Transcript Orbitals

Organic Chemistry
MS.SUPAWADEE SRITHAHAN
DEPARTMENT OF CHEMISTRY
MAHIDOL WITTAYANUSORN SCHOOL
CONTENTS
INTRODUCTION
CLASSIFICATION NAMING AND PROPERTIES OF ORGANIC COMPOUND
BONDING OF ORGANIC COMPOUND
ALKANE & CYCLOALKANE
ALKENE & CYCLOALKENE
ALKYNE & CYCLOALKYNE
2
INTRODUCTION
Structure and Bonding
Organic Chemistry
 “Organic” – until mid 1800’s referred to compounds
from living sources (mineral sources were
“inorganic”)
 Wöhler in 1828 showed that urea, an organic
compound, could be made from a minerals
 Today, organic compounds are those based on
carbon structures and organic chemistry studies
their structures and reactions



Includes biological molecules, drugs, solvents, dyes
Does not include metal salts and materials (inorganic)
Does not include materials of large repeating
molecules without sequences (polymers)
4
Atomic Structure
5
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
6
Atomic Orbitals
Electrons surrounding atoms are concentrated into
regions of space called atomic orbitals.




Four different kinds of orbitals ; 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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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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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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Electronic Configurations of Atoms
1S2
2S2
2p6
3S2
3p6
3d10
4S2
4p6
4d10
4f14
5S2
5p6
5d10
5f14
6S2
6p6
6d10
6f14
7S2
7p6
7d10
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Write electron configurations of Carbon atom
2p
2s
1s
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6C……………………………………………….......
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Molecular Orbitals
 Covalent bond
Electrostatic Interactions
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Valences of Carbon
 Carbon has four valence electrons (2s2 2p2), forming
four bonds (CH4)
13
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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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 (s) 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 energy
 พลังงานพันธะ คือ พลังงานที่ใช้ในการสลายพันธะระหว่างอะตอมของธาตุ
ภายในโมเลกุลที่อยูใ่ นสถานะก๊าซออกเป็ นอะตอมเดี่ยว ๆ
 เช่น
 พลังงานพันธะเฉลี่ย คือ พลังงานเฉลี่ยที่ใช้สลายพันธะแต่ละพันธะในคู่
อะตอมเดียวกัน
D(C-H) = (1660/4) kJ/mol = 415 kJ/mol
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Bond energy
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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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Bond Lengths
21
Bond Lengths
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Bond length and Bond strength
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electron configurations of Carbon atom
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6C
2p
2s
1s
Not CH2
CH4
Why?
Hybridization
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sp3 Hybridization of Carbon
Ground state
Excited state
Promotion of electron
sp3-hybridization state
Hybridization
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Hybridization: sp3 Orbitals
 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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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 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 of Nitrogen
 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
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Hybridization of Oxygen
 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°
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sp2 Hybridization of Carbon
Ground state
Excited state
Promotion of electron
sp2-hybridization state
Hybridization
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Hybridization: sp2 Orbitals
 sp2 hybrid orbitals: 2s orbital combines with two 2p
orbitals, giving 3 orbitals (spp = sp2)
 sp2 orbitals are in a plane with120° angles; trigonal
planar
 Remaining p orbital is perpendicular to the plane
90
120
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Bonds From sp2 Hybrid Orbitals
 Two sp2-hybridized orbitals overlap to form a s bond
 p orbitals overlap side-to-side to formation a pi () bond
 sp2–sp2 s bond and 2p–2p  bond result in sharing four
electrons and formation of C-C double bond
 Electrons in the s bond are centered between nuclei
 Electrons in the  bond occupy regions are on either side of a
line between nuclei
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The Orbital of Ethene
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Bonding in Ethylene
 H atoms form s 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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Hybridization: sp Orbitals
 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 s
bond
 pz orbitals from each C form a pz–pz  bond by
sideways overlap and py orbitals overlap similarly
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Orbitals of Acetylene
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Bonding in Acetylene
 Sharing of six electrons forms C C
 Two sp orbitals form s bonds with hydrogens
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Bond Polarity
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Polarity
 Polarity refers to a separation of positive and negative
charge.
 In a nonpolar bond, the bonding electrons are shared
equally.
H2,
Cl2:
 In a polar bond, electrons are shared unequally.
HCl:
Electronegativity
 Electronegativity refers to the ability of an atom in
a molecule to attract shared electrons.
 The Pauling scale of electronegativity:
Bond Polarity
A polar bond can be pictured using partial
charges: +

H
2.1
Cl
 = 0.9
3.0
Electronegativity
Difference
Bond Type
0 - 0.5
Nonpolar
0.5 - 2.0
Polar
2.0 
Ionic
Bond Polarity
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Molecule Polarity
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Types of Interactions
 1. Intramolecular force





Covalent bond
Ionic bond
Metallic bond
Stearic replusion
Intramolecular Hydrogen Bond
 2. Intermolecular force


Van de Waals force
Hydrogen bond
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Intramolecular forces
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Stearic replusion
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Intramolecular Hydrogen Bond
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Intermolecular forces
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Ion – Dipole Forces
* Between a charged ion and polar molecule
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Dipole-Dipole forces
* Between neutral polar molecule
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London dispersion forces
* Between non polar/non polar molecules
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Hydrogen bonding
* Between hydrogen and an electronegative
atom such as F, O or N
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Structural of organic compounds
CH3CH2CH2COOH
1. Dot structure
2. Dash Formula
3. Condensed formula
4. Partial Condensed Formula
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Structural of organic compounds
CH3CH2CH2COOH
5. Line-angle formula or bond line formula
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6. Three-dimensional formulas
Bonds that project upward out of the plane of the paper
Bonds that lie behind the plane
Bonds that lie in the plane of the page
H
C
H
H
H
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Sample Problem
Rewrite each of the following condensed structural formulas, as dash
formulas as :
H
H
H
H
C
C
H
H
C
C
C
H H C H H
H
H
H
C
H
H
H
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Write dash formulas for each of the following bond-line formulas:
D.
A.
OH
B.
OH
E.
OH
C.
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