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Organic Organic
ChemistryChemistry MS.SUPAWADEE SRITHAHANDEPARTMENT OF CHEMISTRY
MAHIDOL WITTAYANUSORN SCHOOL
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CONTENTS
INTRODUCTIONCLASSIFICATION NAMING AND PROPERTIES OF ORGANIC COMPOUNDBONDING OF ORGANIC COMPOUNDALKANE & CYCLOALKANEALKENE & CYCLOALKENEALKYNE & CYCLOALKYNE
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Organic ChemistryOrganic 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)
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ShellsShells
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
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Atomic OrbitalsAtomic OrbitalsElectrons surrounding atoms are
concentrated into regions of space called atomic atomic
orbitalsorbitals.. Four different kinds of orbitals ; ss, , pp, , dd, and , and ff 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-Orbitalsp-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 ConfigurationsElectron 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 Atoms1S2
2S2
3S2
4S2
5S2
6S2
7S2
2p6
3p6
4p6
5p6
6p6
7p6
3d10
5d10
6d10
7d10
4d10 4f14
6f14
5f14
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Valences of CarbonValences of Carbon
Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4)
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Valences of NitrogenValences of Nitrogen
Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3)
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Non-bonding electronsNon-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 TheoryValence 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 EnergyBond 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
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D(C-H) = (1660/4) kJ/mol = 415 kJ/mol
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Bond LengthBond 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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sp3 Hybridization of Carbon
Ground state Excited state sp3-hybridization state
HybridizationPromotion of electron
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Hybridization: Hybridization: spsp3 3 OrbitalsOrbitals
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 MethaneTetrahedral 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 EthaneThe 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 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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spsp22 Hybridization Hybridization of Carbonof Carbon
Ground state Excited state sp2-hybridization state
HybridizationPromotion of electron
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Hybridization: Hybridization: spsp2 2 OrbitalsOrbitals
sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2)
sp2 orbitals are in a plane with120° angles; trigonal trigonal planarplanar
Remaining p orbital is perpendicular to the plane
90120
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Bonds From spBonds From sp22 Hybrid Orbitals 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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Bonding in 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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Hybridization: Hybridization: spsp OrbitalsOrbitals
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 AcetyleneOrbitals 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 AcetyleneBonding in Acetylene
Sharing of six electrons forms C C Two sp orbitals form bonds with hydrogens
Polarity Polarity
Polarity refers to a separation of positive and negative charge.
In a nonpolar bond, the bonding electrons are shared equally.
HCl:
In a polar bond, electrons are shared unequally.
H2,Cl2:
ElectronegativityElectronegativity
Electronegativity refers to the ability of an atom in a molecule to attract shared electrons.
The Pauling scale of electronegativity:
QuickTime Movie
Bond PolarityBond Polarity
A polar bond can be pictured using partial charges:
= 0.9
ElectronegativityDifference Bond Type
0 - 0.5 Nonpolar
0.5 - 2.0 Polar
2.0 Ionic
2.1 3.0
+
H Cl
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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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Structural of organic compoundsStructural of organic compounds
1. Dot structure 2. Dash Formula
3. Condensed formula 4. Partial Condensed Formula
CHCH33CHCH22CHCH22COOHCOOH
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Structural of organic compoundsStructural of organic compounds
5. Line-angle formula or bond line formula
CHCH33CHCH22CHCH22COOHCOOH
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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
HH
H
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Sample Problem
Rewrite each of the following condensed structural formulas, as dash formulas as :
C
H
H
H
C
H
C
C
H
H
C
C
H
HH
HH
C
H
H
H
H
H
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