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Organic Chemistry
7th Edition
Paula Yurkanis Bruice
Chapter 1
Remembering General
Chemistry:
Electronic Structure
and Bonding
Paula Yurkanis Bruice
University of California,
Santa Barbara
© 2014 Pearson Education, Inc.
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Organic and Inorganic
Early Humans can tell the difference between the kinds of materials in
their world.
“You can live on roots and berries,” they might have said, “but you can’t
eat dirt.”
“You can stay warm by burning tree branches, but
you can’t burn rocks.”
In 1807 Jö ns Jakob Berzelius gave names to
the two kinds of materials.
Compounds from living organisms contain an
unmeasurable vital force—the essence of life
and they are “organic.”
Compounds derived from minerals—those
lacking that vital force were called “inorganic.”
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Organic Chemistry
• Carbon-containing compounds were once considered “organ
compounds” available only from living organisms such as proteins,
enzymes, vitamins, lipids, carbohydrates, and nucleic acids.
• The synthesis of the simple organic compound urea in 1828 (by
Friedrich Wö hler) showed that organic compounds can be prepared in the
laboratory from non-living material (ammonium cyanate).
• Today, organic natural products are routinely
• synthesized in the laboratory.
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Why Carbon?
• Carbon neither gives up nor accepts electrons because it
is in the center of the second periodic row.
• Consequently, carbon forms bonds with other carbons
and other atoms by sharing electrons.
• The capacity of carbon to form bonds in this fashion
makes it the building block of all living organisms.
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Why Study Organic Chemistry?
• Since carbon is the building block of all living
organisms, a knowledge of Organic Chemistry is a
prerequisite to understanding Biochemistry, Medicinal
Chemistry, Chemical Engineering, Polymer, and
Pharmacology.
• Indeed, Organic Chemistry is a required course for
studying Chemical Engineering, Pharmacy, Medicine,
and Dentistry.
• Admission into these professional programs is highly
dependent on your performance in Organic Chemistry.
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Examples of Organic
Compounds Used as Drugs
Methotrexate, Anticancer Drug
5-Fluorouracil,
Colon Cancer Drug
Tamiflu, Influenza
Drug
AZT, HIV Drug
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Examples of Organic
Compounds Used as Drugs
Haldol, Antipsychotic
Elavil, Antidepressant
Prozac, Antidepressant
Viagra, Treats
Erectile Dysfunction
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Natural Versus Synthetic
Are natural substances (those made in Nature) are superior to
synthetic ones?
Synthesized penicillin or estradiol,(여성 성호르몬인 에스트로젠의 주성분)
it is exactly the same in all respects as the compound in nature.
Synthesized analogs of morphine
(compounds with structures similar to but
not identical to that of morphine) that have
pain-killing effects like morphine but,
unlike morphine, are not habit Forming.
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1.1 The Structure of an Atom
• An atom consists of electrons, positively
charged protons, and neutral neutrons.
• Electrons form chemical bonds.
• Atomic number: numbers of protons in its nucleus
• Mass number: the sum of the protons and neutrons
of an atom
• Isotopes have the same atomic number but
different mass numbers.
• The atomic weight: the average weighted mass of its atoms
• Molecular weight: the sum of the atomic weights of all the atoms
in the molecule
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Carbon
98.89% of naturally occurring carbon atoms contain six neutrons:
a mass number of 12
1.11% have seven neutrons: a mass number of 13
The atomic mass of 12C is 12.0000 amu; the atomic mass of 13C is 13.0034
amu. Therefore, the atomic weight of carbon is 12.011 amu.
(0.9889 x 12.0000) + (0.0111 x 13.0034) = 12.011
An atomic mass unit (amu) is defined as exactly 1/12 of the mass of 12C.
A trace amount of 14C (six protons and eight neutrons) and this isotope of carbon is
radioactive (decaying with a half-life of 5730 years).
As long as a plant or animal is alive, it takes in as much 14C as it excretes or exhales.
When it dies, it no longer takes in 14C, so the 14C in the organism slowly decreases.
Then, the age of a substance derived from a living organism can be determined by its
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content.
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14C
1.2 The Distribution of Electrons in an Atom
• Quantum mechanics uses the mathematical equation of wave
motions to characterize the motion of an electron around a
nucleus.
• Wave functions or orbitals tell us the energy of the electron and
the volume of space around the nucleus where an electron is
most likely to be found.
• The atomic orbital closer to the nucleus has the lowest energy.
• Degenerate orbitals have the same energy.
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The ground-state electronic configuration describes the orbitals
occupied by the atom’s electrons with the lowest energy
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The following principles determine which orbitals
electrons occupy:
• The Aufbau principle: an electron always goes to the available orbital
with the lowest energy
• The Pauli exclusion principle: only two electrons can occupy one atomic
orbital and the two electrons have opposite spin
• Hund’s rule: electrons will occupy empty degenerated orbitals before
pairing up in the same orbital
Electrons in inner shells (those below the outermost shell) are called core
electrons and they do not participate in chemical bonding.
Electrons in the outermost shell are called valence electrons and they determine
an element’s chemical properties.
So, the chemical behavior of an element depends on its electronic configuration.
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1.3 IONIC AND COVALENT BONDS
Lewis’s theory: an atom will give up, accept, or share electrons in
order to achieve a filled outer shell or an outer shell that contains
eight electrons → octet rule
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Ionic Bonds : Formed by the Transfer of Electrons
Attractive forces between opposite charges are called electrostatic
attractions
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Covalent Bonds : Formed by Sharing Electrons
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• Equal sharing of electrons: nonpolar covalent bond (e.g., H2)
• Sharing of electrons between atoms of different electronegativities:
polar covalent bond (e.g., HF)
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Polar covalent bonds
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A Polar Bond Has a Dipole Moment
• A polar bond has a negative end and a positive end
dipole moment (D) = m = e x d
(e) : magnitude of the charge on the atom
(d) : distance between the two charges
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Electrostatic Potential Maps
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1.4 HOW THE STRUCTURE OF A COMPOUND
IS REPRESENTED: Lewis Structure
Formal charge = number of valence electrons –
(number of lone pair electrons +1/2 number of bonding electrons)
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Nitrogen has five valence electrons
Carbon has four valence electrons
Hydrogen has one valence electron and halogen has seven
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Important Bond Numbers
Neutral
Cationic
Anionic
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Non-Octet Species
• In the 3rd and 4th rows, expansion beyond the octet to 10
and 12 electrons is possible.
Sulfuric Acid
Periodic Acid
Phosphoric Acid
• Reactive species without an octet such as radicals,
carbocations, carbenes, and electropositive atoms (boron,
beryllium).
Nitric Oxide
Radical
Radical,
Mammalian
Signaling Agent
Carbocation
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Carbene
Borane
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Kekulé Structures
Condensed Structures
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1.5 ATOMIC ORBITALS
An orbital tells us the volume of space around the nucleus where an
electron is most likely to be found
The s Orbitals
An electron in a 1s orbital can be anywhere within the 1s sphere.
2s orbital has a region where the probability of finding an electron falls to zero.
This is called a node, or, more precisely a radial node since this absence of
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electron density lies at one set distance from the nucleus.
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A node is a consequence of the wavelike properties of an electron.
There are two types of waves: traveling waves and standing waves.
Traveling waves move through space; light is an example of a traveling
wave.
A standing wave is confined to a limited space. A vibrating string of a
guitar is an example of a standing wave—the string moves up and down.
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The p Orbitals
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1.6 Molecular Orbitals
• Molecular orbitals belong to the whole molecule.
• s bond: formed by overlapping of two s orbitals.
• Bond strength/bond dissociation: energy required to break a bond or
energy released to form a bond.
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The change in energy that occurs as two 1s atomic orbitals
approach each other.
The internuclear distance at minimum energy is the length of the
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H-H covalent bond.
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In-phase overlap : bonding MO
out-of-phase overlap : antibonding MO
The strongest covalent bonds are formed by electrons that
occupy the molecular orbitals with the lowest energy.
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Therefore that H2+ would not be as stable as H2 because has only one
electron in the bonding orbital.
He2 does not exist: because each He atom would bring two
electrons, He2 would have four electrons: two filling the lower
energy bonding molecular orbital and the remaining two filling the
higher energy antibonding molecular orbital.
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Sigma bond (s) is formed by end-on overlap of two
p orbitals:
A s bond is stronger than a p bond
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Pi bond (p) is formed by sideways overlap of two parallel
p orbitals:
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1.7 HOW SINGLE BONDS ARE FORMED IN
ORGANIC COMPOUNDS
Bonding in Methane
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Hybridization of One s and Three p Orbitals
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The orbitals used in bond formation determine the bond angles
• Tetrahedral bond angle: 109.5°
• Electron pairs spread themselves into space as far from
each other as possible
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The Bonds in Ethane
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Hybrid Orbitals of Ethane
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1.8 Bonding in Ethene: A Double Bond
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Diamond, Graphite, Graphene
Diamond is hardest of all substances: carbon bonded to others vis
sp3 orbitals.
Graphite is a slippery and soft solid: carbon atoms are sp2 hydridized.
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1.9 Bonding in Ethyne: A Triple Bond
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1.10 THE BONDS IN THE METHYL CATION,
THE METHYL RADICAL, AND THE METHYL ANION
Bonding in the Methyl Cation
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Bonding in the Methyl Radical
Bonding in the Methyl Anion
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1.11 THE BONDS IN AMMONIA AND
IN THE AMMONIUM ION
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1.12 THE BONDS IN WATER
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1.13 THE BOND IN A HYDROGEN HALIDE
Fluorine, chlorine, bromine, and iodine are known as the halogens, so HF,
HCl, HBr, and HI are called hydrogen halides.
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1.14 HYBRIDIZATION AND MOLECULAR GEOMETRY
orbitals used in bond formation determine the bond angle
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1.15 SUMMARY
If it forms no p bonds, it is sp3 hybridized;
if it forms one p bond, it is sp2 hybridized;
if it forms two bonds, it is sp hybridized.
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• The shorter the bond, the stronger it is
• The greater the electron density in the region of orbital overlap, the
stronger is the bond
• The more s character, the shorter and stronger is the bond
• The more s character, the larger is the bond angle
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• The shorter the bond, the stronger it is
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• The shorter the bond, the stronger it is
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1.16 THE DIPOLE MOMENTS OF MOLECULES
The vector sum of the magnitude and the direction of the individual
bond dipole determines the overall dipole moment of a molecule
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