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

INTRODUCTIONINTRODUCTION

Structure and BondingStructure and Bonding

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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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Atomic StructureAtomic Structure

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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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Write electron configurations of Carbon atom

126C……………………………………………….......

1s

2s

2p

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Molecular OrbitalsMolecular Orbitals

Covalent bond

Electrostatic Interactions

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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 energy

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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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Bond Lengths

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Bond Lengths

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Bond length and Bond strength

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electron configurations of Carbon atom

126C

Not CH2 CH4

1s

2s

2p

Why?Why?Hybridization

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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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The Orbital of EtheneThe Orbital of Ethene

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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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Orbitals of AcetyleneOrbitals of Acetylene

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Bonding in AcetyleneBonding in Acetylene

Sharing of six electrons forms C C Two sp orbitals form bonds with hydrogens

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Bond Polarity

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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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 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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Write dash formulas for each of the following bond-line formulas:

OH

OH

OH

A.

B.

C.

D.

E.