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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.

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Welcome to Organic Chemistry 234!

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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.


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.


Examples of Organic Compounds Used as Drugs

Methotrexate, Anticancer Drug

5-Fluorouracil, Colon Cancer Drug

Tamiflu, Influenza Drug

AZT, HIV Drug


Examples of Organic Compounds Used as Drugs

Haldol, Antipsychotic

Elavil, Antidepressant

Prozac, Antidepressant

Viagra, TreatsErectile Dysfunction


Fall 2012

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?

  • 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.


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


  • Equal sharing of electrons: nonpolar covalent bond

  • (e.g., H2)

  • Sharing of electrons between atoms of different

  • electronegativities: polar covalent bond (e.g., HF)


A polar covalent bond has a slight positive charge on one

end and a slight negative charge on the other


(e) : magnitude of the charge on the atom

(d) : distance between the two charges

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


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


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:


Nitrogen has five valence electrons

Carbon has four valence electrons

Hydrogen has one valence electron and halogen has

seven


Important Bond Numbers

Neutral

Cationic

Anionic


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,

Mammalian Signaling Agent

Radical

Carbocation

Carbene

Borane


Practice Problems

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


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


The p Orbitals


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.


In-phase overlap forms a bonding MO; out-of-phase

overlap forms an antibonding MO:


Sigma bond (s) is formed by end-on overlap of two

p orbitals:

A s bond is stronger than a p bond


Pi bond (p) is formed by sideways overlap of two parallel

p orbitals:


Bonding in Methane


Hybridization of One s and Three p Orbitals


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


The Bonds in Ethane


Hybrid Orbitals of Ethane


Bonding in Ethene: A Double Bond


Bonding in Ethyne: A Triple Bond


Bonding in the Methyl Cation


Bonding in the Methyl Radical


Bonding in the Methyl Anion


Bonding in Water


Bonding in Ammonia and in the Ammonium Ion


Bonding in Hydrogen Halides


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


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


An Acid/Base Equilibrium

Ka: The acid dissociation constant.

The stronger the acid, the larger its Ka value and the smaller its pKa value.


The Most Common Organic Acids Are Carboxylic Acids


Protonated alcohols and protonated carboxylic acids are very strong acids


An amine can behave as an acid or as a base


Strong Acids / Bases React to Form Weak Acids / Bases


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


The stability of a base is affected by its size and its

electronegativity


  • When atoms are very different in size, the stronger

  • acid will have its proton attached to the largest atom

size overrides electronegativity


  • When atoms are similar in size, the stronger acid will

  • have its proton attached to the more electronegative atom


Substituents Affect the Strength of an Acid


  • Inductive electron withdrawal increases the acidity of a

  • conjugate acid


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


A Summary of the Factors That Determine Acid Strength

  • Size: As the atom attached to the hydrogen increases

  • in size, the strength of the acid increases

2. Electronegativity


3. Hybridization

4. Inductive effect


5. Electron delocalization


Lewis Acids and Bases

  • Lewis acid: non-proton-donating acid; will accept two

  • electrons

  • Lewis base: electron pair donors


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