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
7
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
8
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
10
Write electron configurations of Carbon atom
2p
2s
1s
12
6C……………………………………………….......
11
Molecular Orbitals
Covalent bond
Electrostatic Interactions
12
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)
14
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
15
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
16
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)
17
Bond energy
พลังงานพันธะ คือ พลังงานที่ใช้ในการสลายพันธะระหว่างอะตอมของธาตุ
ภายในโมเลกุลที่อยูใ่ นสถานะก๊าซออกเป็ นอะตอมเดี่ยว ๆ
เช่น
พลังงานพันธะเฉลี่ย คือ พลังงานเฉลี่ยที่ใช้สลายพันธะแต่ละพันธะในคู่
อะตอมเดียวกัน
D(C-H) = (1660/4) kJ/mol = 415 kJ/mol
18
Bond energy
19
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
20
Bond Lengths
21
Bond Lengths
22
Bond length and Bond strength
23
electron configurations of Carbon atom
12
6C
2p
2s
1s
Not CH2
CH4
Why?
Hybridization
24
sp3 Hybridization of Carbon
Ground state
Excited state
Promotion of electron
sp3-hybridization state
Hybridization
25
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
28
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°
30
sp2 Hybridization of Carbon
Ground state
Excited state
Promotion of electron
sp2-hybridization state
Hybridization
31
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
32
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
34
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
38
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
44
Molecule Polarity
45
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
47
Stearic replusion
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Intramolecular Hydrogen Bond
49
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
55
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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