Transcript Slide 1

Welcome to
Organic Chemistry 234!
How Should I Study?
• Do not memorize everything!
• Practice writing mechanisms and “talking” yourself through the steps.
• Learn to ask the right questions.
• Form a small study group (2-3 people).
• Work as many problems as you can.
• Do not hesitate to visit me during office hours for assistance.
• A free tutoring service is available through the LRC.
What is Organic Chemistry?
• It is the study of carbon-containing compounds
Why Carbon?
• Carbon neither gives up nor accepts electrons because it is in the
center of the second periodic row.
• Consequently, carbon forms bonds with other carbons and other
atoms by sharing electrons.
• The capacity of carbon to form bonds in this fashion makes it the
building block of all living organisms.
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Why Study Organic Chemistry?
• Since
carbon is the building block of all living organisms, a
knowledge of Organic Chemistry is a prerequisite to
understanding Biochemistry, Medicinal Chemistry, and
Pharmacology.
• Indeed, Organic Chemistry is a required course for studying
Pharmacy, Medicine, and Dentistry.
• Admission into these professional programs is highly
dependent on your performance in Organic Chemistry.
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Examples of Organic Compounds Used as Drugs
Methotrexate, Anticancer Drug
AZT, HIV Drug
5-Fluorouracil, Colon
Cancer Drug
Tamiflu, Influenza
Drug
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Examples of Organic Compounds Used as Drugs
Haldol, Antipsychotic
Elavil, Antidepressant
Prozac, Antidepressant
Viagra, Treats
Erectile Dysfunction
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Spring 2011
Dr. Halligan
CHM 234
Chapter 1
• Electronic Structure and Bonding
• Acids and Bases
“Speaking Organic Chemistry”
• What are some of the fundamentals of organic
chemistry that we will cover in Chapter 1?
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The periodic table
Bonding
Lewis structures
Delocalized electrons and Resonance Structures
Orbital Hybridization
The art of drawing structures and comprehending organic
compounds
Trends in electronegativity
Determination of formal charges
The use of molecular models to represent compounds
Acids and Bases
Structure and Bonding
Note: Sections 1.1 and 1.2 on the structure of an atom can be reviewed in the textbook.
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Ionic, Covalent, and Polar Bonds
• Bonds formed between two oppositely charged ions are
considered ionic. These attractive forces are called
electrostatic attractions.
• In addition to NaCl, what are some examples of compounds
with ionic bonds?
Covalent Bonding
• In covalent bonding, electrons are shared rather than
transferred.
• Most elements tend to form covalent bonds rather than ionic
bonds because a gain or loss of multiple electrons (to achieve
the octet) is too high in energy.
e.g. carbon would have to lose 4 electrons or gain 4 electrons
in order to participate in ionic bonding.
• What are some examples of compounds with covalent bonds?
Common Bonding Patterns in Organic Compounds and Ions
Atom
Valence
Electrons
B
3
C
4
Positively
Charged
Neutral
(no octet)
+
C
Negatively
Charged
B
B
C
C
N
N
O
O
Cl
Cl
(no octet)
N
O
halogen
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+
N
6
7
O
+
+
Cl
• Equal sharing of electrons: nonpolar covalent bond
(e.g., H2)
• Sharing of electrons between atoms of different
electronegativities: polar covalent bond (e.g., HF)
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A polar covalent bond has a slight positive charge on one
end and a slight negative charge on the other
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A Polar Bond Has a Dipole Moment
• A polar bond has a negative end and a positive end
dipole moment (D) = m = e x d
(e) : magnitude of the charge on the atom
(d) : distance between the two charges
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Molecular Dipole Moment
The vector sum of the magnitude and the direction of the individual
bond dipole determines the overall dipole moment of a molecule
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Electrostatic Potential Maps
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Lewis Structures
• Lewis structures are representations of compounds in which lines and dots
are used to indicate electrons. A bond line is equal to 2 electrons.
• Keep in mind the number of valence electrons that each atom should have
(i.e. In which group is the atom located?).
• If the atoms in a molecule are to contain charges, think about
electronegativity and which atoms will better bear the particular charge.
Formal Charge
• Formal charge is the charge assigned to individual atoms in a Lewis
structure.
• By calculating formal charge, we determine how the number of electrons
around a particular atom compares to its number of valence electrons.
Formal charge is calculated as follows:
• The number of electrons “owned” by an atom is determined by its
number of bonds and lone pairs.
• An atom “owns” all of its unshared electrons and half of its shared
electrons.
Formal Charge
• Determine the formal charge for each atom in the following
molecule:
H
O
H
H
Nitrogen has five valence electrons
Carbon has four valence electrons
Hydrogen has one valence electron and halogen has
seven
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Important Bond Numbers
Neutral
Cationic
Anionic
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Non-Octet Species
• In
the 3rd and 4th rows, expansion beyond the octet to 10
and 12 electrons is possible.
Sulfuric Acid
Periodic Acid
Phosphoric Acid
• Reactive
species without an octet such as radicals,
carbocations, carbenes, and electropositive atoms
(boron, beryllium).
Nitric Oxide Radical,
Radical
Mammalian
Signaling Agent
Carbocation
Carbene
Borane
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Practice Problems
• Count the number of carbon atoms in each of the
following drawings.
O
O
a
b
c
O
OH
d
e
f
How to Draw Line Angle Structures
• Carbon atoms in a straight chain are drawn in a zigzag format.
• When drawing double bonds, try to draw the other bonds as
far away from the double bond as possible.
• When drawing each carbon atom in a zigzag, try to draw all of
the bonds as far apart as possible.
• In line angle structures, we do draw any H’s that are
connected to atoms other than carbon.
• It is good practice to draw in the lone pairs for heteroatoms.
An orbital tells us the volume of space around the nucleus
where an electron is most likely to be found
The s Orbitals
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The p Orbitals
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Molecular Orbitals
• Molecular orbitals belong to the whole molecule.
• s bond: formed by overlapping of two s orbitals.
• Bond
strength/bond dissociation: energy required to
break a bond or energy released to form a bond.
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In-phase overlap forms a bonding MO; out-of-phase
overlap forms an antibonding MO:
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Sigma bond (s) is formed by end-on overlap of two
p orbitals:
A s bond is stronger than a p bond
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Pi bond (p) is formed by sideways overlap of two parallel
p orbitals:
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Bonding in Methane
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Hybridization of One s and Three p Orbitals
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The orbitals used in bond formation determine the
bond angles
• Tetrahedral bond angle: 109.5°
• Electron pairs spread themselves into space as far from
each other as possible
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The Bonds in Ethane
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Hybrid Orbitals of Ethane
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Bonding in Ethene: A Double Bond
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Bonding in Ethyne: A Triple Bond
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Bonding in the Methyl Cation
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Bonding in the Methyl Radical
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Bonding in the Methyl Anion
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Bonding in Water
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Bonding in Ammonia and in the Ammonium Ion
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Bonding in Hydrogen Halides
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Summary
• The shorter the bond, the stronger it is
• The greater the electron density in the region of orbital
overlap, the stronger is the bond
• The more s character, the shorter and stronger is the
bond
• The more s character, the larger is the bond angle
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Brønsted–Lowry Acids and Bases
• Acid donates a proton
• Base accepts a proton
• Strong reacts to give weak
• The weaker the base, the stronger is its conjugate acid
• Stable bases are weak bases
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An Acid/Base Equilibrium


[H 3O ][ A ]
Ka 
[ H 2O][ AH]
LogK a  pK a

Ka: The acid dissociation constant.
The stronger the acid, the larger its Ka value and
the smaller its pKa value.

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The Most Common Organic Acids Are Carboxylic Acids
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Protonated alcohols and protonated carboxylic
acids are very strong acids
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An amine can behave as an acid or as a base
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Strong Acids / Bases React to Form Weak
Acids / Bases
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The Structure of an Acid Affects Its Acidity
• The weaker the base, the stronger is its conjugate
acid
• Stable bases are weak bases
• The more stable the base, the stronger is its conjugate
acid
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The stability of a base is affected by its size and its
electronegativity
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• When atoms are very different in size, the stronger
acid will have its proton attached to the largest atom
size overrides electronegativity
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• When atoms are similar in size, the stronger acid will
have its proton attached to the more electronegative atom
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Substituents Affect the Strength of an Acid
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• Inductive electron withdrawal increases the acidity of a
conjugate acid
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Acetic acid is more acidic than ethanol
The delocalized electrons in acetic acid are shared by
more than two atoms, thereby stabilizing the conjugated
base
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A Summary of the Factors That Determine Acid Strength
1. Size: As the atom attached to the hydrogen increases
in size, the strength of the acid increases
2. Electronegativity
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3. Hybridization
4. Inductive effect
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5. Electron delocalization
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Lewis Acids and Bases
• Lewis acid: non-proton-donating acid; will accept two
electrons
• Lewis base: electron pair donors
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