Transcript Lecture Resource ()

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 Chemistry
• 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
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 atoms
in 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 Dipole
• A polar bond has a negative end and a positive end
dipole moment (D) = m = e x d
(e) : magnitude of the charge on the atom
(d) : distance between the two charges
Lewis Structure
Formal charge =
number of valence electrons –
(number of lone pair electrons +1/2 number of bonding electrons)
Important Bond Numbers
one bond
two bonds
three bonds
four bonds
H
F
Cl
Br
O
N
C
I
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
An sp-Hybridized Carbon
• The bond angle in the sp2 carbon is 120°
• The sp2 carbon is the trigonal planar carbon
Bonding in Ethyne: A Triple Bond
• A triple bond consists of one s bond and two p bonds
• Bond angle of the sp carbon: 180°
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
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
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
H2O + HA
H3O+ + A-
[H3O+][A-]
Ka =
[H2O][HA]
pKa = -log Ka
Ka: The acid dissociation constant
The Henderson–Hasselbalch Equation
The pH indicates the concentration of hydrogen ions (H+)
pK a

HA 
 pH  log
A 

• 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
• 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
O
CH3COH
CH3CH2OH
pKa = 4.76
acetic acid
pKa = 15.9
ethanol
The delocalized electrons in acetic acid are shared by
more than two atoms, thereby stabilizing the conjugated
base
O
CH3CO-
O
CH3CO-
Lewis Acids and Bases
• Lewis acid: non-proton-donating acid; will accept two
electrons
• Lewis base: electron pair donors