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Organic Chemistry6th Edition
Paula Yurkanis Bruice
Chapter 1
Electronic Structure and
Bonding
Acids and Bases
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• Carbon-containing compounds were once considered “organ compounds” available only from living organisms.
• The synthesis of the simple organic compound urea in 1828 showed that organic compounds can be prepared in the laboratory from non-living material.
• Today, organic natural products are routinely synthesized in the laboratory.
Organic Chemistry
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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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The Structure of an Atom
• An atom consists of electrons, positively charged protons,and neutral neutrons.
• Electrons form chemical bonds.
• Atomic number: numbers of protons in its nucleus
• Mass number: the sum of the protons and neutrons of an atom
• Isotopes have the same atomic number but different mass numbers.
• The atomic weight: the average weighted mass of its atoms
• Molecular weight: the sum of the atomic weights of all the atomsin the molecule
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The Distribution of Electrons in an Atom
• Quantum mechanics uses the mathematical equation of wavemotions to characterize the motion of an electron around a nucleus.
• Wave functions or orbitals tell us the energy of the electron and the volume of space around the nucleus where an electron ismost likely to be found.
• The atomic orbital closer to the nucleus has the lowest energy.
• Degenerate orbitals have the same energy.
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The ground-state electronic configuration describes the orbitalsoccupied by the atom’s electrons with the lowest energy
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• The Aufbau principle: an electron always goes to theavailable orbital with the lowest energy
• The Pauli exclusion principle: only two electrons can occupy one atomic orbital and the two electrons have opposite spin
• Hund’s rule: electrons will occupy empty degenerated orbitals before pairing up in the same orbital
The following principles determine which orbitalselectrons occupy:
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Lewis’s theory: an atom will give up, accept, or share electrons inorder to achieve a filled outer shell or an outer shell that containseight electrons
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Ionic Bonds Are Formed by the Transfer of Electrons
Attractive forces between opposite charges are called electrostatic attractions
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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) = µ = e x d(e) : magnitude of the charge on the atom
(d) : distance between the two charges
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Lewis Structure
Formal charge = number of valence electrons –(number of lone pair electrons +1/2 number of bonding electrons)
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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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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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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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Molecular Orbitals
• Molecular orbitals belong to the whole molecule.
• σ 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 (σ) is formed by end-on overlap of two p orbitals:
A σ bond is stronger than a π bond
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Pi bond (π) is formed by sideways overlap of two parallel 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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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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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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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 Kavalue and the smaller its pKa value.
Ka =[H3O
+ ][ A− ][H2O][ AH ]
LogKa = pKa
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Protonated alcohols and protonated carboxylic acids arevery strong acids
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The Structure of an Acid Affects Its Acidity
• The weaker the base, the stronger is its conjugateacid
• Stable bases are weak bases
• The more stable the base, the stronger is its conjugateacid
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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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• 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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• Lewis acid: non-proton-donating acid; will accept two electrons
• Lewis base: electron pair donors
Lewis Acids and Bases
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Basicity and Drug Design: Methotrexate, Substituting Nitrogen for Oxygen