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Chapter 1
Introduction
• Organic Chemistry - is the study of carbon
compounds (ranging from simple
compounds such as CH4 to
complex ones such as proteins,
carbohydrates, DNA, etc…).
• Torbern Bergman (1770) – was the first to make
a distinction between organic compounds (“those
that can only be derived from living things”) and
inorganic compounds (“those that were found in
the nonliving world”).
Vitalism
The belief in a life force outside the
physical and chemical laws.
• Michel-Eugene Chevreul (1816) – converted an
organic substance (fat) to other organic substances
(fatty acids and glycerin) without the help of an
outside vital force.
Animal fat
NaOH
Soap + Glycerin
H2O
Soap
H3O+
“Fatty acids”
• Friedrich Wohler (1828) – accidentally
converted the inorganic salt ammonium
cyanate into the organic substance urea.
• Stanley Miller (1953) – used electrical discharges
to trigger reactions in a primitive “atmosphere” of
H2O, H2, NH3 (ammonia), and CH4 (methane). His
experiment demonstrated that the spontaneous
synthesis of organic compounds could have been
an early stage in the origin of life.
Mechanism
The belief that all natural
processes are governed by
physical and chemical laws.
• Vitalism - is the belief in a life force outside
the physical and chemical laws.
• Mechanism - is the belief that all natural
processes are governed by physical and
chemical laws.
Carbon Atoms: The Most Versatile
Building Blocks of Molecules
•
Carbon is tetravalent. It has 4 electrons in its
valence shell.
• Carbon completes its valence shell by sharing
electrons with other atoms in four covalent
bonds.
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
•
Carbon can form a covalent bond with another
carbon atom, forming the carbon skeleton of
organic compounds (e.g. C2H6, C2H4, …).
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
•
Carbon can bond to a variety of atoms, including
O, H, and N (e.g. CO2, urea CO(NH2)2…)
Carbon dioxide, CO2
O
C
O
O
H
Urea, CO(NH2)2
C
N
N
H
H
H
• Variation in carbon skeletons contributes to the
diversity of organic molecules.
• Carbon skeletons vary in:
 Length
 Shape and arrangement (branching or ring)
 Presence and location of double and/or triple bonds
Fig. 4.4. Variations in carbon skeletons
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
• Hydrocarbons - are organic molecules
consisting only of carbon (C)
and hydrogen (H) atoms.
Hydrocarbons
Aliphatic Hydrocarbons
Alkanes
Alkenes
Aromatic Hydrocarbons
Alkynes
I. Atomic Structure
A.
Overview
B.
Orbitals
C.
Electron Configuration
A. Overview
Atoms are composed of smaller parts, called
subatomic particles:
 Protons (p+)
 Neutrons (n0)
 Electrons (e-)
Protons (p+)
 are positively charged
 are tightly packed with the neutrons to form a
dense core, called the atomic nucleus (10-1410-15 m in diameter), at the center of the atom
 have a mass of 1 dalton (amu = 1.7 x 10-24 g)
Neutrons (n0)
 are uncharged
 are tightly packed with the protons to form the
atomic nucleus at the center of the atom
 have a mass of 1 dalton (amu = 1.7 x 10-24 g)
Electrons (e-)
 are negatively charged
 move around the nucleus at a distance of
about 10-10 m and at a speed of light forming a
cloud
 have a mass of 1/2000 dalton
Structure of an atom
 Positively charged nucleus (protons and neutrons) is
very dense and small (10-15 m)
 Negatively charged electrons form a cloud (10-10 m)
around nucleus
• Diameter is about 2  10-10 m (200 picometers
(pm)) [1 Å (angstrom) = 10-10 m = 100 pm]
Atomic Number and Mass Number
An atom is described by its:
• Atomic number (Z) = # of protons
• Mass number (A) = # of protons + # of neutrons
All the atoms of a given element have the same
atomic number.
Isotopes -
are atoms of the same element
with different numbers of neutrons
and thus different mass numbers.
Atomic Mass - is the weighted average mass in
atomic mass units (amu) of an
element’s naturally occurring
isotopes (= atomic weight).
B. Orbitals
Orbital
 is a three-dimensional space where an
electron is found 90-95% of the time.
 Each electron orbital holds up two electrons.
How are electrons distributed in an atom?
• Erwin Schrödinger (1926) – proposed a wave
equation, a mathematical expression based on
the concept that electrons show properties not
only of particles but also of waves.
Quantum Mechanics
A branch of science that studies particles
and their associated waves.
• Quantum mechanics: describes electron
energies and locations by a wave equation.
• The solution to Schrödinger equation led to a
wave function, y.
• Wave function, y - gives the shape of the
electron orbital.
• Each wave function,y, is an orbital.
• The square of the wave function y2 - gives the
probability of finding the electron.
y2 gives the electron density for the atom.
y2 is called the probability density.
• Electron cloud has no specific boundary so the
most probable area is shown.
Shapes of Atomic Orbitals for Electrons
• Based on those derived for a hydrogen atom, there are
four different kinds of orbitals for electrons: s, p, d, and
f.
• s and p orbitals are most important in organic chemistry.
spherical,
nucleus at center
dumbbell-shaped,
nucleus at middle
cloverleaf-shaped,
nucleus at middle
Orbitals and Shells
• Orbitals are grouped in shells of increasing size and
energy
• Different shells contain different numbers and kinds of
orbitals
• Each orbital can be occupied by two electrons
3rd shell contains an s orbital (3s),
three p orbitals (3p), and five d orbitals
(3d), 18 electrons
2nd shell contains one s orbital (2s)
and three p orbitals (2p), eight electrons
1st shell contains one s orbital, 1s, and
holds only two electrons
p-Orbitals
• In each shell 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, a
node.
C. Electron Configuration
Ground-state electron configuration of an atom:
• is a list of orbitals occupied by its electrons
• is the lowest-energy arrangement
• can be determined by following three “rules”
Rules for Electron Configuration
1. The Aufbau Principle. Lowest-energy orbitals fill
first: 1s  2s  2p  3s  3p  4s  3d
2. The Pauli Exclusion Principle. No more than
two electrons can occupy an orbital, and they
must be of opposite spin,  and .
3. Hund’s Rule. If two or more empty orbitals of
equal energy are available, electrons occupy
each with spins parallel until all orbitals have
one electron.
Practice Problem: Give the ground-state electron configuration
for each of the following elements:
a. Boron
b. Phosphorus
c. Oxygen
d. Chlorine
Practice Problem: How many electrons does each of the
following elements have in its outermost
electron shell?
a. Potassium
b. Aluminum
c. Krypton
II. Chemical Bonding
A.
Development of Chemical
Bonding Theory
B.
Nature of Chemical Bonding
A. Development of Chemical Bonding
Theory
• Kekulé and Couper (1858) independently
observed that carbon always has four bonds.
• Extensions to the Kekulé-Couper theory
include:
• Multiple bonding
• Ring formation
• van't Hoff and Le Bel (1874) proposed that:
 The four bonds of carbon have specific
spatial directions.
 Atoms surround carbon as corners of a
tetrahedron.
Note that a dashed line
indicates a bond is behind
the page
Note that a wedge indicates a
bond is coming forward
Practice Problem: Draw a molecule of chloroform, CHCl3, using
solid, wedged, and dashed lines to show its
tetrahedral geometry
Practice Problem: Convert the following representation of
ethane, C2H6, into a conventional drawing
that uses solid, wedged, and dashed lines to
indicate tetrahedral geometry around each
carbon
B. Nature of Chemical Bonding
• Atoms form bonds because the compound that
results is more stable (has less energy) than
the separate atoms.
• Stable molecule results at completed shell,
octet (eight dots) for main-group atoms (two for
hydrogen).
Octet Rule
• Atoms often gain, lose, or share electrons to
achieve a configuration of the noble gas
closest to them.
• Atoms tend to gain, lose, or share electrons until they are
surrounded by 8 valence electrons (4 electron pairs).
Major types of Bonds
1. Ionic Bond - results from total electron transfer.
- is an electrostatic attraction
between ions.
2. Covalent Bond - results from electron sharing.
- first described by G. N. Lewis
(1916).
Lewis Structures (or electron-dot structures)
Valence electrons – are those that reside in the
incomplete, outermost orbital of an atom
(valence shell).
Lewis structures - show valence electrons of an
atom as dots. The # of electrons available for
bonding are indicated by unpaired dots.
Drawing Lewis Structures
1. Add the valence electrons from all atoms.
2. Identify the central atom (usually the one with the
highest molecular mass and closest to the center of
the periodic table).
3. Place the central atom in the center of the molecule
and add all other atoms around it.
4. Place one bond (two electrons) between each pair of
bonded atoms.
5. Complete the octet for all atoms surrounding the
central atom.
6. Complete the octet for the central atom. Use double
or triple bonds if necessary.
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.
Valences of Carbon
• Carbon has four valence electrons (2s2 2p2),
forming four bonds (CH4).
Valences of Oxygen
• Oxygen has six valence electrons (2s2 2p4) but
forms two bonds (H2O).
Valences of Nitrogen
• Nitrogen has five valence electrons (2s2 2p3) but
forms only three bonds (NH3).
Non-bonding electrons
• Valence electrons not used in bonding are
called nonbonding electrons, or lone-pair
electrons.
 Example: Nitrogen atom in ammonia (NH3)
Kekulé Structures (or line-bond structures)
Kekulé structures - show a two-electron covalent
bond as a line drawn between atoms. Lonepairs of nonbonding valence electrons are often
not shown.
Practice Problem: What are likely formulas for the following
substances?
a. GeCl?
b. AlH?
c. CH?Cl2
d. SiF?
e. CH3NH?
Practice Problem: Write both Lewis and line-bond structures for
the following substances, showing all
nonbonding electrons:
a. CHCl3, chloroform
b. H2S, hydrogen sulfide
c. CH3NH2, methylamine
d. NaH, sodium hydride
e. CH3Li, methyl lithium
Practice Problem: Why can’t an organic molecule have the
formula C2H7?
III. Covalent Bonding
A.
Valence Bond Theory
B.
Hybridization
C.
Molecular Orbital Theory
A.
Valence Bond Theory
Valence Bond Theory
 is one of two models that describes covalent
bond formation.
 describes a covalent bond as resulting from
the overlap of two atomic orbitals
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.
Valence Bond Theory
• H–H bond results from the overlap of two singly
occupied hydrogen 1s orbitals
• H-H bond is cylindrically
symmetrical, sigma (s)
bond
Bond Energy
• Reaction 2 H·  H2 releases 436 kJ/mol
• Product has 436 kJ/mol less energy than 2H:
H–H has bond strength of 436 kJ/mol.
(1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)
Bond Length
• It is the distance
between nuclei that
leads to maximum
stability
• If too close, they repel
because both are
positively charged
• If too far apart,
bonding is weak
B.
Hybridization
• Hybridization - is a process through which a
hybrid orbital is formed.
• Hybrid orbital - is an orbital derived from a
combination of atomic orbitals.
• Hybrid orbitals, such as sp3, sp2, and sp hybrids of C, are
strongly directed and form stronger bonds than unhybridized
atomic orbitals, s or p.
Hybridization:
sp3 Orbitals and the Structure of
Methane
• Carbon has 4 valence electrons (2s2 2p2)
• In CH4, all C–H bonds are identical (tetrahedral)
even though C uses two kinds of orbitals.
How does C form four identical C–H bonds in CH4?
• sp3 hybrid orbitals: s orbital and three p orbitals
combine to form four equivalent, unsymmetrical,
tetrahedral orbitals (sppp = sp3), Linus Pauling
(1931)
Why does C form four identical C–H bonds in CH4?
• When s orbital hybridizes with three p orbitals, the
resultant sp3 hybrid orbitals are unsymmetrical
about the nucleus and strongly oriented.
• This is due to opposite algebraic signs of p lobes.
• This allows to form strong bonds.
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.
Hybridization:
sp3 Orbitals and the Structure of
Ethane
• Two C’s bond to each other by s overlap of an sp3
orbital from each
• Three sp3 orbitals on each C overlap with three H’s
1s orbitals to form six C–H bonds
Hybridization:
sp3 Orbitals and the Structure of
Ethane
• C–H bond strength in ethane 420 kJ/mol
• C–C bond is 154 pm long and strength is 376 kJ/mol
• All bond angles of ethane are tetrahedral
Hybridization: sp2 Orbitals and the Structure
of Ethylene
• sp2 hybrid orbitals: 2s orbital 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
90
120
Bonds from sp2 Hybrid Orbitals
• Two sp2-hybridized orbitals overlap to form a
sigma (s) bond
• p orbitals overlap side-to-side to form a pi ()
bond
Bonds from sp2 Hybrid Orbitals
• sp2–sp2 s bond and 2p–2p  bond share four
electrons and form C-C 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
Structure of Ethylene
• H atoms form s bonds with four sp2 orbitals
• H–C–H and H–C–C bond angles are about 120°
• C=C double bond in ethylene is shorter and stronger
than single bond in ethane
• Ethylene C=C bond length 133 pm (Ethane C–C 154 pm)
Hybridization: sp Orbitals and the Structure
of Acetylene
• sp hybrid orbitals: Carbon 2s orbital hybridizes
with a single p orbital giving two sp hybrids
• two p orbitals remain unchanged
Hybridization: sp Orbitals and the Structure
of Acetylene
• sp orbitals are linear, 180° apart on x-axis
• Two p orbitals are perpendicular on the y-axis and
the z-axis
Orbitals of Acetylene
• Two sp hybrid orbitals from each C form sp–sp
s bond
• pz orbitals from each C form a pz–pz  bond by
sideways overlap and py orbitals overlap
similarly to form py–py  bond
Bonding in Acetylene
• Sharing of six electrons forms C  C
• Two sp orbitals form s bonds with hydrogens
Hybridization:
Nitrogen and Oxygen
• Atoms such as nitrogen and oxygen hybridize to
form strong, oriented bonds
• Nitrogen in NH3
• Oxygen in H2O
Hybridization of Nitrogen in Ammonia
• H–N–H bond angle in
ammonia (NH3) is 107.3°
• N’s orbitals (sppp) hybridize
to form four sp3 orbitals
• One sp3 orbital is occupied
by two nonbonding
electrons, and three sp3
orbitals have one electron
each, forming bonds to H
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°
To assign hybridization:
• Draw the Lewis structure for the molecule or ion
• Determine the electron domain geometry
• Specify the hybrid orbitals needed to accommodate
the electron pairs based on their geometric
arrangement
– Example: the hybrid orbitals used by N in NH3 molecule are
predicted by first writing the Lewis structure, and finally the
geometry.
To assign hybridization:
# attached atoms
+ # nonbonding
electron pairs
(nondelocalized)
Hybridization
Bond Angles
Molecular Geometry
2
sp
180o
Linear
3
sp2
120o
Trigonal Planar
109.5o
Tetrahedral,
pyramidal,
or bent
4
sp3
Practice Problem: Draw a line-bond structure for propane,
CH3CH2CH3. Predict the value of each bond
angle, and indicate the overall shape of the
molecule
Practice Problem: Convert the following molecular model of
hexane, a component of gasoline, into a linebond structure
Practice Problem: Draw a line-bond structure for propene,
CH3CH=CH2, indicate the hybridization of
each carbon, and predict the value of each
bond angle
Practice Problem: Draw a line-bond structure for 1,3-butadiene
H2C=CH-CH=CH2: indicate the hybridization
each carbon, and predict the value of each
bond angle
Practice Problem: Draw both a Lewis structure and a line-bond
structure for acetaldehyde, CH3CHO
Practice Problem: Shown below is a molecular model of aspirin
(acetylsalicyclic acid). Identify the
hybridization of each carbon atom, and tell
which atoms have lone pairs of electrons
Practice Problem: Draw a line-bond structure for propyne,
CH3CCH2, indicate the hybridization of
each carbon, and predict the value of each
bond angle
Practice Problem: Draw Lewis and line-bond structures for
formaldimine,CH2NH. How many electrons
are shared in the carbon-nitrogen bond?
What is the hybridization of the nitrogen
atom?
Practice Problem: What geometry do you expect for each of the
following atoms?
a. The oxygen atom in methanol
b. The nitrogen atom in trimethylamine
c. The phosphorus atom in :PH3
C.
Molecular Orbital Theory
Molecular Orbital Theory
 is one of two models that describes covalent
bond formation.
 describes a covalent bond as resulting from
combination of atomic orbitals to give
molecular orbitals, which belong to the entire
molecule
• Molecular Orbital (MO) - is a region where
electrons are most likely to be found in a molecule.
• It has a specific energy and general shape
• There are two types of orbital combination:
• Additive combination
• Substractive combination
• Additive combination: (bonding) MO is lower in
energy
• Substractive combination: (antibonding) MO is
higher in energy
Molecular Orbitals in Ethylene
• The  bonding MO is from combining p orbital
lobes with the same algebraic sign
• The  antibonding MO is from combining
lobes with opposite signs
• Only bonding MO is occupied
Chapter 1
Summary
• Organic chemistry – chemistry of carbon
compounds
• Atom: positively charged nucleus surrounded by
negatively charged electrons
• Electronic structure of an atom described by wave
equation
– Electrons occupy orbitals around the nucleus.
– Different orbitals have different energy levels and
different shapes
• s orbitals are spherical, p orbitals are dumbbellshaped
Summary
• Covalent bonds - electron pair is shared
between atoms
• Valence bond theory - electron sharing occurs
by overlap of two atomic orbitals
• Molecular orbital (MO) theory - bonds result
from combination of atomic orbitals to give
molecular orbitals, which belong to the entire
molecule
Summary
• Sigma (s) bonds - Circular cross-section and
are formed by head-on interaction
• Pi () bonds – “dumbbell” shape from
sideways interaction of p orbitals
Summary
• 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
Summary
• Atoms such as nitrogen and oxygen hybridize
to form strong, oriented bonds
– The nitrogen atom in ammonia and the oxygen
atom in water are sp3-hybridized