Acid-Base and Donor-Acceptor
Chemistry
Chapter 6
Acid-Base Concepts
• Arrhenius concept
– Acids from hydrogen ions (H3O+) in aqueous
solutions and bases from hydroxide ions in aqueous
solutions.
• HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)
[H 3O  ][Cl ]
Ka 
[HCl]
• NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
Cannot be used for nonaqueous solutions.
Acid-Base Concepts
• Brønsted-Lowry concept
– Acid – a species with a tendency to lose a hydrogen
ion.
– Base – a species with a tendency to gain a hydrogen
ion.
H3O+ + NO2-  H2O + HNO2
Conjugate acid-base pairs differ by a proton.
A reaction always favors the formation of weaker
acids.
Acid-Base Concepts
• Lewis concept
– A base is an electron-pair donor and an acid is an electron-pair
acceptor. The product is often called an adduct.
– BF3 + NH3  F3BNH3 (draw and discuss)
• MO treatment (HOMO-LUMO interaction)
– Figure 6-1 and Figure 6-2 (illustrate with Spartan)
• Boron has nearly a tetrahedral geometry in the adduct.
– Frontier orbitals are those at the occupied-unoccupied frontier
(commonly the HOMO-LUMO).
• In most acid-base reactions, a HOMO-LUMO combination forms new
HOMO and LUMO orbitals of the product.
– Commonly, the boiling points raise significantly upon
formation of the adduct.
– Adducts involving metal ions are called coordination
compounds.
Lewis Concept
• NH3 + H+  NH4+ (C3v  Td), Figure 6-4.
– The HOMO of NH3 has A1 symmetry and the LUMO of H+
can be assumed to have A1 symmetry.
• HOMO (NH3) = -10.4 eV, LUMO (H+) = -11.3 eV
– Upon formation of the adduct, NH4+, the symmetry changes
from C3v to Td and a new set of orbitals are formed.
• The a1 orbital of NH3 (previously nonbonding) is stabilized by this
interaction. It becomes part of the triply-degenerate t2 orbital set.
• 4 bonding and 4 antibonding orbitals are produced from the minimal
basis set (a1 and t2 sets).
– There is a net lowering in energy when the adduct forms since
the largely nonbonding a1 becomes a bonding t2.
Illustrate this with Spartan software.
Hydrogen Bonding – A Frontier
Molecular Orbital Approach
• FHF- The MOs are be generated by combining the MOs
of HF with F-.
– The px and py lone pair orbitals on the fluorines are essentially
nonbonding.
• No matching orbitals on the hydrogen atom.
– Overlap exists between the /* of HF and the  of F-.
– Three MOs are produced
• Low-energy MO that is bonding (no nodes between atoms).
• Intermediate-energy MO that is close to nonbonding.
• High-energy MO that is usually antibonding (usally possessing nodes
between atoms).
Hydrogen Bonding – BHA
• The pattern for bonding is essentially the same (Fig. 6-7).
– The resulting BHA structure is lower energy.
• Three possibilities for B + HA interaction (Fig. 6-8).
– Poor match of HOMO-LUMO energies; no H-bonding occurs.
Occupied reactant orbitals are lower in energy.
• H2O + CH4
– Good match of HOMO-LUMO energies; H-bonding occurs.
Occupied product orbitals are lower in energy.
• H2O + HF
– Very poor match of HOMO-LUMO energies; proton transfer.
BH+A is lower than B + HA
• H2O + HCl
The Electronic Spectra of I2
• The changes in the electronic energy levels of I2 cause
changes in visible color and the spectra.
• The g*  u* transition (Figure 6-9).
– Gas phase I2 is violet.
• Absorption removes the middle yellow, green, and blue parts of the
visible spectrum, leaving red and violet at the opposite ends
(complimentary colors).
– In nondonor solvents (i.e. hexane), the color remains the same.
There is essentially no interaction.
– In donor solvents (I.e. methanol), the color becomes brown.
• Interaction of 9u* with the donor orbital produces a lower energy
bonding orbital and a higher-energy unoccupied orbital. The
absorbance color (higher-energy) shifts toward the blue and brown is
transmitted.
The Electronic Spectra of I2
• In water, I2 is yellow-brown due to the average donor
characteristics of water.
• I- is a good donor and the observed color is brown upon
formation of I3-.
• As the donor capability increases, the energy of the g*
 u* transition increases.
• Charge transfer is due to the g  u*.
– The transition causes the electron to move largely from a donor
orbital to an acceptor orbital.
h
I 2  donor [I 2 ]  [donor ]
CT
Examine Figure 6-10 to examine changes in the transition bands.
Hard and Soft Acids and Bases
• Factors other than acid/base strength determine
the acid/base reactivity.
– Silver halide solubility;
2O
AgX (s) H

 Ag  (aq)  X  (aq)
• Ksp’s: AgF=205, AgCl=1.810-10, AgBr=5.2 10-13, AgI = 8.3 10-17
• Solvation: F- is much better solvated.
• Degree of Ag-X interaction also plays a role.
• Hard soft acid base (HSAB) theory can help explain the data
– Hard acids and hard bases are small and nonpolarizable.
– Soft acids and soft bases are large and polarizable.
– Hard/hard and soft/soft interactions are most favorable.
• Polarizable – easily distorted by other charged ions.
HSAB Theory
• AgX data
– Ag+ is large and polarizable.
– Softness of halides; I->Br->Cl->F– AgI has the strongest interaction, and, therefore, the lowest
solubility.
• Softness is also associated with covalent bonding.
• Color of the AgX compounds depend on the HOMOLUMO energy difference.
– Large difference – energy is in the UV range. The compound
would be colorless or white.
• AgCl and AgF involve hard ions. Electron transfers between very
different energy levels.
HSAB Theory
• Color of the AgX compounds depend on the HOMOLUMO energy difference.
– Small difference – low-energy transition moves the
absorption into the visible region.
• AgI and AgBr involve soft, covalent interactions. Covalency
generally implies electron sharing between similar orbital energies.
These compounds are usually colored.
– Color and low solubility usually go with soft-soft
interactions; colorless compounds and high solubility
usually go with hard-hard interactions (not always). An
exception is the lithium halide compounds (next slide).
The Lithium Halides
• Solubility trend; LiBr>LiCl>LiI>LiF
– Li+ and F- can both be considered as hard ions. The
interaction between them is predicted to be stable producing
an ionic and soluble compound.
– Favorable hard-hard interactions overcome solubility.
– LiBr and LiCl are more soluble because of less favored
interactions.
– LiI is out of order due to poor I- solvation.
The solubility property is used in some cases to support HSAB.
An expected trend in solubility does not necessarily indicate
that the HSAB theory is invalid. The LiF interaction is stil
the strongest among the lithium halides. Many factors must
be considered when interpreting solubility.
Pearson’s Hard and Soft Acids
and Bases
Most of the hard-soft distinction depends on polarizability.
• Hard acids and bases are small, compact, and nonpolarizable.
– Most metal ions are hard acids (class a ions, Fig. 6-11).
– Hardness of a metal ion generally increases with its charge.
• Soft acids are generally large and polarizable (class b ions).
– Generally, more electrons and larger sizes lead to softer
behavior.
– Soft acids can often be characterized as having d
electrons/orbitals for  bnding.
• Tl(I) versus Tl(III) – Tl(III) has more class be character even though the
charge on the metal is greater. Why?
Pearson’s Hard and Soft Acids
and Bases
• Reactions favor hardness matches (Fig. 6-11, Table 6-3
and Table 6-4).
– Hard=small, compact charge, and nonpolarizable.
• M+3, O-2
– Soft=large, polarizable.
• M0, S2Example and Exercise 6-4.
– Comparison is easiest within a column of the periodic table.
Any solubility predictions based on HSAB theory must be
considered tentative. Solvent and other interactions must be
considered carefully.
HOMO-LUMO Considerations
• Hard-hard interactions can be viewed as simple
electrostatic interactions with the LUMO far above the
HOMO in energy.
– Little change in orbital energies.
• Soft-soft interactions involved HOMO-LUMO energies
that are much closer in energy.
– Large change in orbital energies.
Many studies indicate that the hard-hard interaction is generally
stronger than the soft-soft interaction.
Interpreting Fig. 6-12 should be done with caution. This is a
simplistic interpretation.
Quantitative Measures, Drago
and Weyland
• Includes electrostatic and covalent factors to account
more fully for reactivity of an acid-base system.
– -H=EAEB+CACB
• H is the enthalpy of reaction, A+BAB
• E is a measure of the capacity for electrostatic interactions.
• C is a measure of the tendency to form covalent bonds.
– I2 is the reference acid (CA and EA = 1); N,Ndimethylacetamide and diethyl sulfide are the reference bases.
• I2 has very little electrostatic attraction for bases (small EA value);
however, it has a strong tendency to form covalent bonds (high CA
values)
Table 6-7
Drago and Weyland
• I2 + C6H6  I2C6H6
• Another example
• Drago emphasizes to factors involved in acid-base
strength
– Electrostatic and covalency
• If E and C values have been tabulated, Drago’s method
can predict acid-base strength fairly accurately. If no
data is available, Pearson’s HSAB method can be useful
for rough predictions. Neither method covers every
case. Additionally, neither method includes solvation
effects.
Acid and Base Strength
• Binary hydrogen compounds (HF, HCl, HBr)
– HF<<HCl<HBr
– Acidity increases going down a column even
though the electronegativity decreases.
Explanation: The larger conjugate bases have lower
charge densities and smaller attraction for the
hydrogen ion.
Weak conjugate base  strong acid
Acid and Base Strength
• Binary hydrogen compounds
– The acidity increases across a period (the
electronegativity also increases).
• HF>H2O>NH3
Explanation: The negative charge of the conjugate
base is spread out over more lone pairs (electrons) as
the base gets larger. The larger the number of lone
pairs, the lower the attraction for protons.
Acid-Base Strength, Inductive
Effects
• Electronegativity effects can can change the
acidity and basicity.
– PF3 is a weaker base than PH3 due to the electrons
being drawn toward the electronegative fluorines.
NMe3>NHMe2>NH2Me>NH3 (base strength)
Alkyl groups tend to increase the electron density on
the center atom.
Electron-contributing and electron-withdrawing
capabilities of the ligands need to be considered.
Acid-Base Strength, Oxyacids
• HClO4>HClO3>HClO2>HClO
– pKa  9-7n
• n = number of nonhydrogenated oxygen atoms.
Explanation: Nonhydrogenated oxygens draw electron density
from the central atom. The central atom, in turn, draws
electron density away from he hydrogenated oxygen. The net
result is a weaker O-H bnd.
Other explanation: The charge of the conjugate base is spread
over all the nonhydrogenated oxygens. Stability of the
conjugate base increases with the number of these oxygens
producing a weaker base.
Acidity of Cations in Aqueous Solutions
• In general, stronger acids form from metal ions with larger charges
and smaller radii.
– [Fe(H2O)6]3+ + H2O  [Fe(H2O)5OH]2+ + H3O+
• Solubility of the metal hydroxide is also a measure of cation
acidity. The stronger the cation acid, the less soluble the
hydroxide.
– 3+ metal ions form hydroxides that precipitate.
• Fe+3(aq) + 3H2O(l)  Fe(OH)3(s) + 3H+(aq)
– 2+ d-block ions (and Mg+2) from hydroxide precipitates in
neutral or slightly acidic solutions
– Alkali (e.g. Na+) and alkaline earth metals produce no pH
changes (spectators).
• The free metal cation is no longer detectable for highly-charged
species.
– UO2+ and CrO42-
Steric Effects
• Repulsions by bulky groups make reactions less
favorable for adduct formation.
• Different types of strain (Brown):
– F (front) – bulky groups interfere directly with the approach.
– B (back) – bulky groups interfere with each other when they
bend away from the other molecule forming the adduct.
– I (internal) – electronic differences within similar molecules.
Examples
Solvation
• In aqueous solution the basicities have the order
– Me2NH>MeNH2>Me3N>NH3 Why?
– The reduced solvation of their protonated cations.
• Solvation energies; RNH3+>R2NH2+>R3NH+ for the
reaction RnH4-nN+(g) + H2O  RnH4-nN+(aq). Solvation
is dependent upon the number of hydrogen atoms
available to hydrogen bond.
– Competition between induction and solvation
produce the observed order for the amines.
Nonaqueous Solvents
• Any acid stronger than H3O+ reacts completely
with H2O to form H3O+. The is called the
leveling effect; acids or bases are brought down
to the limiting conjugate acid or base of the
solvent.
– HCl, HBr, HClO4, and HNO3 all have the same
strength in water.
• Likewise, the strongest base in water is OH-.
– Water can differentiate weak acids and bases.
Nonaqueous Solvents
• Acetic acid can differentiate the strengths of many
strong acids.
– H2SO4 + HOAc  H2OAc+ + HSO4– In acetic acids, HClO4>HCl>H2SO4>HNO3
• Inert solvents allow a broad range of acids and
bases to be differentiated according to strength.
The do not form acid or base species readily.
– Figure 6-17.