Structure and

Bonding
What is organic chemistry?
⚫ The study of carbon and its compounds.
⚫ living organisms
⚫ proteins
⚫ DNA
⚫ food
⚫ clothes
⚫ medicines

2
Why organic chemistry?
⚫ carbon has unique chemistry
⚫ bonds to every other element
⚫ bonds to itself in long chains
⚫ organic chemistry involves enormous variety
⚫ many possible structures
⚫ many possible reactions

3
Periodic Table

4
 H N O F
C P S Cl
Organic Chemistry Br

Today : I

⚫ Chemistry of the compounds of carbon…( …

and H N O P S F Cl Br I )
Originally :
⚫ Chemistry of compounds derived from plant
and animal (“organic”) sources
⚫ “Chemistry of Life”

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Origin of Modern Organic
Chemistry
VITALISM
⚫ (prominent idea of the 1700’s)

⚫ Belief that certain chemicals,

ORGANIC CHEMICALS,
⚫ could only be made by living organisms.

INORGANIC CHEMICALS
⚫ were found primarily in the earth as mineral
deposits, but could also be prepared by man.

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Origin of Modern Organic
Chemistry
⚫ The study of carbon containing compounds
First organic synthesis, 1828.

Friedrich Wohler
1800 - 1882

….. defeated vitalism 7

Atomic Structure
⚫ Structure of an atom
⚫ Positively charged nucleus (very dense, protons and
neutrons) and small (10-15 m)
⚫ Negatively charged electrons are in a cloud (10-10 m)
around nucleus
⚫ Diameter is about 2  10-10 m (200 picometers (pm))
[the unit angstrom (Å) is 10-10 m = 100 pm]

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Mass Number
⚫ Mass Number: total number of protons +
neutrons (nucleons) in an atom.
Mass Number = # protons + # neutrons
mass number = 6 p + 6 n = 12 amu

12
6
C
number of protons

12
6C
Carbon-12,
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Atomic Structure: Orbitals
⚫ Quantum mechanics: describes electron energies
and locations by a wave equation
⚫ Wave function solution of wave equation
⚫ Each Wave function is an orbital,
⚫ A plot of 2 describes where electron is most likely
to be located
⚫ Electron cloud has no specific boundary so we show
most probable area

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Shapes of Atomic Orbitals for
Electrons
⚫ Four different kinds of orbitals for electrons based
on those derived for a hydrogen atom
⚫ Denoted s, p, d, and f
⚫ s and p orbitals most important in organic chemistry
⚫ s orbitals: spherical, nucleus at center
⚫ p orbitals: dumbbell-shaped, nucleus at middle

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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
⚫ First shell contains one s orbital, denoted 1s, holds only two electrons
⚫ Second shell contains one s orbital (2s) and three p orbitals (2p), eight
electrons
⚫ Third shell contains an s orbital (3s), three p orbitals (3p), and five d
orbitals (3d), 18 electrons

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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

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Atomic Structure: Electron
Configurations
⚫ Ground-state electron configuration of an atom lists orbitals
occupied by its electrons. Rules:
1. Lowest-energy orbitals fill first: 1s → 2s → 2p → 3s → 3p → 4s
→ 3d (Aufbau (“build-up”) principle)
2. Electron spin can have only two orientations, up  and down .
Only two electrons can occupy an orbital, and they must be of
opposite spin (Pauli exclusion principle) to have unique wave
equations
3. If two or more empty orbitals of equal energy are available,
electrons occupy each with spins parallel until all orbitals have
one electron (Hund's rule).

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Electron Structure of Atoms

15
Examples-Orbital Diagram

1s 2s
H 
He 
Li  
Be  

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Electronic Configuration
1s 2s 2p Configuration

Li   1s22s1

Be   1s22s2

B    1s22s22p1

HUND’S RULE. When placing electrons in

a set of orbitals having the same energy, we
place them singly as long as possible.
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Valence Electrons

Electrons are divided between core and

valence electrons.
Na 1s2 2s2 2p6 3s1
Core = [Ne] and valence = 3s1

Br [Ar] 3d10 4s2 4p5

Core = [Ar] 3d10 and valence = 4s2 4p5
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Development of Chemical
Bonding Theory
⚫ Kekulé and Couper independently observed that
carbon always has four bonds
⚫ van't Hoff and Le Bel 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 19
The Nature of the Chemical Bond
⚫ Atoms form bonds because the compound that results is more
stable than the separate atoms
⚫ Ionic bonds in salts form as a result of electron transfers
⚫ Organic compounds have covalent bonds from sharing electrons
(G. N. Lewis, 1916)
⚫ Lewis structures shown valence electrons of an atom as dots
⚫ Hydrogen has one dot, representing its 1s electron
⚫ Carbon has four dots (2s2 2p2)
⚫ Stable molecule results at completed shell, octet (eight dots) for
main-group atoms (two for hydrogen)

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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

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Valences of Carbon
⚫ Carbon has four valence electrons (2s2 2p2), forming
four bonds (CH4)

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Valences of Oxygen
⚫ Oxygen has six valence electrons (2s2 2p4) but
forms two bonds (H2O)

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Valences of Nitrogen
⚫ Nitrogen has five valence electrons (2s2 2p3) but
forms only three bonds (NH3)

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Non-bonding electrons
⚫ Valence electrons not used in bonding are called
nonbonding electrons, or lone-pair electrons
⚫ Nitrogen atom in ammonia (NH3)
⚫ Shares six valence electrons in three covalent bonds
and remaining two valence electrons are nonbonding
lone pair

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Forms of Chemical Bonds
⚫ There are 2 extreme forms of
connecting or bonding atoms:
⚫ Ionic—complete transfer of
electrons from one atom to
another
⚫ Covalent—electrons
shared between atoms
⚫ Most bonds are somewhere
in between.

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Ionic Bonding
⚫ Ionic bonding: results from the electrostatic
attraction between cations and anions.
⚫ Formation of an ionic bond can be viewed as
a transfer of electrons.

Na + Cl Na + Cl
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Bond and Lone Pairs
⚫ Electrons are distributed as shared or BOND
PAIRS and unshared or LONE PAIRS.

••
•
H Cl
••
•

lone pair (LP)

shared or
bond pair
This is called a LEWIS
ELECTRON DOT structure.
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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
⚫ H–H bond results from the overlap of
two singly occupied hydrogen 1s
orbitals
⚫ H-H bond is cylindrically symmetrical,
sigma () bond

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Bond Energy
⚫ Reaction 2 H· → H2 releases 436 kJ/mol
⚫ Product has 436 kJ/mol less energy than two atoms:
H–H has bond strength of 436 kJ/mol. (1 kJ =
0.2390 kcal; 1 kcal = 4.184 kJ)

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Bond Length
⚫ 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

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Hybridization
⚫ The Linear Combination of Atomic Orbitals
(LCAO) to form a new set of atomic orbitals
with the same total electron capacity and with
properties and energies intermediate
between the original unhybridized orbitals.

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Hybridization: sp3 Orbitals and the
Structure of Methane
⚫ Carbon has 4 valence electrons (2s2 2p2)
⚫ In CH4, all C–H bonds are identical (tetrahedral)
⚫ sp3 hybrid orbitals: s orbital and three p orbitals
combine to form four equivalent, unsymmetrical,
tetrahedral orbitals (sppp = sp3), Pauling (1931)

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Tetrahedral Structure of
Methane
⚫ sp3 orbitals on C overlap with 1s orbitals on 4 H
atom 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.

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Hybridization: sp3 Orbitals and the
Structure of Ethane
⚫ Two C’s bond to each other by  overlap of an sp3 orbital from each
⚫ Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H
bonds
⚫ 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

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Hybridization: sp2 Orbitals and the
Structure of Ethylene
⚫ sp2 hybrid orbitals: 2s orbital combines with two 2p
orbitals, giving 3 orbitals (spp = sp 2)
⚫ sp2 orbitals are in a plane with 120° angles
⚫ Remaining p orbital is perpendicular to the plane

120
90

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Bonds From sp2 Hybrid
Orbitals
⚫ Two sp2-hybridized orbitals overlap to form a  bond
⚫ p orbitals overlap side-to-side to formation a pi ()
bond
⚫ sp2–sp2  bond and 2p–2p  bond result in sharing
four electrons and formation of C-C double bond
⚫ Electrons in the  bond are centered between nuclei
⚫ Electrons in the  bond occupy regions are on either
side of a line between nuclei

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Structure of Ethylene
⚫ H atoms form  bonds with four sp2 orbitals
⚫ H–C–H and H–C–C bond angles of about 120°
⚫ C–C double bond in ethylene shorter and stronger
than single bond in ethane
⚫ Ethylene C=C bond length 133 pm (C–C 154 pm)

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-Bonding in C2H4
⚫ The unused p-orbital on each C atom
contains an electron and this p-orbital
overlaps the p-orbital on the neighboring
atom to form the  bond.

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Sigma Bonds in C2H4

H H
2
120 C C sp
H H

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Multiple Bonding in C2H4

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Hybridization: sp Orbitals and the
Structure of Acetylene
⚫ C-C a triple bond sharing six electrons
⚫ Carbon 2s orbital hybridizes with a single p orbital
giving two sp hybrids
⚫ two p orbitals remain unchanged
⚫ sp orbitals are linear, 180° apart on x-axis
⚫ Two p orbitals are perpendicular on the y-axis and
the z-axis

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Orbitals of Acetylene
⚫ Two sp hybrid orbitals from each C form sp–sp 
bond
⚫ pz orbitals from each C form a pz–pz  bond by
sideways overlap and py orbitals overlap similarly

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Bonding in Acetylene
⚫ Sharing of six electrons forms CC
⚫ Two sp orbitals form  bonds with hydrogens

44
Hybridization of Nitrogen and
Oxygen
⚫ Elements other than C can
have hybridized orbitals
⚫ H–N–H bond angle in
ammonia (NH3) 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 45
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°

46
Molecular Orbital Theory
⚫ Robert Mullikan (1896-
1986)
⚫ valence electrons are
DELOCALIZED
⚫ valence electrons are in
orbitals (called
molecular orbitals)
spread over entire
molecule

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Molecular Orbital Theory
Molecular orbitals …
⚫ describe regions of space in a molecule
⚫ Specific shape, size and energy
⚫ are formed by combining atomic orbitals
⚫ lower in energy than the starting atomic
orbitals are bonding
⚫ higher in energy are nonbonding

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Molecular Orbital Theory
⚫ Terminology
⚫ ground state = lowest energy
⚫ excited state = NOT lowest energy
⚫  = sigma bonding MO
⚫ * = sigma antibonding MO
⚫  = pi bonding MO
⚫ * = pi antibonding MO

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Molecular Orbital Theory
⚫ A molecular orbital (MO): where electrons are most likely to
be found (specific energy and general shape) in a molecule
⚫ Additive combination (bonding) MO is lower in energy
⚫ Subtractive combination (antibonding) forms MO is higher

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Molecular Orbital Theory
⚫ MO Theory pictures the formation of a
covalent bond as a mathematical
combination of atomic orbitals.

Antibonding MO

Bonding MO

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Molecular Orbitals
destructive and constructive interference of
atomic orbitals

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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

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Benzene - C6H6

sp2 hybridize the C atoms and combine the

unused p-orbitals into molecular orbitals.
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Benzene - C6H6

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Benzene - C6H6

56
Benzene - C6H6

57
Benzene - C6H6

58
Kekule Structures and
Molecular Formulas
C6H6O C9H8O4

H H O
H C OH H C O C CH3
C C C C
C C C C
H C H H C COOH
H H
O
O
OH
OH
59

O
 H3C CH2
CH3CH2CH2CH2CHCH2CH2 CH2 CH3

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Structure Forms
Kekulé or
Wedge/dash

C3H8 H3C CH2 CH3 H H

H H
skeletal
CH3CH2CH3 H H H
condensed H

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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 dumbbell-shaped
⚫ 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
⚫ Sigma () bonds - Circular cross-section and are formed by head-on interaction
⚫ Pi () bonds – “dumbbell” shape from sideways interaction of p orbitals
⚫ 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
⚫ 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

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