Welcome to Organic Chemistry 234!

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

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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, Chemical Ecology 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

Fall 2012 Dr. HalliganCHM 234

• Electronic Structure and Bonding• Acids and Bases

Chapter 1

“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

Structure and Bonding

10 Note: Sections 1.1 and 1.2 on the structure of an atom can be reviewed in the textbook.

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

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

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

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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

Practice Problems

• Count the number of carbon atoms in each of the following drawings.

O

O

OH

O

a b c

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.

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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