organic chemistry i chm 201
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Organic Chemistry I CHM 201. William A. Price, Ph.D. Introduction and Review: Structure and Bonding. Atomic structure Lewis Structures Resonance Structural Formulas Acids and Bases. Electronic Structure of the Atom. - PowerPoint PPT PresentationTRANSCRIPT
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Organic Chemistry ICHM 201
William A. Price, Ph.D.
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Introduction and Review: Structure and Bonding
Atomic structureLewis Structures
ResonanceStructural Formulas
Acids and Bases
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Chapter 1 5
Electronic Structure of the Atom
• An atom has a dense, positively charged nucleus surrounded by a cloud of electrons.
• The electron density is highest at the nucleus and drops off exponentially with increasing distance from the nucleus in any direction.
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Orbitals are Probabilities
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2s Orbital Has a Node
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The p Orbital
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The 2p Orbitals• There are three 2p
orbitals, oriented at right angles to each other.
• Each p orbital consists of two lobes.
• Each is labeled according to its orientation along the x, y, or z axis.
Chapter 1 9
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px, py, pz
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Electronic Configurations
• The aufbau principle states to fill the lowest energy orbitals first.
• Hund’s rule states that when there are two or more orbitals of the same energy (degenerate), electrons will go into different orbitals rather than pairing up in the same orbital.
Chapter 1 11
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Electronic Configurations of Atoms• Valence electrons are electrons on the
outermost shell of the atom.
Chapter 1 12
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Covalent Bonding• Electrons are shared between the atoms to complete
the octet.• When the electrons are shared evenly, the bond is
said to be nonpolar covalent, or pure covalent.• When electrons are not shared evenly between the
atoms, the resulting bond will be polar covalent.
Chapter 1 13
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Lewis Dot Structure of Methane
C.. . .
carbon - 4 valence e hydrogen - 1 valence e
H.
1s22s22p2 1s
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Tetrahderal Geometry
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CH4 NH3
H2O Cl2
Lewis Structures
C
H
H
H
HN H
H
H
O HH ClCl
Carbon: 4 e4 H@1 e ea: 4 e 8 e
Nitrogen: 5 e3 H@1 e ea: 3 e 8 e
Oxygen: 6 e2 H@1 e ea: 2 e 8 e
2 Cl @7 e ea: 14 e
Chapter 1 16
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Bonding PatternsValence
electrons (group #)
# Bonds # Lone Pair Electrons
C
N
O
Halides(F, Cl, Br, I)
4 4 0
5 3 1
6 2 2
7 1 3
Chapter 1 17
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Bonding Characteristics of Period 2 Elements
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Lewis structures are the way wewrite organic chemistry.
Learning now to draw themquickly and correctly will helpyou throughout this course.
HINT
Chapter 1 20
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Multiple Bonding
• Sharing two pairs of electrons is called a double bond.• Sharing three pairs of electrons is called a triple bond.
Chapter 1 21
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Convert Formula into Lewis Structure
• HCN• HNO2
• CHOCl• C2H3Cl
• N2H2
• O3
• HCO3-
• C3H4
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Formal Charges
H3O+ NO+
OH H
H
Formal charge = [group number ] – [nonbonding electrons ] – ½ [shared electrons]
N O
6 – 2 – ½ (6) = +1 6 – 2 – ½ (6) = +1
5 – 2 – ½ (6) = 0+
+
• Formal charges are a way of keeping track of electrons.• They may or may not correspond to actual charges in the
molecule.
Chapter 1 23
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Common Bonding Patterns
Chapter 1 24
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Work enough problems tobecome familiar with thesebonding patterns so you canrecognize other patterns as
being either unusual or wrong.
HINT
Chapter 1 25
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Electronegativity Trends Ability to Attract the Electrons in a Covalent Bond
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Dipole Moment
• Dipole moment is defined to be the amount of charge separation (d) multiplied by the bond length (m).
• Charge separation is shown by an electrostatic potential map (EPM), where red indicates a partially negative region and blue indicates a partially positive region.
Chapter 1 27
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Methanol
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Dipole Moment (m) is sum of the Bond Moments
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Nonpolar CompoundsBond Moments Cancel Out
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Nitromethane
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Nitromethane has 2 Formal Charges
CH3NO2 C N
O
O
H
HH
Formal Charge = [Group #] - [# bonds] - [# non-bonding electrons]
N = 5-4-0 = +1
O = 6-1-6 = -1
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Both Resonance Structures Contribute to the Actual Structure
CH3NO2
C N
O
O
H
HHH
H
H
N
O
O
C
2 Equivalent Resonance Structures
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Dipole Moment reflects Both Resonance Structures
C N
O
O
H
HHH
H
H
N
O
O
C
Resonance Hybrid
C
H
HH
O
O
N
d
d
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Resonance Rules• Cannot break single (sigma) bonds• Only electrons move, not atoms
3 possibilities:– Lone pair of e- to adjacent bond position
• Forms p bond
p bond to adjacent atom p bond to adjacent bond position
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Curved Arrow Formalism Shows flow of electrons
C N
O
O
H
HHH
H
H
N
O
O
C
Arrows depict electron pairs moving
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Resonance Forms• The structures of some compounds are not
adequately represented by a single Lewis structure.• Resonance forms are Lewis structures that can be
interconverted by moving electrons only. • The true structure will be a hybrid between the
contributing resonance forms.
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Resonance Forms
Resonance forms can be compared using the following criteria, beginning with the most important:1. Has as many octets as possible.2. Has as many bonds as possible.3. Has the negative charge on the most
electronegative atom.4. Has as little charge separation as possible.
Chapter 1 38
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Two Nonequivalent Resonance Structures in Formaldehyde
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Major and Minor Contributors
• When both resonance forms obey the octet rule, the major contributor is the one with the negative charge on the most electronegative atom.
N C O N C OMAJOR MINOR
The oxygen is more electronegative,so it should have more of the negativecharge.
Chapter 1 40
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Resonance Stabilization of IonsPos. charge is “delocalized”
CH
HC
C
H
H
H
CC
H
H
H
H
HC
H
HCd
d
H
H
H
CC
resonance hybrid
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Solved Problem 2
Draw the important resonance forms for [CH3OCH2]+. Indicate which structure is the major and minor contributor or whether they would have the same energy.
The first (minor) structure has a carbon atom with only six electrons around it. The second (major) structure has octets on all atoms and an additional bond.
Solution
Chapter 1 42
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Solved Problem 3Draw the resonance structures of the compound below. Indicate which structure is the major and minor contributor or whether they would have the same energy.
Both of these structures have octets on oxygen and both carbon atoms, and they have the same number of bonds. The first structure has the negative charge on carbon, the second on oxygen. Oxygen is the more electronegative element, so the second structure is the major contributor.
Solution
Chapter 1 43
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Resonance Forms for the Acetate Ion
• When acetic acid loses a proton, the resulting acetate ion has a negative charge delocalized over both oxygen atoms.
• Each oxygen atom bears half of the negative charge, and this delocalization stabilizes the ion.
• Each of the carbon–oxygen bonds is halfway between a single bond and a double bond and is said to have a bond order of 1½.
Chapter 1 44
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Condensed Structural Formulas Lewis Condensed
CH3CH3H C C H
H
H
H
H
1 2
• Condensed forms are written without showing all the individual bonds.
• Atoms bonded to the central atom are listed after the central atom (CH3CH3, not H3CCH3).
• If there are two or more identical groups, parentheses and a subscript may be used to represent them.
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Drawing Structures
CC
CC
H
H
H
H
HH H
HH
H
Butane, C4H10
CC
C
C
H
H H
HH
H H
HHH
Methylpropane, C4H10
CH3CH2CH2CH3 CH3CH(CH3)CH3
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Octane Representations
CC
CC
CC
CC
H
HH
H H
H H
H H
H H
H H
HHH
HH
CH3CH2CH2CH2CH2CH2CH2CH3
CH3(CH2)6CH3
Lewis structure condensed structural formula
C8H18 is molecular formula but there are 18 possible structural isomers
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Line-Angle Structures are Often Used as a Short-hand
Line-angle or "zig-zag" structures
Lewis structure
CC
CC
CC
CC
H
HH
H H
H H
H H
H H
H H
HHH
HH
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Line-Angle Structures
CC
CC
CC
C
CC
C
CC
HH
H
H
H
H
H HH
H H
H
H
H
H
H
HH
HH
H
H H
HH
H
HH
H
H
H
H
H HH
H H
H
H
H
H
H
HH
HH
H
H H
HH
H
C12H26
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HH
H
H
H
H
H HH
H H
H
H
H
H
H
HH
HH
H
H H
HH
H
At the end of every line and at the intersection of any lines there is a carbon atom with 4 bonds. Hydrogen atoms
are mentally supplied to fill the valency to 4.
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Line-Angle structure Superimposed on Lewis Structure
CC
CC
CC
C
CC
C
CC
HH
H
H
H
H
H HH
H H
H
H
H
H
H
HH
HH
H
H H
HH
H
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Line-Angle Drawings
H C C C
H
H
H
H
C
H
C
H
H
H H
H
C
O
H
O
H
1 2 3 4 5 61 2 3 4 5 6
• Atoms other than carbon must be shown.• Double and triple bonds must also be shown.
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For Cyclic Structures, Draw the Corresponding Polygon
Cyclohexane
CC
CC
C
C
HHH
HH
HHHH
HH
H
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Some Steroids
HOCholesterol
C27H44O
OH
OTestosterone
C19H26O2
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Definitions of Acids/Bases
Arrhenius acid - forms H3O+ in H2O
Bronsted-Lowry acid - donates a H+ (proton)
Lewis acid - accepts an electron pair to form a new bond
Arrhenius base - forms OH- in H2O
Bronsted-Lowry base - accepts a H+ (proton)
Lewis base - donates an electron pair to form a new bond
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Dissociation in H2OArrhenius Acid forms H3O+
Bronsted-Lowry Acid donates a H+
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Brønsted-Lowry Acids and BasesBrønsted-Lowry acids are any species that donate a proton.
Brønsted-Lowry bases are any species that can accept a proton.
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Conjugate Acids and Bases
• Conjugate acid: when a base accepts a proton, it becomes an acid capable of returning that proton.
• Conjugate base: when an acid donates its proton, it becomes capable of accepting that proton back.
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Acid Strength defined by pKa
HCl + H2O H3O + Cl
Keq = [H3O ][Cl ][HCl][H2O]
Ka = Keq[H2O] =[H3O ][Cl ]
[HCl]= 107
pKa = -log(Ka) = -7
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Stronger Acid Controls Equilibrium
HCl + H2O H3O + Clacid base conjugate conjugate
acid basepKa = -7 -1.7
stronger weaker
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Reaction Described with Arrows
H Cl + OHH O
H
HH+ Cl
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Equilibrium Reactions
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Identify the Acid and Base
CH3OH + OCOH
O
CH3O + HOCOH
O
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Equilibrium Favors Reactants
CH3OH + OCOH
O
CH3O + HOCOH
O
acid base conj. base conj. acid
pKa 15.5 6.5
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The Effect of Resonance on pKa
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Effect of Electronegativity on pKa
• As the bond to H becomes more polarized, H becomes more positive and the bond is easier to break.
Chapter 1 66
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Effect of Size on pKa
• As size increases, the H is more loosely held and the bond is easier to break.
• A larger size also stabilizes the anion.
Chapter 1 67
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Lewis Acids and Lewis Bases
• Lewis bases are species with available electrons than can be donated to form a new bond.
• Lewis acids are species that can accept these electrons to form new bonds.
• Since a Lewis acid accepts a pair of electrons, it is called an electrophile.
Chapter 1 68
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Nucleophiles and Electrophiles
• Nucleophile: Donates electrons to a nucleus with an empty orbital (same as Lewis Base)
• Electrophile: Accepts a pair of electrons (same as Lewis Acid)
• When forming a bond, the nucleophile attacks the electrophile, so the arrow goes from negative to positive.
• When breaking a bond, the more electronegative atom receives the electrons.
Chapter 1 69
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Nucleophiles and Electrophiles
Chapter 1 70