Download Structure and Bonding

Chapter 1
Structure and Bonding
© 2006 Thomson Higher Education
Scientific Revolution
Scientific revolution leads to:
• Safer and more effective medicines
• Cures for genetic diseases
• Increased life span
• Improved quality of life
Scientific advances in medicine and biology require
an understanding of organic chemistry
Organic Chemistry
Organic Chemistry
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Historically
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The chemistry of compounds found in living organisms
Prostaglandins
• Large class if molecules found in tissues and fluids of the body
• Cyclooxygenase reaction of arachidonic acid produces
prostaglandin H2 (PGH2) as the initial step in the body’s
response to inflammation
• Active site of enzyme contains heme group and tyrosine
Carbon
Organic chemistry
• All organic compounds contain the element
carbon
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4A element
Shares four electrons
Forms four strong covalent bonds
Bonds to other carbons to create chains and
rings
Not all carbon compounds are derived from
living organisms
Over 99% of 26 million known compounds
contain carbon
1.1 Atomic Structure: The Nucleus
Structure of atom
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Nucleus
• Positively charged
• Made up of protons (positively charged) and neutrons (neutral)
• Small (10-14 to 10-15m)
• Contains essentially all the mass of the atom
Electron cloud
• Negatively charged electrons in cloud around nucleus
• Atomic diameter is about 2 Angstroms (Å)
-10 m = 100 pm
• The unit angstrom is 10
• Atomic diameter in SI units is 2  10-10 m (200 picometers (pm))
Atomic Structure: The Nucleus
Atomic number (Z)
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Number of protons in the atom's nucleus
Mass number (A)
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Number of protons plus neutrons in the atom’s nucleus
All the atoms of a given element have the same
atomic number
Isotopes
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Atoms of the same element (same Z) that have
different numbers of neutrons and therefore different
mass numbers (different A)
Atomic mass (atomic weight)
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Weighted average mass in atomic mass units (amu) of
an element’s naturally occurring isotopes
1.2 Atomic Structure: Orbitals
Quantum Mechanical Model
• Behavior of a specific electron in an atom described
by mathematical expression called a wave equation
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Wave equation is similar to mathematical expression
used to describe motion of waves in fluids
• Solution of wave equation is called a wave function
• Wave function is an orbital
• Orbital denoted by Greek letter psi, 
• A plot of 
2
describes volume of space around
nucleus the electron is most likely to occupy
• Electron cloud has no specific boundary
Atomic Structure: Orbitals
Four different kinds of orbitals for electrons
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Denoted s, p, d, and f
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s and p orbitals most important in organic chemistry
• s orbitals
• Spherical, nucleus at center
• p orbitals
• Dumbbell-shaped, nucleus at middle
• d orbitals
• Four cloverleaf-shaped and one dumbbell-doughnut
Atomic Structure: Orbitals
Orbitals are grouped in electron shells of increasing size and
energy
Electron Shell
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A group of an atom’s electrons with the same principal quantum
number
Each orbital can be occupied by two electrons
First shell contains one s orbital, denoted 1s, which holds only two
electrons
Second shell contains four orbitals, one s orbital (2s) and three p
orbitals (2p), which hold a total of eight electrons
Third shell contains nine orbitals, one s orbital (3s), three p orbitals
(3p), and five d orbitals (3d), which hold a total of 18 electrons
Atomic Structure: Orbitals
p Orbital
• In each shell, beginning with the second, there are three
perpendicular p orbitals, px, py, and pz, of equal energy
• Lobes of a p orbital are separated by a region of zero electron
density, a node
• Each lobe has a different algebraic sign, + and -, represented by
different colors
• Algebraic signs are not charges
1.3 Atomic Structure: Electron
Configuration
Ground-state electron configuration
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The most stable, lowest-energy electron
configuration of a molecule or atom
Three rules:
1. Aufbau principle
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Lowest-energy orbitals fill first: 1s  2s  2p  3s 
3p  4s  3d
2. Pauli Exclusion Principle
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Electron spin can have only two orientations, up 
and down 
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Only two electrons can occupy an orbital, and they
must be of opposite spin to have unique wave
equations
Atomic Structure: Electron
Configuration
Hund's rule
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If two or more empty orbitals of equal energy are
available, electrons occupy each orbital with parallel
spins until all orbitals have one electron
1.4 Development of Chemical Theory
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In 1858 August Kekulé and Archibald Couper independently
proposed that carbon is tetravalent (always forms four bonds)
Emil Erlenmeyer proposed a carbon-carbon triple bond for
acetylene
Alexander Crum Brown proposed a carbon-carbon double bond
In 1865 Kekulé suggested that carbon chains can double back
to form rings of atoms
In 1874 Jacobus van’t Hoff and Joseph Le Bel proposed four
atoms to which carbon is bonded sit at the corner of a regular
tetrahedron
Development of Chemical Theory
Atoms bond because the compound that results is more
stable and lower in energy than the separate atoms
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Energy is released from the chemical system when a
bond forms
Energy is consumed by the system when a bond
breaks
Valence shell
• Outer most electron shell of an atom
• Eight electrons in valence shell (an electron octet)
impart special stability to noble-gas elements in 8A
• Main group elements are governed by their tendency
to take on electron configuration of the nearest noble
gas
Development of Chemical Theory
Ionic compounds
• Some elements achieve an octet configuration by gaining or
losing electrons
• When an electron is gained or lost from a neutral atom an ion
is formed
• Ions are charged because they have different numbers of
protons and electrons
• Ions are held together by an electrostatic attraction, like in
Na+ Cl-, forming an ionic bond
Carbon achieves an octet configuration by sharing
electrons
Covalent Bond
• A bond formed by sharing electrons between atoms
Molecule
• A neutral collection of atoms held together by covalent bonds
Development of Chemical Theory
Lewis structures (electron-dot structures)
• Representations of covalent bonds in molecules
Kekulé structures (line-bond structures)
Development of Chemical Theory
Number of covalent bonds depends on how many additional valence
electrons needed to reach noble-gas configuration
• H (1s) needs one more electron to attain (1s2)
• N (2s22p3) needs three more electrons to attain (2s22p6)
Lone-pair electrons
• Valence-shell electron pairs not used for bonding
• Lone-pair electrons can act as nucleophiles and react with
electrophiles
Worked Example 1.1
Predicting the Number of Bonds to Atoms in a
Molecule
How many hydrogen atoms does phosphorus
bond to in phosphine, PH??
Worked Example 1.1
Predicting the Number of Bonds to Atoms in a
Molecule
Strategy
• Identify the periodic group of phosphorus,
and tell from that how many electrons (bonds)
are needed to make an octet.
Worked Example 1.1
Predicting the Number of Bonds to Atoms in a
Molecule
Solution
• Phosphorus, like nitrogen, is in group 5A of
the periodic table and has five valence
electrons. It thus needs to share three more
electrons to make an octet and therefore
bonds to three hydrogen atoms, giving PH3.
1.5 The Nature of Chemical Bonds:
Valence Bond Theory
Valence bond theory
• Bonding theory that describes a
covalent bond as resulting from
the overlap of two atomic
orbitals
• Electrons are paired in the
overlapping orbitals and are
attracted to nuclei of both
atoms, thus bonding the two
atoms together
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H–H bond results from the
overlap of two singly occupied
hydrogen 1s orbitals
H-H bond is cylindrically
symmetrical
Bonds formed by head-on
overlap of two atomic orbitals
along a line drawn between the
nuclei are sigma (s) bonds
The Nature of Chemical Bonds:
Valence Bond Theory
Bond strength
• H2 molecule has 436 kJ/mol less energy than the
starting 2 H atoms, the product is more stable than
the reactant and the H-H bond has a strength of 436
kJ/mol
• Conversely, the bond dissociation energy of H2 is 436
kJ/mol because it requires 436 kJ/mol of energy to
break the H2 bond
The Nature of Chemical Bonds:
Valence Bond Theory
There is an optimum distance between nuclei that leads to
maximum stability called the bond length
Bond length
• The distance between
nuclei at the minimum
energy point
• Because a covalent bond
is dynamic, like a spring,
the characteristic bond
length is the equilibrium
distance between the
nuclei of two atoms that
are bonded to each other
1.6 sp3 Hybrid Orbitals and the
Structure of Methane
Carbon has four valence electrons (2s22p2) that form four
bonds
Methane CH4
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All four carbon-hydrogen bonds in methane are identical and
are spatially oriented toward the corners of a regular
tetrahedron
sp3 hybrid orbitals
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A hybrid orbital derived from the combination of an s atomic
orbital with three p atomic orbitals
• Linus Pauling (1931) showed mathematically how s orbitals
and p orbitals on an atom can combine, hybridize, to form
four equivalent atomic orbitals with tetrahedral orientation
called sp3 hybrids
sp3 Hybrid Orbitals and the Structure
of Methane
Two lobes of a p orbital have different algebraic signs (+ and -) in
corresponding wave functions
• When a p orbital hybridizes with an s orbital
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Positive p lobe adds to s orbital generating larger lobe of sp3 hybrid
Negative p lobe subtracts from s orbital generating smaller lobe of sp3
hybrid
Resulting asymmetrical sp3 hybrid is strongly oriented in one direction
Larger lobe of sp3 hybrid can overlap more effectively with an orbital
from another atom forming much stronger bonds than s or p orbitals
sp3 Hybrid Orbitals and the Structure
of Methane
When each of the four identical sp3 hybrid orbitals of
carbon overlaps with the 1s orbital of a hydrogen atom,
four identical C-H bonds are formed and methane results
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Each C-H bond in methane has a strength of 436 kJ/mol and a
length of 109 pm
The four C-H bonds of methane have a specific geometry and
describe a characteristic bond angle
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The angle formed between two adjacent bonds
1.7 sp3 Hybrid Orbitals and the
Structure of Ethane
Orbital hybridization accounts for the bonding together of carbon
atoms into chains and rings
Ethane C2H6
• Tetrahedral
• Bond angles are near 109.5º
• Carbon-carbon single bond
• Formed by s overlap of
sp3 hybrids from each
carbon
• The remaining sp3
hybrids of each carbon
overlap with 1s orbitals of
three hydrogen atoms to
form six carbon-hydrogen
bonds
1.8 sp2 Hybrid Orbitals and the
Structure of Ethylene
Ethylene C2H4
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Carbon-carbon double bond
• Four shared electrons
Planar (flat)
Bond angles 120º
sp2 hybrid orbitals
• A hybrid orbital derived by
combination of an s atomic
orbital with 2p atomic orbitals
• One p orbital remains nonhybridized
sp2 Hybrid Orbitals and the
Structure of Ethylene
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s bond in ethylene formed by head-on overlap of two sp2
hybrid orbitals
• Two non-hybridized 2p orbitals overlap sideways forming a
p bond
• Carbon-carbon double bond is shorter and stronger than
carbon-carbon single bond
Worked Example 1.2
Predicting the Structures of Simple Molecules
from Their Formulas
Commonly used in biology as a tissue
preservative, formaldehyde, CH2O, contains a
carbon-oxygen double bond. Draw the linebond structure of formaldehyde, and indicate
the hybridization of the carbon atom.
Worked Example 1.2
Predicting the Structures of Simple Molecules
from Their Formulas
Strategy
• We know that hydrogen forms one covalent
bond, carbon forms four, and oxygen forms
two. Trial and error, combined with intuition,
is needed to fit the atoms together.
Worked Example 1.2
Predicting the Structures of Simple Molecules
from Their Formulas
Solution
• There is only one way that two hydrogens, one
carbon, and one oxygen can combine
• Like the carbon atoms in ethylene, the carbon atom
in formaldehyde is sp2-hybridized
1.9 sp Hybrid Orbitals and the
Structure of Acetylene
Acetylene
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Linear
Carbon-carbon triple bond
• Six shared electrons
• Bond angles are 180º
sp hybridized orbital
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A hybrid orbital derived
from the combination of
one s and one p atomic
orbital
• The two sp hybrids are
oriented at an angle of
180º to each other
• Two 2p orbitals remain
non-hybridized
sp Hybrid Orbitals and the Structure
of Acetylene
1.10 Hybridization of Nitrogen,
Oxygen, Phosphorus, and Sulfur
Covalent bonds formed by other elements can also be
described using hybrid orbitals
Nitrogen
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Methylamine CH3NH2
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Organic derivative of ammonia and the substance responsible for
the odor of rotting fish
Bond angles are close to the 109.5º tetrahedral angle found in
methane
Nitrogen hybridizes to form four sp3 orbitals
3
• One of the four sp orbitals is occupied by two nonbonding
electrons
Hybridization of Nitrogen, Oxygen,
Phosphorus, and Sulfur
Oxygen
• Methanol CH3OH
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Methyl alcohol
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Bonds are close to the109.5º tetrahedral angle
Two of the four sp3 hybrid orbitals on oxygen are
occupied by nonbonding electron lone pairs
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Hybridization of Nitrogen, Oxygen,
Phosphorus, and Sulfur
Phosphorus
• Most commonly encountered in biological molecules
in organophosphates
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compounds that contain a phosphorus atom bonded to
four oxygens with one of the oxygens also bonded to
carbon
• Methyl phosphate CH3OPO32• sp3 hybrid orbitals on phosphorus
Hybridization of Nitrogen, Oxygen,
Phosphorus, and Sulfur
Sulfur
• Commonly encountered in biological molecules
• Thiols
• Have a sulfur atom bonded to one hydrogen and
one carbon
• Sulfides
• Have a sulfur atom bonded to
two carbons
• Methanethiol CH3SH
• Produce by various bacteria
• Simplest example of a thiol
• sp3 hybridization
• Dimethyl Sulfide (CH3)2S
• Simplest example of a sulfide
• sp3 hybridization
1.11 The Nature of Chemical Bonds:
Molecular Orbital Theory
Molecular Orbital Theory (MO)
• A description of covalent bond formation as resulting from a
mathematical combination of atomic orbitals (wave functions)
to form molecular orbitals
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Additive combination of two 1s orbitals
• Leads to formation of a low energy s bonding MO
• Egg shaped
Subtractive combination of two 1s orbitals
• Leads to formation of a high energy s* antibonding MO
• Elongated dumbbell shaped with a node
The Nature of Chemical Bonds:
Molecular Orbital Theory
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Additive combination of two 2p orbitals
• Leads to formation of a low energy p bonding MO
• The p bonding MO is formed by combining p orbital lobes with the
same algebraic sign
• No node between nuclei
Subtractive combination of two 2p orbitals
• Leads to formation of a high energy p* antibonding MO
• The p* antibonding MO is formed by combining p orbital lobes
with different algebraic signs
• Node between nuclei
Only the bonding p MO is occupied
1.12 Drawing Chemical Structures
Condensed Structures
• A shorthand way of writing structures in which
carbon-hydrogen and carbon-carbon bonds are
understood rather than explicitly shown
• 2-Methylbutane
Drawing Chemical Structures
Skeletal Structure
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A shorthand way of writing structures in which carbon
atoms are assumed to be at each intersection of two
lines (bonds) and at the end of each line
Three rules:
1. Carbon atoms not usually
shown, they are assumed
2. Hydrogen atoms bonded
to carbon are assumed
3. Atoms other than carbon
and hydrogen are shown
Note: Sometimes the writing order of
atoms is inverted to make bonding
connections clearer
Drawing Chemical Structures
Worked Example 1.3
Interpreting Line-Bond Structures
Carvone, a substance responsible for the odor of
spearmint, had the following structure. Tell how many
hydrogens are bonded to each carbon, and give the
molecular formula of carvone.
Worked Example 1.3
Interpreting Line-Bond Structures
Strategy
• The end of a line represents a carbon atom
with 3 hydrogens, CH3
• A two-way intersection is a carbon atom with
2 hydrogens, CH2
• A three-way intersection is a carbon atom
with 1 hydrogen, CH
• A four-way intersection is a carbon atom with
no attached hydrogens
Worked Example 1.3
Interpreting Line-Bond Structures
Solution