1-1 1 copyright © 2000 by john wiley & sons, inc. all rights reserved. introduction to organic...

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Copyright © 2000 by John Wiley & Sons, Inc. All rights reserved.

Introduction to Introduction to Organic Organic

ChemistryChemistry2 ed2 ed

William H. BrownWilliam H. Brown

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Covalent Covalent Bonding & Bonding & Shapes of Shapes of MoleculesMolecules

Chapter 1 Chapter 1

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Organic ChemistryOrganic Chemistry• The study of the compounds of carbon• Over 10 million organic compounds have been

identified• about 1000 new ones are discovered or synthesized

and identified each day!

• C is a small atom • it forms single, double, and triple bonds• it is intermediate in electronegativity (2.5)• it forms strong bonds with C, H, O, N, and some metals

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Electronic Structure of AtomsElectronic Structure of Atoms• Structure of atoms

• small dense nucleus, diameter 10-14 - 10-15 m, which contains most of the mass of the atom

• extranuclear space, diameter 10-10 m, which contains positively-charged electrons

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Electronic Structure of AtomsElectronic Structure of Atoms• Electrons are confined to regions of space

called principle energy levels (shells)• each shell can hold 2n2 electrons (n = 1,2,3,4......)

Shell

Number of Electrons ShellCan Hold

Relative Energiesof Electrons in These Shells

3218 8 2

4321

higher

lower

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Electronic Structure of AtomsElectronic Structure of Atoms• Shells are divided into subshells called orbitals, which are

designated by the letters s, p, d,........• s (one per shell) • p (set of three per shell 2 and higher) • d (set of five per shell 3 and higher) .....

Shell Orbitals Contained in That Shell

3

2

1 1s

2s, 2px, 2py, 2pz

3s, 3px, 3py, 3pz, plus five 3d orbitals

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Shapes of Atomic OrbitalsShapes of Atomic Orbitals• All s orbitals have the

shape of a sphere, with its center at the nucleus. Of the s orbitals, a 1s is the smallest, a 2s is larger, and a 3s is larger still

• A p orbital consists of two lobes arranged in a straight line with the center at the nucleus

x

y

z

a pz orbital

x

y

z

an s orbital

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Electron Structure of AtomsElectron Structure of Atoms

Rule 1Rule 1: Orbitals fill in order of increasing energy, from lowest to highest

Rule 2Rule 2: Each orbital can hold up to two electrons

Rule 3Rule 3: When orbitals of equivalent energy are available but there are not enough electrons to fill them completely, one electron is added to each equivalent orbital before a second is added to any one of them

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Lewis StructuresLewis Structures• Gilbert N. Lewis• Valence shellValence shell: the outermost electron shell of an

atom• Valence electronsValence electrons: electrons in the valence shell

of an atom; these electrons are used in forming chemical bonds

• Lewis structure Lewis structure • the symbol of the atom represents the nucleus and all

inner shell electrons• dots represent valence electrons

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Lewis StructuresLewis Structures• Lewis structures for elements 1-18 of the periodic table

N

•

•OB

H

Li Be

Na Mg

He

IA IIA VA VIA VIIA VIIIAIIIA IVA

Cl

F

S

Ne

Ar

C

SiAl P

••

••

••

••••

••

•

•

••

•

••

•• •

•

•

•

•

•

•

•

•

•

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Lewis Model of BondingLewis Model of Bonding• Atoms bond together so that each atom in the

bond acquires the electron configuration of the noble gas closest it in atomic number• an atom that gains electrons becomes an anion• an atom that loses electrons becomes a cation

• Ionic bondIonic bond: a chemical bond resulting from the electrostatic attraction of an anion and a cation

• Covalent bondCovalent bond: a chemical bond resulting from two atoms sharing one or more pairs of electrons

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ElectronegativityElectronegativity• ElectronegativityElectronegativity: a measure of the force of an atom’s attraction

for the electrons it shares in a chemical bond with another atom• Pauling scale

• increases from left to right within a period• increases from bottom to top in a group

H Cl2.1 3.0

δ+ δ-

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ElectronegativityElectronegativityClassification of chemical bonds

Difference in ElectronegativityBetween Bonded AtomsType of Bond

less than 0.5

0.5 to 1.9greater than 1.9

nonpolar covalentpolar covalent

ionic

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Lewis StructuresLewis Structures• To write a Lewis structure

• determine the number of valence electrons in the molecule or ion

• determine the arrangement of atoms• connect the atoms by single bonds• arrange the remaining electrons so that each atom has

a complete valence shell• show bonding electrons as a single bond• show nonbonding electrons as a pair of dots• atoms share 1 pair of electrons in a single bond, 2

pairs in a double bond, and 3 pairs in a triple bond

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Lewis StructuresLewis Structures

Hydrogen chloride

Methane Ammonia

Water

H O

H

H

H NH C

H

H

H Cl

H

H

H2O (8)

NH3 (8)CH4 (8)

HCl (8)

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Lewis StructuresLewis Structures

Carbonic acidMethanal

Acetylene

Ethylene

H

C C

H

C C HH

HO

CC

HO

H

O

H

HH

O

C2H4 (12)

C2H2 (10)

CH2O (12) H2CO3 (24)

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Lewis StructuresLewis Structures• In neutral molecules containing C, H, N, O, and

halogen• hydrogen has one bond• carbon has 4 bonds and no unshared electrons• nitrogen has 3 bonds and 1 unshared pair of electrons• oxygen has 2 bonds and 2 unshared pairs of electrons• halogen has 1 bond and 3 unshared pairs of electrons

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Formal ChargeFormal Charge• Formal chargeFormal charge: the charge on an atom in a

molecule or polyatomic ion• write a Lewis structure for the molecule or ion• assign each atom all of its unshared (nonbonding)

electrons and one-half its shared (bonding) electrons• compare this number with the number of valence

electrons in the neutral, unbonded atom

# of valence electrons in

unbonded atom

all unsharedelectrons

one half of all shared

electrons+Formal

charge=

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Formal ChargeFormal Charge• if the number assigned to the bonded atom is less than that assigned to the unbonded atom, the atom has a positive formal charge• if the number is greater, the atom has a negative formal charge # of valence

electrons in unbonded atom

all unsharedelectrons

one half of all shared

electrons+Formal

charge=

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Formal ChargeFormal Charge• Example: Example: draw Lewis structures and show all formal charges for these ions

NH4+ CH3NH3

+

HCO3- CO3

2-OH-

CH3CO2-CH3

-

CH3OH2+

BF4-

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VSEPR ModelVSEPR Model

4 regions of e- density(tetrahedral, 109.5°)

••••

••••

••

••

••

••

HH

NH

C

HH H

H

H

C C

H

O C

H

C

HH

O

H

CH C HO

3 regions of e- density(trigonal planar, 120°)

2 regions of e- density(linear, 180°)

••

HH

O

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VSEPR ModelVSEPR ModelExample: Example: predict all bond angles for these molecules and ionsNH4

+ CH3NH2

H2CO3 HCO3-CH3CH=CH2

CH3CO2HCH3CHO

CH3OH

BF4-

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Polar & Nonpolar MoleculesPolar & Nonpolar Molecules• A molecule is polar if

• it has polar bonds and• its center of positive charge is at a different place

than its center of negative charge

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Polar & Nonpolar MoleculesPolar & Nonpolar Molecules

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ResonanceResonance• For many molecules and ions, no single Lewis

structure provides a truly accurate representation

Ethanoate ion(Acetate ion)

C

O

O

CH3C

O

O

CH3

-

-

and

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ResonanceResonance• Linus Pauling - 1930s

• many molecules and ions are best described by writing two or more Lewis structures

• individual Lewis structures are called contributing structures

• connect individual contributing structures by a double-headed (resonance) arrow

• the molecule or ion is a hybrid of the various contributing structures

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ResonanceResonanceExamplesExamples:

N

O

O

N

O

O

-

-

Ethanoate ion(equivalent contributing structures)

C

O

O

CH3 C

O

O

CH3

-

Nitrite ion(equivalent contributing structures)

-

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ResonanceResonance• Curved arrowCurved arrow: a symbol used to show the

redistribution of valence electrons• In using curved arrows, there are only two

allowed types of electron redistribution:• from a bond to an adjacent atom• from an atom to an adjacent bond

• Electron pushing by the use of curved arrows is a survival skill in organic chemistry. Learn it well!

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ResonanceResonance• All acceptable contributing structures must1. have the same number of valence electrons

2. obey the rules of covalent bonding• no more than 2 electrons in the valence shell of H • no more than 8 electrons in valence shell of 2nd

period elements• 3rd period elements may have up to 12 electrons

in their valence shells

3. differ only in distribution of valence electrons

4. have the same number of paired and unpaired electrons

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ResonanceResonance• Examples of ions and a molecule best

represented as resonance hybrids

carbonate ion CO32-

acetate ion

CH3COCH3acetone

CH3CO2-

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Valence Bond ModelValence Bond Model• a covalent bond is formed

when a portion of an atomic orbital of one atom overlaps a portion of an atomic orbital of another atom.

• in forming the covalent bond in H-H, for example, there is overlap of the 1s orbitals of each hydrogen

H H

area of overlap(the covalent bond)

• •

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Hybrid OrbitalsHybrid Orbitals• The Problem:

• 2s and 2p atomic orbitals would give bond angles of approximately 90°

• instead we observe approximately 109.5°,120°, and 180°

• A Solution• hybridization of atomic orbitals• 2nd row elements use sp3, sp2, and sp hybrid

orbitals for bonding

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Hybrid OrbitalsHybrid Orbitals• We study three types of hybrid atomic orbitals

spsp33 (1 s orbital + 3 p orbitals -> four sp3 orbitals)

spsp22 (1 s orbital + 2 p orbitals -> three sp2 orbitals)

spsp (1 s orbital + 1 p orbital -> two orbitals)

• Overlap of hybrid atomic orbitals can form two types of bonds, depending on the geometry of the overlap bonds bonds are formed by “direct” overlap

bonds bonds are formed by “parallel” overlap

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spsp33 Hybrid Orbitals Hybrid Orbitals• each sp3 hybrid orbital has

two lobes of unequal size• the four sp3 hybrid orbitals are

directed toward the corners of a regular tetrahedron at angles of 109.5°

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spsp22 Hybrid Orbitals Hybrid Orbitals• each sp2 hybrid orbital has

two lobes of unequal size • the three sp2 hybrid

orbitals are directed toward the corners of an equilateral triangle at angles of 120°

• the unhybridized 2p orbital is perpendicular to the plane of the sp2 hybrid orbitals

sp2 hybrid

orbitals

unhybridized2p orbital

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sp Hybrid Orbitalssp Hybrid Orbitals• each sp hybrid orbital

has two lobes of unequal size

• the two sp hybrid orbitals lie in a line at an angle of 180°

• the two unhybridized 2p orbitals are perpendicular to each other and to the line through the two sp hybrid orbitals

x

y

z

sp hybridorbitals

unhybridized 2p orbitals lie on the y and z axes

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Hybrid OrbitalsHybrid OrbitalsHybrid-ization

Types of Bonds to CarbonExample

sp3 four sigma bonds

sp2 three sigma bondsand one pi bond

sp two sigma bondsand two pi bonds

H H

H H

H-C C-H

H

C C

H

H H

H-C-C-H Ethane

Ethylene

Acetylene

Name

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Functional GroupsFunctional Groups• Functional groupFunctional group: an atom or group of atoms

within a molecule that shows a characteristic set of physical and chemical properties

• Functional groups are important for three reason; they are• the units by which we divide organic compounds into

classes• the sites of characteristic chemical reactions• the basis for naming organic compounds

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Functional GroupsFunctional Groups• Hydroxyl (OH) group of alcohols

An alcohol (Ethanol)

H-C-C-O-H

HH

H H

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Functional GroupsFunctional Groups• Carbonyl group of aldehydes and ketones

An aldehyde A ketone

O O

CH3-C-H CH3-C-CH3

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Functional GroupsFunctional Groups• Carboxyl group of carboxylic acids

or or

O

CH3-C-O-H CH3COOH CH3CO2H

Acetic acid

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AminesAmines• Amino group of 1°, 2° and 3° amines

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Covalent Bonding Covalent Bonding & Shapes of & Shapes of MoleculesMolecules

End Chapter 1End Chapter 1

1-1 1 copyright © 2000 by john wiley & sons, inc. all rights reserved. introduction to organic...

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