56 CHAPTER 1 Electronic Structure and Bonding • Acids and Bases
bondsin organic compounds are bonds, a double bond
consists of one bond and one bond, and a triple bond
consists of one bond and two bonds. Triple bonds are
shorter and stronger than double bonds, which are shorter
and stronger than single bonds. To form four bonds, carbon
promotes an electron from a 2sto a 2porbital. C, N, and O
form bonds using hybrid orbitals. The hybridizationof C,
N, or O depends on the number of bonds the atom forms:
No bonds means that the atom is hybridized, one
bond indicates that it is hybridized, and two bonds
signifies that it is sphybridized. Exceptions are carboca-
tions and carbon radicals, which are hybridized. The
more scharacter in the orbital used to form a bond, the
shorter and stronger the bond is and the larger the bond
angle is. Bonding and lone-pair electrons around an atom
are positioned as far apart as possible.
An acidis a species that donates a proton, and a baseis a
species that accepts a proton. A Lewis acidis a species that
accepts a share in an electron pair; a Lewis baseis a species
that donates a share in an electron pair.
Acidityis a measure of the tendency of a compound to
give up a proton. Basicityis a measure of a compound’s
affinity for a proton. The stronger the acid, the weaker is its
conjugate base. The strength of an acid is given by the acid
dissociation constant(Ka).Approximate pKavalues are as
sp^2sp^2 pp sp^3 pps ps ps follows: protonated alcohols, protonated carboxylic acids,
protonated water carboxylic acids protonated
amines alcohols and water The pHof a solution
indicates the concentration of positively charged hydrogen
ions in the solution. In acid–base reactions, the equilibrium
favors reaction of the strong and formation of the weak.
The strength of an acid is determined by the stability of
its conjugate base: The more stable the base, the stronger is
its conjugate acid. When atoms are similar in size, the more
acidic compound has its hydrogen attached to the more elec-
tronegative atom. When atoms are very different in size, the
more acidic compound has its hydrogen attached to the larg-
er atom. Inductive electron withdrawalincreases acidity;
acidity decreases with increasing distance between the elec-
tron-withdrawing substituent and the ionizing group.
Delocalized electronsare electrons shared by more than
two atoms. A compound with delocalized electrons has
resonance. The resonance hybridis a composite of the
resonance contributors, which differ only in the location of
their lone-pair and electrons.
The Henderson–Hasselbalch equationgives the rela-
tionship between and pH: A compound exists primarily
in its acidic form in solutions more acidic than its value
and primarily in its basic form in solutions more basic than
its value.pKapKapKap'10; '15.6 0; '5;Key Terms
acid (p. 39)
acid–base reaction (p. 40)
acid dissociation constant (p. 41)
acidity (p. 40)
antibonding molecular orbital (p. 21)
atomic number (p. 4)
atomic orbital (p. 5)
atomic weight (p. 4)
aufbau principle (p. 6)
base (p. 39)
basicity (p. 40)
bond (p. 8)
bond dissociation energy (p. 21)
bond length (p. 21)
bonding molecular orbital (p. 21)
bond strength (p. 21)
buffer solution (p. 53)
carbanion (p. 14)
carbocation (p. 14)
condensed structure (p. 16)
conjugate acid (p. 40)
conjugate base (p. 40)
constitutional isomer (p. 18)
core electrons (p. 7)
covalent bond (p. 9)
degenerate orbitals (p. 5)
delocalized electrons (p. 50)
debye (D) (p. 12)
(Ka)dipole (p. 12)
dipole moment (p. 12)
double bond (p. 30)
electronegative (p. 8)
electronegativity (p. 10)
electropositive (p. 7)
electrostatic attraction (p. 8)
electrostatic potential map (p. 11)
equilibrium constant (p. 41)
excited-state electronic configuration (p. 6)
formal charge (p. 13)
free radical (p. 14)
ground-state electronic configuration (p. 6)
Heisenberg uncertainty principle (p. 18)
Henderson–Hasselbalch equation (p. 51)
Hund’s rule (p. 7)
hybrid orbital (p. 26)
hydride ion (p. 9)
hydrogen ion (p. 9)
inductive electron withdrawal (p. 47)
ionic bond (p. 8)
ionic compound (p. 9)
ionization energy (p. 7)
isotopes (p. 4)
Kekulé structure (p. 16)
Lewis acid (p. 54)
Lewis base (p. 54)
Lewis structure (p. 13)1 m 2lone-pair electrons (p. 13)
mass number (p. 4)
molecular weight (p. 4)
molecular orbital (p. 20)
molecular orbital (MO) theory (p. 20)
node (p. 18)
nodal plane (p. 19)
nonbonding electrons (p. 13)
nonpolar covalent bond (p. 10)
nonpolar molecule (p. 26)
octet rule (p. 7)
orbital (p. 4)
orbital hybridization (p. 26)
organic compound (p. 2)
Pauli exclusion principle (p. 6)
pH (p. 41)
pi bond (p. 23)
(p. 41)
polar covalent bond (p. 10)
proton (p. 9)
proton-transfer reaction (p. 40)
quantum mechanics (p. 4)
radial node (p. 18)
radical (p. 18)
resonance (p. 51)
resonance contributors (p. 50)
resonance hybrid (p. 50)
sigma bond (p. 20)(s)pKa(p)