interactions.
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Hybrid Orbitals Bonding in Complex Molecules
1-8
Mixing of atomic orbitals from the same atom
results in new atomic orbitals of different
energy and directionality. Sp hybrids produce
linear structures. An incorrect structure for
BeH2 is predicted if 2s and 2p orbitals of Be are
overlapped with the 1s orbitals of H
42
Mixing the 2s orbital with one of the 2p orbitals
of Be results in two new hybrid sp orbitals, each
made up of 50 s and 50 p character. The
resulting bond angle is 180o which corresponds
with the observed bond angle in the BeH2 molecule.
Hybridization does not change the number of
orbitals on the atom. In this case two atomic
orbitals are replaced by two new hybrid orbitals.
The two un-hybridized p orbitals are still
available to hold electrons.
43
Sp2 hybrids create trigonal structures.
Hybridization of a 2s and two 2p orbitals results
in three new hybrid orbitals that point to the
corners of an equilateral triangle. The remaining
p orbital points up and down, perpendicular to
each of the three hybrid orbitals. Bond angles in
molecules using sp2 hybridization are
approximately 120o.
The molecule BH3 is isoelectric with the methyl
cation, CH3. Both involve sp2 hybridization
about the central atom.
44
Sp3 hybridizaton explains the shape of
tetrahedral carbon compounds.
When the 2s and all three 2p orbitals are
hybridized, four hybrid orbitals called sp3
orbitals are formed. These orbitals point to the
corners of a regular tetrahedron. Bond angles in
molecules using sp3 hybridization are
approximatelly 109.5o
45
Hybrid orbitals may contain lone electron pairs
ammonia and water.
Not all hybrid orbitals participate in bond
formation. Some may contain lone pairs of
electrons.
The bond angles in ammonia are 107.3o and that in
water are 104.5o, both close to 109.5o. The
slightly smaller bond angles in ammonia and water
are due to the slightly larger volume
requirements for lone pair electrons, which
forces the remaining bonding pair electrons
closer together.
46
Pi bonds are present in ethene (ethylene) and
ethyne (acetylene).
Molecules containing double or triple bonds
contain unhybridized p orbitals that overlap
lengthwise rather than end-on.
47
Structures and Formulas of Organic Molecules
1-9
To establish the identity of a molecule, we
determine its structure.
The empirical formula of a substance specifies
the kinds and ratios of elements present in the
substance. The empirical formula can be from an
elemental analysis of the substance. More than
one substance can have the same empirical
formula. Each of these substances will have its
own set of unique physical and chemical
properties, however. Substances having the same
empirical formula but different connectivity of
atoms are called constitutional or structural
isomers.
48
A chemist may be able to identify an unknown
substance if its properties match those of a
substance already determined. New substances
require other methods of identification such as
X-ray crystallography, or various forms of
spectroscopy. Two ways of representing the
structures of know molecules are ball-and-stick
models and space-filling models.
49
Several types of drawings are used to represent
molecular structures. Kekulé
Condensed Bond-Line
50
Tetrahedral carbon structures can be accurately
represented in three dimensions using the
dashed/wedged line notation.
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The Big Picture
1

• The importance of Coulombs Law
• Atomic attraction
• Relative electronegativity
• Electron repulsion model for shapes of molecules
• Choice of dominant resonance contributors
• The tendency of electrons to spread out
  (delocalize)
• Resonance forms
• Bonding overlap
• The correlation of the valence electron count
  with the Aufbau Principle.
• Associated stability of the elements in noble
  gas-octet-closed-shell configurations obtained by
  bond formation.
• The characteristic shapes of atomic and molecular
  orbitals
• Provides a feeling for the location of the
  reacting electrons around the nuclei.

52

• The overlap model for bonding
• Allows a judgment of energetics, directions and
  overall feasibility of reactions.

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Important Concepts
1

• Organic Chemistry Chemistry of carbon and its
  compounds.
• Coulombs Law Relates attractive or repulsive
  force between charges to the distance between
  them.
• Ionic Bonds Result from coulombic attraction of
  oppositely charged ions.
• Covalent Bonds Result from electron sharing
  between two atoms.
• Bond Length Average distance between two
  covalently bonded atoms
• Polar Bonds Formed between atoms of differing
  electronegativity

54
Important Concepts
1

• Shape of Molecules Strongly Influenced by
  electron repulsion.
• Lewis Structures Describe bonding using valence
  electron dots. Hydrogen receives a duet while
  other atoms receive an octet. Charge separation
  should be minimized but may be enforced by the
  Octet Rule.
• Resonance Forms When a structure is described
  by two or more Lewis structures differing only in
  their electron positions. The actual molecule is
  an average of the resonance forms. Some resonance
  structures may be more important that others.
• De Broglie Relation Relates wavelength of an
  electron to its mass and velocity.

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Important Concepts
1

• Wave Equations Describe motions of electrons
  about the nucleus. Solutions are called orbitals.
  These describe probabilities of finding the
  electrons in particular regions of space.
• s Orbital Spherical. P-orbital Figure Eight.
  Each orbital can hold two electrons of opposite
  spin. With increasing energy, the number of
  nodes in an orbital increases.
• Aufbau Principle Building electronic
  configurations by adding one electron at a time
  to the atomic orbitals, starting with those of
  lowest energy. (Pauli exclusion principle,
  Hunds Rule).

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Important Concepts
1

• Molecular Orbital Two overlapping atomic
  orbitals form either a bonding or an antibonding
  molecular orbital. The number of molecular
  orbitals equals the number of atomic orbitals
  overlapped.
• ? Bonds Formed when atomic orbitals overlap
  along the bond axis. ? bonds Formed from
  p-orbitals overlapping perpendicular to the bond
  axis.
• Hybrid Orbitals Formed by mixing of orbitals on
  the same atom. sp 2 linear orbitals, sp2 3
  trigonal orbitals , sp3 4 tetrahedral orbitals.
  Atomic orbitals not hybridized remain unchanged.
  Hybrid orbitals can contain either bonding or
  lone pair electrons.

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Important Concepts
1

• Elemental Analysis Determines ratios of types
  of atoms in a compound. Molecular Formula
  Actual number of atoms of each type.
• Constitutional Isomers (Structural Isomers)
  Same molecular formula but different connectivity
  of atoms. Different properties.
• Condensed and Bond-Line Formulas Abbreviated
  representations of molecules. Dashed-Wedged Line
  Drawings Illustrate molecules in three
  dimensions.