ORGANIC CHEMISTRY
171
Section 201
Chapter one
ORIGIN OF ORGANIC CHEMISTRY
• In the earlier period of development of
chemistry, chemists tried their best to
synthesize organic compounds in the
laboratory,but failed.Their failures led them to
believe that organic compounds could be
prepared only from living organisms ( plants
and animals ) through force called vital force
and theory known as Vital force theory.
In
1828 Freidrich Wohler, a German chemist
accidentally obtained urea, (NH2)2 CO, an organic
compound found in the urine of mammals. In fact,
Wohler tried to prepare ammonium cyanate, a
substance with mineral origin, by heating ammonium
sulphate and potassium cyanate, ammonium cyanate
rearranged to urea, a compound which was of organic
nature.
Freidrich Wohler
NH4CNO
Ammonium cyanate
HEATING
NH2CONH2
Urea
This chance discovery of Wohler brought about a
revolution in the field of organic chemistry.
Hennel (1828) prepared ethyl alcohol,
Berthelot (1856) prepared methane etc. in the
laboratory from mineral resources. By the
middle of nineteenth century, the Vital force
theory was completely discarded. Chemists
then never looked back and at present about
ninety five per cent of the organic compounds
are man-made.
Organic Chemistry
• The study of the compounds of carbon.
• Over 10 million compounds have been identified.
– About 1000 new ones are identified each day!
• C is a small atom.
– It forms single, double and triple bonds.
– It is intermediate in electronegativity (2.5).
– It forms strong bonds with C, H, O, N, and some
metals.
Chemical Bonds
Bonds :
forces that hold groups of atoms together
and make them function as a unit.
,determines chemical and
physical properties of
substances
Types of Chemical Bonds
• Ionic Bond
– electrostatic attraction between closely packed
oppositely charged ions
– formed when a metal reacts with a nonmetal
• a metal loses electrons easily to form a positive
ion which is attracted to the negative ion formed
when the nonmetal gains electrons.
+
F
••
+
Na
••
-
•
•
•
•
Na
•
•
••
F
••
– In forming Na+F-, the single 3s electron from Na is
transferred to the partially filled valence shell of F.
Na(1 s 22s 22p 63s 1) + F(1s 2 2s 2 2p 5 )
Na + (1s 22s 22p 6) + F-(1s 2 2s 2 2p 6 )
Covalent Bonding
– electrons are shared between nuclei
• electrons are shared equally between identical atoms.
• The simplest covalent bond is that in H2
H•
•
•
+
•H
H-H
H 0 = -435 kJ (-104 kcal)/mol
• polar covalent bond
• electrons may be shared unequally between two different atoms.
Or Unequal sharing of electrons between atoms in a molecule.(as HF and HCl)
Results in a charge separation in the bond (partial positive and partial negative
charge).
D i fference in
El ectron eg ati vity
Betw een Bo nded Ato ms
Less than 0.5
0.5 to 1.9
Greater than 1.9
Typ e of Bond
N on pol ar cov alent
Pol ar co valent
Io ns f orm
Electron Configuration of Atoms
• Shells are divided into subshells called orbitals, which
are designated by the letters s, p, d, f........
– s (one per shell)
– p (set of three per shell 2 and higher)
– d (set of five per shell 3 and higher) ..…
– The distribution of Orbitals in Shells
S hell
O rb itals Contain ed in Th at S hell
3
3s , 3p x , 3p y , 3p z, p lu s five 3d orbitals
2
2s , 2p x , 2p y , 2p z
1
1s
A molecule is an aggregate of two or more atoms in a definite
arrangement held together by chemical bonds
Molecule, is the collection of two atoms by sharing electrons
H2
H2O
NH3
CH4
A diatomic molecule contains only two atoms
H2, N2, O2, Br2, HCl, CO
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4
2.5
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
Cl-
17 protons
18 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
2.5
Electronegativity
• Electronegativity
– the ability of an atom in a molecule to attract shared
electrons to itself.
• Pauling scale
– Generally increases left to right in a row.
– Generally increases bottom to top in a column.
Electron Configuration of Atoms
• Aufbau (“Build-Up”) Principle:
– Orbitals fill in order of increasing energy from lowest energy
to highest energy.
• Pauli Exclusion Principle:
– No more than two electrons may be present in an orbital. If
two electrons are present, their spins must be paired.
• Hund’s Rule:
– one electron is added to each orbital of equal energy before a
second electron is added to any one of them; the spins of the
electrons should be aligned.
Electron Configuration
• Energy-level diagram; electrons are placed in an
electron configuration. For example, the energy-level
diagram for the ground-state electron configuration of
carbon is 1s2 2s2 2p2. For chlorine: 1s2 2s2 2p6 3s2 3p5.
3p
2p
Energy
Energy
2s
3s
2p
2s
1s
Energy-level
diagram for carbon
(atomic number 6)
1s
Energy-level
diagram for chlorine
(atomic number 17)
Lewis Dot Structures
• Table 1.4 Lewis Dot Structures for Elements 1-18
1A
Na
4A
5A
6A
7A
.
8A
.
.
Be
Mg
:
:
.
B
.
Al
:
.
C:
.
.
. N. :
.
: O. :
.
:..F :
:
.
Si :
.
.
.P :
.
.
: S. :
.
:Cl :
:
: :
He
:N e :
: :
Li
3A
:
H
2A
:A r :
Formal Charge
• Formal charge: The charge on an atom in a molecule or
a polyatomic ion.
• To derive formal charge
1. Write a correct Lewis structure for the molecule or ion.
2. Assign each atom all its unshared (nonbonding) electrons and
one-half its shared (bonding) electrons.
3. Compare this number with the number of valence electrons in
the neutral, unbonded atom.
Formal
charge
N umber of
= valence electrons
in th e neutral,
un bonded atom
All
One h alf of
un shared + all sh ared
electrons
electrons
4. The sum of all formal charges is equal to the total charge
on the molecule or ion.
Formal Charge
• Example: Draw Lewis structures, and show which
atom in each bears the formal charge.
(a) NH2
-
(d) CH3 NH3
(b) HCO3
+
-
(c) CO3
-
(e) HCOO
2-
-
(f) CH3 COO
Structural formula of organic compounds
The structural formula of a chemical compound is a graphical
representation of the molecular structure, showing how the atoms
are arranged. The chemical bonding within the molecule is also
shown.
The structure of carbon and its compound can be
expressed using the Lewis-dot structure This system
identifies how the atoms in a molecule of a specific
compound are attached (bonded) to one another oriented
in space. An atom is indicated by its symbol with a
number of dots representing the number of valency
electrons e.g Hydrogen would be H with a single dot,
Carbon would be a C with four dots
When two valency electrons are paired they are represented by
two adjacent dots. Paired valency electrons are not normally
available for forming bonds with other atoms.
H-O-H
H 2 O (8)
Water
H
H-N-H
H
N H 3 (8)
Ammonia
H
C
C
H
H
C2 H 4 (12)
Ethylene
H
H-C-H
H
CH 4 (8)
Meth ane
H-Cl
HCl (8)
Hyd rogen ch loride
O
H
H-C C-H
C O
C 2H 2 (10)
Acetylen e
H
CH 2O (12)
Formald ehyde
H
O
C
O
H
H 2CO 3 (24)
Carbonic acid
In neutral molecules
hydrogen has one bond.
carbon has 4 bonds and no lone pairs.
nitrogen has 3 bonds and 1 lone pair.
oxygen has 2 bonds and 2 lone pairs.
halogens have 1 bond and 3 lone pairs.
Isomers
• Two compounds are considered isomers if they have the
same molecular formula (i.e. the same numbers and types
of atoms) but different structures.
• There are two types of isomers, structural isomers and
stereoisomers .
• structural isomers If two compounds are considered to
have the same molecular formula but different connections
between atoms (bonding).
• stereoisomers If two compounds are considered to have the
same molecular formula, the same connections between
atoms, but different arrangements of the atoms in three
dimensional space. Geometric isomerism (also known as cis-trans
isomerism or E-Z isomerism) is a form of stereoisomerism.
Consider compounds with the molecular formula C4H10.
There are two valid structures that fit this molecular
formula.They differ in physical properties.
A linear structure: n-butane
A branched structure: iso-butane (or 2-methylpropane)
CH3
CH3 CH2 CH2 CH3
Butane
CH3 CHCH3
Is ob utane
CH3
CH3 CH2 CH2 CH2 CH3 CH3 CH2 CHCH3
Pentan e
Is op entane
CH3
CH3 CCH3
CH3
N eopentan e
Two isomers of butane C4H10:
CH3CH2CH2CH3
n-butane
bp 0 oC
mp –138 oC
d 0.622 g/cc
CH3
CH3CHCH3
bp -12 oC
mp -159 oC
d 0.604 g/cc
isobutane
Resonance
• Linus Pauling - 1930s
– Many molecules and ions are best described by
writing two or more Lewis structures.
– Individual Lewis structures are called contributing
structures.
– Connect individual contributing structures by doubleheaded (resonance) arrows.
– The molecule or ion is a hybrid of the various
contributing structures.
Resonance
:
:
• For many molecules and ions, no single Lewis structure
provides a truly accurate representation.
O:
:O :
H3 C C
and
H3 C C
O:
-
:
:
:O:
-
Ethanoate ion
(acetate ion)
Resonance
• Examples: equivalent contributing structures.
CH3
O:
CH3
O:
C
: O :-
:
N itrite ion
(equivalent con trib uting
s tru ctures)
C
:
:O :-
:
:
O:
: O :-
O:
:N
:
:N
:
:
:
:O: -
A cetate ion
(equ ivalen t contributin g
s tru ctures)
Resonance
• Curved arrow: A symbol used to show the redistribution
of valence electrons.
• In using curved arrows, there are only two allowed
types of electron redistribution:
– from a bond to an adjacent atom.
– from a lone pair on an atom to an adjacent bond.
Resonance
• All contributing structures must
1. have the same number of valence electrons.
2. obey the rules of covalent bonding:
– no more than 2 electrons in the valence shell of H.
– no more than 8 electrons in the valence shell of a 2nd
period element.
3. differ only in distribution of valence electrons; the
position of all nuclei must be the same.
4. have the same number of paired and unpaired
electrons.
Resonance
• The carbonate ion
– Is a hybrid of three equivalent contributing
structures.
– The negative charge is distributed equally among the
three oxygens.
Resonance
• Preference 1: filled valence shells
– Structures in which all atoms have filled valence
shells contribute more than those with one or more
unfilled valence shells.
CH3
+
O
••
C
H
H
Greater contribution;
both carbon and oxygen have
complete valence shells
••
CH3 O
••
+
C
H
H
Les ser contribution;
carbon has only 6 electrons
in its valence shell
Resonance
• Preference 2: maximum number of covalent
bonds
– Structures with a greater number of covalent bonds
contribute more than those with fewer covalent
bonds.
CH3
+
O
••
••
C
H
H
Greater contribution
(8 covalent bonds)
CH3 •O
•
+
C
H
H
Les ser contribution
(7 covalent bonds)
Resonance
• Preference 4: negative charge on the more
electronegative atom.
– Structures that carry a negative charge on the more
electronegative atom contribute more than those
with the negative charge on a less electronegative
atom.
O
(1)
C
H3 C
O
O
CH3
(a)
Less er
con trib ution
(2)
C
H3 C
CH3
(b)
Greater
contribu tion
C
H3 C
CH3
(c)
S hould n ot
be d raw n
Shapes of Atomic s and p Orbitals
• All s orbitals have the
shape of a sphere with
the center of the sphere
at the nucleus.
Shapes of Atomic s and p Orbitals
– Figure 1.9 (a) 3D representations of the 2px,
2py, and 2pz atomic orbitals including nodal
planes.
Shapes of Atomic s and p Orbitals
• Figure 1.9(b) Cartoon representations of the 2px, 2py,
and 2pz atomic orbitals.
Hybridization
The tetravalency shown by carbon is actually due to excited state
of carbon which is responsible for carbon bonding capacity.
If the bond formed is by overlapping then all the bonds will not
be equivalent so a new concept known as hybridization is
introduced which can explain the equivalent character of bonds.
s and p orbital belonging to the same atom having slightly different energies
mix together to produce same number of new set of orbital called as hybrid
orbital and the phenomenon is called as hybridization.
Important characteristics of hybridization
(i)The number of hybridized orbital is equal to number
of orbitals that get hybridized.
(ii) The hybrid orbitals are always equivalent in energy
and shape.
(iii) The hybrid orbitals form more stable bond than
the pure atom orbital.
(iv) The hybrid orbitals are directed in space in same
preferred direction to have some stable arrangement
and giving suitable geometry to the molecule.
Types Of Hybridization
• There are the three kinds of hybridization that
are important in organic chemistry.
• (i) sp3 hybridization: In this case, one s and
three p orbitals hybridise to form four sp3 hybrid
orbitals. These four sp3 hybrid orbitals are
oriented in a tetrahedral arrangement.For
example in methane CH4
(ii) sp2 hybridization: In this case one s and two p
orbitals mix together to form three sp2 hybrid orbitals
and are orientedin a trigonal planar geometry.
Take, for example, ethene (C2H4). Ethene has a
double bond between the carbons.
For this molecule, carbon will sp2 hybridise,
because one π (pi) bond is required for the double
bond between the carbons, and only three σ bonds
are formed per carbon atom. In sp2 hybridisation
the 2s orbital is mixed with only two of the three
available 2p orbitals:
forming a total of 3 sp2 orbitals with one p-orbital
remaining. In ethylene (ethene)or ehtylene the two
carbon atoms form a σ bond by overlapping two sp2
orbitals and each carbon atom forms two covalent
bonds with hydrogen by s–sp2 overlap all with 120°
angles. The π bond between the carbon atoms
perpendicular to the molecular plane is formed by 2p–
2p overlap. The hydrogen-carbon bonds are all of
equal strength and length, which agrees with
experimental data.
(iii) sp hybridization: In this case, one s and one p
orbital mix together to form two sp hybrid orbitals
and are oriented in a linear shape.
In this model, the 2s orbital mixes with only one of
the three p-orbitals resulting in two sp orbitals and
two remaining unchanged p orbitals. The chemical
bonding in acetylene (ethyne) (C2H2) consists of sp–
sp overlap between the two carbon atoms forming a σ
bond and two additional π bonds formed by p–p
overlap. Each carbon also bonds to hydrogen in a
sigma s–sp overlap at 180° angles.