Organic Chemistry 4th Edition

Paula Yurkanis Bruice

Chapter 1

Electronic Structure and

Bonding

Acids and Bases

Irene LeeCase Western Reserve University

Cleveland, OH©2004, Prentice Hall

• Organic compounds are compounds containing carbon

• Carbon neither readily gives up nor readily accepts electrons

• Carbon shares electrons with other carbon atoms as well as with several different kinds of atoms

Organic Chemistry

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

The Distribution of Electrons in an Atom

• Quantum mechanics uses the mathematical equation of wave motions 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 is most likely to be found

• The atomic orbital closer to the nucleus has the lowest energy

• Degenerate orbitals have the same energy

Table 1.1

• The Aufbau principle: electrons occupy the orbitals with the lowest energy first

• 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

• Ionic compounds are formed when an electropositive element transfers electron(s) to an electronegative element

Covalent Compounds• Equal sharing of electrons: nonpolar covalent bond (e.g., H2)

• Sharing of electrons between atoms of different electronegativities: polar covalent bond (e.g., HF)

Electrostatic Potential Maps

• 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

A Dipole

Lewis Structure

Formal charge = number of valence electrons –(number of lone pair electrons +1/2 number of bonding electrons)

Important Bond Numbers

H F ICl Brone bond

Otwo bonds

Nthree bonds

Cfour bonds

The s Orbitals

The p Orbitals

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

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

Sigma bond (σ) is formed by end-on overlap of two p orbitals

A σ bond is stronger than a π bond

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

Bonding in Methane and Ethane: Single Bonds

Hybridization of 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

Hybrid Orbitals of Ethane

Bonding in Ethene: A Double Bond

• The bond angle in the sp2 carbon is 120°

• The sp2 carbon is the trigonal planar carbon

An sp-Hybridized Carbon

Bonding in Ethyne: A Triple Bond

• Bond angle of the sp carbon: 180°

• A triple bond consists of one σ bond and two π bonds

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

• A π bond is weaker than a σ bond

• 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

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

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

H2O + HA H3O+ + A-

[H3O+][A-]Ka =

[H2O][HA]

pKa = -log Ka

The Henderson–Hasselbalch Equation

• A compound will exist primarily in its acidic form at a pH < its pKa

• A compound will exist primarily in its basic form at a pH > its pKa

• A buffer solution maintains a nearly constant pH upon addition of small amount of acid or base

[ ][ ]−+=A

HAlogpHp aK

The pH indicates the concentration of hydrogen ions (H+)

• When atoms are very different in size, the stronger acid will have its proton attached to the largest atom

• When atoms are similar in size, the stronger acid will have its proton attached to the more electronegative atom

• Inductive electron withdrawal increases the acidity of a conjugate acid

Acetic acid is more acidic than ethanol

CH3COH

O

CH3CH2OH

pKa = 4.76 pKa = 15.9acetic acid ethanol

The delocalized electrons in acetic acid are shared by more than two atoms, thereby stabilizing the conjugated base

CH3CO-

O

CH3CO-

O

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

• Lewis base: electron pair donors

Lewis Acids and Bases