welcome to organic chemistry 234! how should i study? do not memorize everything! practice writing...
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Welcome to Organic Chemistry 234!
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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.
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What is Organic Chemistry?• It is the study of carbon-containing compounds
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• 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.
Why Carbon?
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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 5-Fluorouracil, Colon Cancer Drug
Tamiflu, Influenza DrugAZT, HIV Drug
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Examples of Organic Compounds Used as Drugs
Haldol, AntipsychoticElavil, Antidepressant
Prozac, Antidepressant Viagra, TreatsErectile Dysfunction
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Spring 2011 Dr. HalliganCHM 234
• Electronic Structure and Bonding• Acids and Bases
Chapter 1
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“Speaking Organic Chemistry”
• What are some of the fundamentals of organic chemistry that we will cover in Chapter 1?
• 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
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Structure and Bonding
10 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?
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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?
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Common Bonding Patterns in Organic Compounds and Ions
AtomValence Electrons
Positively Charged Neutral
Negatively Charged
B
C
N
O
halogen
3
4
5
6
7
C
N
O
Cl
C
N
O
Cl
C
N
O
Cl
B B
+
+
+
+
(no octet)
(no octet)
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• 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) = = e x d
(e) : magnitude of the charge on the atom
(d) : distance between the two charges
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The vector sum of the magnitude and the direction of the individual bond dipole determines the overall dipole moment of a molecule
Molecular Dipole Moment
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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.
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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.
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Formal Charge
• Determine the formal charge for each atom in the following molecule:
H O
H
H
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Nitrogen has five valence electrons
Carbon has four valence electrons
Hydrogen has one valence electron and halogen hasseven
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Important Bond NumbersNeutral
Cationic
Anionic
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Non-Octet Species
Sulfuric Acid Periodic Acid Phosphoric Acid
• In the 3rd and 4th rows, expansion beyond the octet to 10 and 12 electrons is possible.
• Reactive species without an octet such as radicals, carbocations, carbenes, and electropositive atoms (boron, beryllium).
Nitric Oxide Radical,Mammalian
Signaling Agent
Radical Carbocation Carbene Borane
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Practice Problems
• Count the number of carbon atoms in each of the following drawings.
O
O
OH
O
a b c
d e f
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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.
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The s Orbitals
An orbital tells us the volume of space around the nucleuswhere an electron is most likely to be found
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The p Orbitals
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Molecular Orbitals
• Molecular orbitals belong to the whole molecule.
• 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 () is formed by end-on overlap of two p orbitals:
A bond is stronger than a bond
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Pi bond () 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
Ka: The acid dissociation constant.
The stronger the acid, the larger its Ka value and the smaller its pKa value.
Ka [H3O
][A ][H2O][AH ]
LogKa pKa
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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 acid: non-proton-donating acid; will accept two electrons
• Lewis base: electron pair donors
Lewis Acids and Bases