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Organic Chemistry
4th Edition
Paula Yurkanis Bruice
Chapter 1

Electronic Structure
and
Bonding

Acids and Bases

Irene Lee
Case 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
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
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) = m = e x d
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
I
Cl
Br
one bond
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 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 s bond and two p 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 p bond is weaker than a s 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
An Acid/Base Equilibrium
Ka: The acid dissociation constant
The Henderson–Hasselbalch Equation
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
Lewis Acids and Bases