Chapter 1
ORGANIC CHEMISTRY
STUDY
OF
CARBON
CONTAINING
COMPOUNDS
Compounds from Nature
Synthetic compounds: invented by organic chemists and prepared in
their laboratories
1828
Friedrich Woehler’s urea synthesis
Ammonium isocyanate + heat ------> urea
NH4CNO
NH2CONH2
1828
“I have been able to make urea without aid of kidney of man or dog”.
Some organic chemicals
Medicines
DNA
Active Pharmaceutical Ingredients•
Excipients•
Fuels
Materials
Essential oils
Pigments
AND
C
C
C
C
WITH ITSELF
C
C
CHAINS
C
CH
C
C
C
C
C
CH
C
C
C
C
RINGS
CHAINS WITH BRANCHES
No limit
Electronic Structure of Atoms
• Structure of atoms
– a small dense nucleus,
diameter 10-14 - 10-15 m,
which contains
positively charged
protons, neutrons and
most of the mass of
the atom
– extranuclear space,
diameter 10-10 m,
which contains
negatively charged
electrons
10-10 m
N ucleus (proton s
and neutron s)
Sp ace occup ied
by electron s
Proton
N eutron
10-15 m
Notice: one s orbital in each principal shell
three p orbitals in the second shell (and in higher ones)
five d orbitals in the third shell (and in higher ones)
Rules for Electron
Configurations
Capacities of shells (n) and subshells (l)
Electronic Structure of Atoms
• Electrons are confined to regions of space called
principle energy levels (shells)
– each shell can hold 2n2 electrons (n = 1, 2, 3, 4......)
S h e ll
4
3
2
1
S hell
Nu m be r o f
Re la t ive En e rg i e s
Ele ctro n s S h e llo f Ele ctro n s
C a n H ol d
i n Th es e S h e l ls
h ig h e r
32
18
8
2
l owe r
Orbitals Contained in That Sh ell
3
3s, 3px , 3p y, 3p z , p lus five 3d orb itals
2
1
2s, 2px , 2p y, 2p z
1s
Electronic Structure of Atoms
• Shells are divided into subshells called orbitals,
which are designated by the letters s, p, d,........
– s (one per shell)
– p (set of three per shell 2 and higher)
– d (set of five per shell 3 and higher) .....
S hell
Orbitals Contained in That Sh ell
3
3s, 3px , 3p y, 3p z , p lus five 3d orb itals
2
1
2s, 2px , 2p y, 2p z
1s
Electronic Structure of Atoms
• Rule 1: orbitals fill
from lowest energy to
highest energy
• Rule 2: only two
electrons per orbital,
spins must be paired
• Rule 3: for a set of
orbitals with the same
energy, add one
electron in each before
a second is added in
any one
“Periodic” Behavior of Elements
Flame tests: elements with low first ionization energies are excited
in a flame, and often emit in the visible region of the spectrum
Li
Na
Ca
Sr
K
Ba
Atoms emit energy when electrons fall from higher to lower energy states
Atomic Spectrum of Hydrogen
Electronic Structure of Atoms
• The pairing of electron spins
Lewis Structures
For Nitrogen atom:
Valence shell of Nitrogen= 3
Number of valence electrons of Nitrogen = 5
• Gilbert N. Lewis
• Valence shell: the outermost electron shell of an
atom
• Valence electrons: electrons in the valence shell
of an atom; these electrons are used in forming
chemical bonds
• Lewis structure
– the symbol of the atom represents the nucleus and
all inner shell electrons
– dots represent valence electrons
Lewis Structures
• Lewis structures for elements 1-18 of the Periodic Table
For Nitrogen atom:
Valence shell of Nitrogen= 3
Number of valence electrons= 5
Lewis Model of Bonding
• Atoms bond together so that each atom in the bond
acquires the electron configuration of the noble gas
nearest it in atomic number
– an atom that gains electrons becomes an anion
– an atom that loses electrons becomes a cation
– Ionic bond: a chemical bond resulting from the
electrostatic attraction of an anion and a cation
– Covalent bond: a chemical bond resulting from two
atoms sharing one or more pairs of electrons
• We classify chemical bonds as ionic, polar covalent,
and nonpolar covalent based on the difference in
electronegativity between the atoms
Electronegativity
• Electronegativity: a measure of the force of an
atom’s attraction for the electrons it shares in a
chemical bond with another atom
• Pauling scale
– increases from left to right within a period
– increases from bottom to top in a group
Electronegativity
Electronegativity of atoms (Pauling scale)
Electronegativity
• Electronegativity and chemical bonding
D i fference in
Electronegativity
Betw een Bonded Atoms
less than 0.5
0.5 to 1.9
greater than 1.9
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic
bond!
Ty pe o f Bo nd
no npo lar co val ent
po lar coval ent
ioni c
Coulomb’s Law
“The energy of interaction between a pair of ions is
proportional to the product of their charges,
divided by the distance between their centers”
19
E  (2.31x 10
 Q1Q2 
J  nm)

 r 
What forces that hold atom together within molecules?
Covalent Bonding Forces
H•
+
•H
H-H
H0 = -104 kcal/mol (-435 kJ/mol)
 Electron – electron
repulsive forces
 Proton – proton
repulsive forces
 Electron – proton
attractive forces
Bond Length Diagram
Scientists can determine the internuclear
distances that correspond to the lowest energy
states of molecules
Net repulsion
Net attraction
http://ch301.cm.utexas.edu/simulations/bond-strength/BondStrength.swf
Bond Length and Energy
Bond
C-C
C=C
CC
C-O
C=O
C-N
C=N
CN
Bond type
Single
Double
Triple
Single
Double
Single
Double
Triple
Bond length
(pm)
Bond Energy
(kJ/mol)
154
134
120
347
614
839
143
358
123
143
745
305
138
116
615
891
Bonds between elements become shorter and stronger as
multiplicity increases
Covalent Bonds
• A covalent bond forms when electron pairs are
shared between two atoms whose difference in
electronegativity is 1.9 or less
– an example is the formation of a covalent bond
between two hydrogen atoms
– the shared pair of electrons completes the valence
shell of each hydrogen.
H•
+
•H
H-H
H0 = -104 kcal/mol (-435 kJ/mol)
Polar Covalent Bonds
• In a polar covalent bond
– the more electronegative atom has a partial negative
charge, indicated by the symbol d– the less electronegative atom has a partial positive
charge, indicated by the symbol d+
• in an electron density model
– red indicates a region of high electron density
– blue indicates a region of low electron density
Polar and Nonpolar Molecules
– ammonia and formaldehyde are polar molecules
– acetylene is a nonpolar molecule
In sert elpot of
ammon ia
(page 19)
In sert elpot of
acetylene
(page 20)
dO
dN
H
In sert elpot of
formaldeh yd e
(page 20)
H
H d+
A mmon ia
(p olar)
C
H d+ H
H C C H
Formald ehyde
(p olar)
Acetylene
(nonpolar)
Carbon – Intro and Review
• Atomic Structure
– Atoms – made up of protons, neutrons, electrons
– Isotopes – same # protons; different # neutrons
• Electronic Structure
– Electrons
• determine structure
• give rise to bonding
• behave like waves
• orbitals (s, p)
12
14
6
6
C
C
Orbital overlap to form σ bonds.
Orbital overlap to form p bonds.
Electron Probabilities
and the 1s Orbital
The 1s orbital looks very much like a fuzzy ball,
that is, the orbital has spherical symmetry
The electrons are more concentrated near the center
Spherical symmetry;
probability of finding
the electron is the same
in each direction.
The electron
cloud doesn’t
“end” here …
… the electron just
spends very little
time farther out.
Electron Probabilities
and the 2s Orbital
The 2s orbital has two regions of high electron
probability, both being spherical
The region near the nucleus is separated from the
outer region by a spherical node - a spherical shell in
which the electron probability is zero
The Three p Orbitals
2p
The Five d Orbitals
3d
Rules for Electron
Configurations
Subshell filling order ...
Each subshell must be
filled before moving
to the next level
1s22s22p63s23p6 ...
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
The most stable arrangement of electrons
in subshells is the one with the greatest
number of parallel spins (Hund’s rule).
Ne 1s22s22p6
F 1s22s22p5
O 1s22s22p4
N 1s22s22p3
C 1s22s22p2
Periodic Relationships
The valence shell is the outermost occupied shell
The period number = principal quantum
number, n, of the electrons in the valence shell
Atomic Orbitals
• 1s – 1st orbital
– s type (spherical)
– 1s, 2s, 3s
Atomic Orbitals
2s orbital (spherical)
Atomic Orbitals
• p (2p, 3p…)
– 3 orbitals oriented perpendicular to each
other
– have node (region of 0 e- density)
• nodal plane
2p orbital
Atomic
Orbitals
• p (2p, 3p…)
– 3 orbitals oriented perpendicular to each other
– have node (region of 0 e- density)
• nodal plane
– shape
• dumbbell
Electronic Configuration of Atoms
• Aufbau
– Fill lowest energy orbital 1st
• Hund’s Rule
– 1 e- into each orbital of = energy

Pauli Exclusion
Principle

Electrons in the
same orbital are
spin paired



Chapter 1

Electronic Configurations
WHY DO HYBRIDS ??
• 1. Electron pair repulsions are minimized (= lower
energy)
• 2. Stronger bonds (= lower energy) are formed
• 3. Hybrids have better directionality for forming
bonds
Shapes of Atomic Orbitals
• All s orbitals have the shape of a sphere, with its center at the nucleus
– of the s orbitals, a 1s orbital is the smallest, a 2s orbital is larger, and a 3s
orbital is larger still
Shapes of Atomic Orbitals
– A p orbital consists of two lobes arranged in a
straight line with the center at the nucleus
Orbital Overlap Model
• A covalent bond forms when a portion of an
atomic orbital of one atom overlaps a portion of
an atomic orbital of another atom
– in forming the covalent bond in H-H, for example,
there is overlap of the 1s orbitals of each hydrogen
Hybrid Orbitals
• We will study three types of hybrid atomic
orbitals
sp3 (one s orbital + three p orbitals give four sp3
orbitals)
sp2 (one s orbital + two p orbitals give three sp2
orbitals)
sp (one s orbital + one p orbital give two sp orbitals)
• Overlap of hybrid orbitals can form two types of
bonds, depending on the geometry of the
overlap
 bonds are formed by “direct” overlap
p bonds are formed by “parallel” overlap
sp3 Hybrid Orbitals
– Each sp3 hybrid orbital has two lobes of unequal
size
– The four sp3 hybrid orbitals are directed toward the
corners of a regular tetrahedron at angles of 109.5°
sp3 Hybrid Orbitals
– orbital overlap bonding in water, ammonia, and
methane
2
sp
Hybrid Orbitals
• An sp2 hybrid orbital has two lobes of unequal size
– the three sp2 hybrid orbitals are directed toward the
corners of an equilateral triangle at angles of 120°
– the unhybridized 2p orbital is perpendicular to the
plane of the three sp2 hybrid orbitals
sp2 Hybrid Orbitals
– a carbon-carbon double bond consists of one
sigma () bond and one pi (p) bond
2
sp
Hybrid Orbitals
– a carbon-oxygen double bond also consists of one
sigma () bond and one pi (p) bond
sp Hybrid Orbitals
• Each sp hybrid orbital has two lobes of unequal size
– the two sp hybrid orbitals lie in a line at an angle of 180°
– the two unhybridized 2p orbitals are perpendicular to each
other and to the line through the two sp hybrid orbitals
sp Hybrid Orbitals
– a carbon-carbon triple bond consists of one sigma
() bond and two pi (p) bonds
Hybrid Orbitals
• Summary of orbitals and bond types
Hybridization
sp
3
sp 2
Types of
Bonds to Carbon
fou r s igma bond s
Example
HH
H-C-C-H
HH
H
three sigma bonds
and on e pi bond
tw o sigma b on ds
and tw o p i bonds
Ethan e
H
C
H
sp
N ame
H-C
Ethylene
C
H
C-H Acetylene
Examples of sigma σ bonds formed from
sp3 hybrid orbitals
Orbital overlap to form σ bonds.
Orbital overlap to form p bonds.
..
H
H C H
H N H
H
H
H
H
C H
H
H
N H
H
H
H
..
O
..
O
H
H
Examples of natural acyclic compounds, their sources
(in parentheses), and selected characteristics
Examples of natural heterocyclic compounds having
a variety of heteroatoms and ring sizes.
Examples of natural carbocyclic compounds with rings of
various sizes and shapes.
Isomerism
 The Molecular Formula of a substance gives the number
of different atoms present.
 The Structural Formula indicates how those atoms are
arranged.
 Isomers are molecules with the same number and kinds
of atoms but different arrangements of the atoms.
 Structural (or Constitutional) isomers have the same
molecular formula but different structural formulas.
Constitutional Isomerism
– the potential for constitutional isomerism is enormous
Mol e cu l ar C on s ti tu tion al
Form u la
Isome rs
CH4
C5 H1 2
1
3
C1 0 H2 2
75
C1 5 H3 2
4,347
C2 5 H5 2
36,797,588
C3 0 H6 2
4,111,846,763
World population
is about
6,000,000,000
Condensed Structural Formulas
74
Cyclic Molecules
75
Bond-line Formulas
76
77
Three-Dimensional Formulas
In this representation, bonds that project upward out
of the plane of the paper are indicated by a wedge,
those that lie behind the plane are indicated with a
dashed wedge, and those bonds that lie in the plane
of the page are indicated by a line.
78
writing structural Formulas
 In a continuous chain, atoms are bonded one after
another.
 In a branched chain, some atoms form branches from the
longest continuous chain.
Abbreviated Structural Formulas
Formal Charge
 Here, some molecules one or more atoms maybe charged
+ve or –ve which comes from the chemical reactions.
 Its important to know how to tell where the charge is
located.
H3
+
O
Formal Charge
 The formal charge on an atom in a covalently bonded
molecule or ion is the number of valence electrons in the
neutral atom minus the number of covalent bonds to the
atom and the number of unshared electrons on the atom.
Resonance
 Resonance structures of a molecule or ion are two or more
structures with identical arrangements of the atoms but
different arrangements of the electrons.
 If resonance structures can be written, the true structure of
the molecule or ion is a resonance hybrid of the
contributing resonance structures.
Resonance
 Physical measurements tell us that none of the foregoing
structures accurately describes the real carbonate ion.
 Experimentally, It was found that all three carbon–oxygen
bond lengths are identical: 1.31 Å. This distance is
intermediate between the normal C=O (1.20 Å) and C-O
(1.41 Å)
 The real carbonate ion has a structure that is a resonance
hybrid of the three contributing resonance structures