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Acid-Base Theories
Arrhenius Theory
Svante Arrenhius (1857-1927)
Acid: Substance that produces H+ in water.
Base: Substance that produces OH-1 in water.
HCl(aq)  H+ + Cl- produces H+ in water
NH3(aq)  NH4+ + OH- produces OH- in water
Although NH3 does not contain OH-, hydroxide ions form
when added to water.
Arrhenius acid and base neutralize each other to produce
salt and water:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O)(l)
H+(aq) + OH-(aq)  H2O(l)
Bronsted/Lowry Theory
Johannes Bronsted (1879-1947) Thomas Lowry (1874-1936)
Acid: Substance that can donate proton (H+).
Base: Substance that can accept proton (must
contain lone pair of electrons).
HCl + NH3  NH4+ + ClAcid base CA CB
Acids may be cations, neutral molecules, or anions, while bases may be
anions or neutral molecules. Just as a reduction must always accompany
an oxidation, a proton donor (acid) must accompany a proton acceptor
(base). Once an acid transfers its proton it becomes the conjugate base
(CB) and once a base accepts the proton it becomes the conjugate acid
(CA). Since protons are always transferred in the Arrenhius concept, all
Arrhenius acid/base reactions are also Bronsted-Lowry acid/base
reactions. But if water is not involved (HCl & NH3), the reaction can be
explained by Bronsted/Lowry concept and not Arrenhius.
Solvent system concept of acids and bases:
Acid= cation of solvent via autodissociation
Base = anion of the solvent by autodissociation.
Solutes that increase the concentration of the cation of the
solvent are considered acids and soultes that increase the
concentration of the anion are considered bases
The solvent must be able to behave as both an acid and a
base (amphoteric)
2 H2O  OH- + H3O+
H2SO4 + H2O  H3O+ + HSO4- H2SO4 is an acid
2BrF3  BrF2+ + BrF4SbF5 + BrF3  BrF2+ + SbF6- SbF5 is an acid
KF + BrF3  BrF4- + K+ KF is a base
Lewis Theory
Gilbert Lewis (1875-1946)
Acid: Substance that can accept a pair of electrons from
another atom to form a new bond.
Base: Substance that can donate a pair of electrons to
another atom to form a new bond.
The product of Lewis acid-base reaction referred to as
adduct. The proton itself can act as Lewis acid. Lewis
expands acid/base reactions to include many substances
without H in formula.
F3B + NH3  F3B:NH3
Chapter 6 p166
Chapter 6 p169
Which theories can explain the following?
HI + H2O 
H 3 O + + I-
HI + NH3  NH4+ + II2 + NH3  NH3I+ + I-
I2 + Cl- ICl + IX:- + Y+  Y:X
Lewis Concept
The lone pair in the HOMO of the ammonia molecule combines with
the empty LUMO of the BF3, which has very large, empty orbital
lobes on boron, to form the adduct.
The B-F bonds in the product are bent away from the ammonia into a
nearly tetrahedral geometry around the boron
Boron trifluoride-diethyl ether adduct
Lone pair on the oxygen of the diethyl ether are attached to the
boron. The result is that one of the lone pairs bonds to the boron,
changing the geometry around B from planar to tetrahedral. As a
result, BF3, with a boiling point of -99.9 oC, and diethyl ether, with
a boiling point of 34.5 oC, form an adduct of 125 – 126 oC.
Lewis acid-base adducts involving metal ions are called
coordination compounds.
Ag+ + 2 :NH3  [H3N:Ag:NH3]+
Frontier orbitals and acid-base reactions
HOMO-LUMO interactions
In most acid-base reactions, a HOMO-LUMO combination form
new HOMO and LUMO of the product. Orbitals whose shapes
allow significant overlap and whose energies are similar form
useful bonding and antibonding orbitals. On the other hand, if the
orbital combinations have no useful overlap, no net bonding is
possible and they can not form acid-base product.
Even when the orbital shapes match, several reactions may be
possible, depending on the relative energies. A single species can
act as an oxidant, an acid, a base or a reductant, depending on the
other reactant
HOMO-LUMO interactions
1.
2H2O + Ca  Ca2+ + 2OH- + H2
water as oxidant
2.
nH2O + Cl- [Cl(H2O)n]water as acid
3.
6H2O + Mg2+  [Mg(H2O)6]2+
water as base
4. 2H2O + 2F2  4F- + 4H+ + O2
water as reductant
A base has an electron pair in a HOMO of suitable symmetry to
interact with the LUMO of the acid. The better the energy match
between the base’s HOMO and the acid’s LUMO, the stronger the
interaction.
Hydrogen bonding
4 nodes
3 nodes
2 nodes
The lowest orbital is
distinctly bonding, with all
three component orbitals
contributing and no nodes
between the atoms. The
middle (HOMO) orbitals is
essentially nobonding, with
nodes through each of the
nuclei. The highest energy
orbital (LUMO) is
antibonding, with nodes
between each pair of atoms
For unsymmetrical Hbonding B + HA  BHA,
the pattern is similar
Electronic spectra
Charge transfer: the transition
transfers an electron from an
orbital that is primarily of donor
composition to one that is
primarily of acceptor composition
I2Donor  [I2]- [Donor]+
Hard and Soft Acid and Bases
Ag F(s) + H2O  Ag+(ag) + F-(aq)
Ksp = 205
Ag Cl(s) + H2O  Ag+(ag) + Cl-(aq)
Ksp = 1.8 x 10-10
Ag Br(s) + H2O  Ag+(ag) + Br-(aq)
Ksp = 5.2 x 10-13
Ag I(s) + H2O  Ag+(ag) + I-(aq)
Ksp = 8.3 x 10-17
1.
Solvation of the ions is a factor in these reactions, with fluoride ion being
much strongly solvated than the other anions.
2.
Related to HSAB in which iodide is much softer (more polarizable) than
the others and interacts more strongly with silver ions, a soft cation. The
result is a more covalent bond.
Colors
AgI yellow
AgBr slightly yellow
AgCl and AgF white
Color depends on the difference in energy between occupied and
unoccupied orbitals. A large difference results in absorption in the
ultraviolet region of the spectrum; a smaller difference in energy
levels moves the absorption into the visible region. Compounds
absorbing violet appear to be yellow; as the absorption moves
toward lower energy, the color shifts and become more intense.
Black indicates very broad and very strong absorption.
Color and low solubility typically go with soft-soft interactions;
colorless compounds and high solubility generally go with hardhard interactions.
Color and low solubility typically go with soft-soft interactions;
colorless compounds and high solubility generally go with hardhard interactions, although some hard-hard combination have low
solubilities.
LiBr> LiCl > LiI > LiF
The solubilities show a strong hard-hard interaction in LiF that
overcomes the solvation of water, but the weaker hard-soft
interactions of the other halides are not strong enough to prevent
solvation and these halides are more soluble than LiF.
Fajan's Rules (Polarization)
Polarization will be increased by:-
1. High charge and small size of the cation
Ionic potential ?Z+/r+ (= polarizing power)
2. High charge and large size of the anion
The polarizability of an anion is related to the deformability of its electron cloud
(i.e. its "softness")
3. An incomplete valence shell electron configuration
noble gas configuration of the cation better shielding less polarizing power
i.e. charge factor in (1) should be effective nuclear charge
e.g. Hg2+ (r+ = 102 pm) is more polarising than Ca2+ (r+ = 100 pm)
Four rules can be summarized:
1. For a given cation, covalent character increases with increase in
size of the anion.
2. For a given anion, covalent character increases with decrease in
size of the cation.
3. Covalent character increase with increasing charge on either ion.
4. Covalent character is greater for cations with nonnoble gas
electronic configuration.
Q1. Ag2S is much less soluble than Ag2O
A1. Rule 1: S2- is much larger than O2Q2. Fe(OH)3 is much less soluble than Fe(OH)2
A2. Rule 3: Fe3+ has a larger charge than Fe2+
These rules are helpful in predicting behavior of specific
cation-anion interaction, but not enough
1. Li series does not fit
2. Solubility MgCO3 > CaCO3 >SrCO3 >BaCO3
Rule 2 predicts the reverse of the order. The difference
lies in the aquation of the metal ions. Mg2+ (small with
higher charge density) attracts water molecules much
more strongly than the others, with Ba2+ (large with
smaller charge density) the least strongly solvated.
Ahrland, Chatt and Davies:
Class (a) ions: Most metals
Class (b) ions: Cu2+, Pd2+, Ag+, Pt2+, Au+, Hg22+, Hg2+, Tl+, Tl3+, Pb2+,
and heavier transition metal ions
The class (b) ions form halides whose solubility is in the
order F- > Cl- > Br- > I-. The solubility of Class (a)
halide is in the reverse order. The calss (b) metals ions
also have a larger enthalpy of reaction with P donor than
with N donor, again the reverse order of the Class (a)
metal ion recations.
Class (b) – having d electrons available for  bonding
Tl(III) show stronger Class (b) character than Tl(I)
because Tl(I) has two 6s electrons that screen the 5d
electrons and keep them form being fully available for 
bonding
Pearson’s Principle:
Hard Lewis acids prefer to bind to hard Lewis bases; soft
Lewis acids prefer to bind to soft Lewis bases
Class (a)– hard acids
Class (b)– soft acids
Hard and Soft Acids and Bases (HSAB)
Let A be a Lewis acid, and B a base
Measure log K for the reaction
A + B  AB
If for B = halide, the order of
log K is
I– < Br– < Cl– < F–
then A is called a hard acid
If for B = halide, the order of
log K is
I– > Br– > Cl– > F–
then A is called a soft acid
Hard metal ions form their most stable complexes with
Hard Bases
Hard Bases: contain the smaller electronegative atoms,
especially O, N, F and Cl.
The bonding between a Hard Lewis Acid and a
Hard Lewis Base is predominantly ionic
Soft metal ions form their most stable complexes with
Soft Bases
Soft Bases: contain the larger, more polarisable and less
electronegative atoms, especially S, Se, P, C and As.
The bonding between a Soft Lewis Acid and a
Soft Lewis Base is predominantly covalent
Chapter 6 p183
Chapter 6 p184
The smaller drop in energy in
the hard-hard case does not
indicate small interaction.
The hard-hard interaction
depends on a longer range
electrostatic force, and this
interaction can be quite
strong. Many comparisons of
hard-hard and soft-soft
interactions indicates that the
hard-hard combination is
stronger and is the primary
driving force for the reaction.
Quantitative measure
Absolute hardness  = (I –A)/2
Mulliken’s definition of electronegativity
 = (I + A)/2
Softness  = 1/
Chapter 6 p189
Chapter 6 p189
Chapter 6 p191
Measurement of Acid-Base Interactions
1. Change in boiling points
2. Direct calorimetric methods or temperature dependent
of equivalent constants can be used to measure
enthalpies and entropies
3. Gas phase measurements of the formation of
protonated species
4. IR
5. NMR
6. UV-vis
Thermodynamic measurements
The enthalpy and entropy of ionization of a weak acid HA can be found by
measuring (1) the enthalpy of reaction with NaOH, (2) the enthalpy of reaction
of a strong acid (HCl) with NaOH and the equivalent constant for dissociation
of the acid.
H1o
(1) HA + OH-  A- + H2O
H2o
(2) H3O+ + OH-  2H2O
Ka
(3) HA + H2O ⇋ H3O+ A-
Ka
H3o = H1o - H2o
S3o = S1o - S2o
G3o = -RTlnKa = H3o - T S3o
Ln Ka = - H3o/RT + S3o /R
H3o
On a plot of Ka vs 1/T, the
slope is - H3o/R and the
intercept is S3o/R
Chapter 6 p193
Proton Affinity
BH+(g) B(g) + H+
proton affinity = H
A large proton affinity means it is difficult yo remove the hydrogen ion;
this means that B is a strong base and BH+ is a weak acid .
1. The alkali metal hydroxide, which are of equal basicity in aqueous
solution have gas phase basicities in the order LiOH < NaOH < KOH
<CsOH. This order matches the increase in the electron-releasing ability
of the cation in these hydroxides.
2. Pyridine and analine are stronger base the ammonia in the gas phase,
but they are weaker than ammonia in aqueous solution., presumably
because the interaction of the ammonium ion with water is more favorable
than the interaction with pyridinium or anilinium ions,
In binary acids, such as the hydrogen halides, the strength of the acid is
determined by the strength of the H–X bond. For a series such as the hydrogen
halides, the strength of the H–X bond decreases as the size of X increases. In
terms of acidic strength, HF < HCl < HBr < HI
As we move from left to right across a row in the periodic table, there is less
change in bond strength. In this case, what determines acid strength is the polarity
of the H- X bond. The electronegativity of elements increases from left to right
across a period in the periodic table. As the electronegativity of X increases, the
polarity of the H- X bond increases, increasing acidity. Acidity of the second row
hydrides varies as
CH4 < NH3 < H2O < HF
Chapter 6 p195
Inductive effects
An atom like fluorine which can pull the bonding pair away from
the atom it is attached to is said to have a negative inductive effect.
Most atoms that you will come across have a negative inductive
effect when they are attached to a carbon atom, because they are
mostly more electronegative than carbon.
You will come across some groups of atoms which have a slight
positive inductive effect - they "push" electrons towards the carbon
they are attached to, making it slightly negative.
Inductive effects are sometimes given symbols: -I (a negative
inductive effect) and +I (a positive inductive effect).
Inductive effects (electron releasing and withdrawing)
PF3 is a much weaker base than PH3
Base strength NMe3 > NHMe2 > NH2Me > NH3
But the acid strength BF3 < BCl3 ≤ BBr3
BF3 and BCl3 have significant  bonding that increase the
electron density on B
Inductive Effect
Effect
Acid
Base
Electron
strengthen weaken
withdrawing by
inductive effect
(-I)
electron
releasing by
inductive effect
(+I)
Examples
-NO2
and
X
O
weaken
and
strengthen
O
R
Strength of Oxyacids
Many bases contain the OH- ion, but O–H groups are found in acids as
well. Whether an OH compound is a base or an acid depends on
whether the OH groups are combined with a metal or a nonmetal. For
instance, sodium is a metal; NaOH is a base. Chlorine is a nonmetal;
HClO is an acid. Acids that contain one or more O–H bonds are called
oxyacids.
The strength of an oxyacid depends on the electronegativity of the
central nonmetal to which the OH groups are bound and on the number
of oxygen atoms bound to the central nonmetal atom. For a series of
oxyacids with the same number of oxygen atoms, the acidity increases
with the electronegativity of the nonmetal. The table below gives such a
series and the corresponding Ka values.
For a series of acids with the same central nonmetal atom, the
acidity increases with the number of oxygen atoms bound to the
central atom. (This also relates increasing acidity to increasing
oxidation number on the central atom.)
Oxyacid Strength
• More electronegative E, more ionic O-H bond,
stronger acid
– H2SO4 > H3PO4
– HNO3 > H2CO3
H-O-E-
• Less electronegative E, O-H more covalent, E-O
more ionic and more likely to beak in water
– Which bond do you expect to ionize in NaOH in water?
Oxyacid Strength
• More oxygens on central atom
– Withdraw e- from O-E bond, making H-O more
ionic
– Negative charge spread out over larger
– anion, reducing charge density, reducing
O
attraction for H+
• More oxygens, stronger acid
H-O-EO O
Oxyacid Strength
• H2SO4 > H2SO3
• HNO3 > HNO2
• HClO4 > HClO3 > HClO2 > HClO
Acidity of cations in aqueous solution
Many positive ions exhibit acidic behavior in solution.
In general, metal ions with large charges and smaller radii are
stronger acid.
Solubility of the metal hydroxide is a measure of
cation acidity. The stronger the acid, the less soluble
the hydroxide
Steric effect
H. C. Brown: Molecules have F (front) strain and B strain (back)
strain depending on whether the bulky groups interfere directly with
the approach of an acid and a base to each other. He also called
effects from electronic differences within similar molecules I
(internal) strain.
Gas phase measurements of proton affinity: Me3N > Me2NH > MeNH2 > NH3, on the
basis of electron donation by the methyl groups and resulting increased electron
density and basicity of the nitrogen.
When larger acid are used, the order changes: (B strain)
Solvation and acid-base strength
In aqueous solution: Basicity : Me2NH > MeNH2 > Me3N > NH3
Et2NH > EtNH2 = Et3N > NH3
Solvation energies for the reaction
RnH4-nN+(g) + H2O  RnH4-nN+(aq)
are in the order RNH3+ > R2NH2+ > R3NH+
Solvation is dependent on the number of H atoms available for Hbonding to water to form H-O---H-N H-bonds.. With fewer H atoms
for H-bonding, the more highly substituted molecules are less basic.
Nonaqueous solvents and acid-base strength
HOAc + H2O ⇋ H3O+ + OAc- (about 1.3 % in 0.1M solution )
HCl + H2O ⇋ H3O+ + Cl- ( 100 % in 0.1M solution)
NH3 + H2O ⇋ NH4+ + OH- (about 1.3 % in 0.1M solution )
Na2O + H2O ⇋ 2Na+ + 2OH- ( 100 % in 0.1M solution)
Water is amphoteric. The strongest acid possible in water is H3O+ and the
strongest base is OH-.
In glacial acetic acid solvent (100 % acetic acid) only the strongest acids can
force another H ion onto the acetic acid molecule, but acetic acid will react
readily with any base, forming the conjugate acid of the base and the acetic ion:
H2SO4 + HOAc ⇋ H2OAc+ + HSO4NH3 + HOAc ⇋ NH4+ + OAcThe strongest base possible in pure acetic acid is the acetate ion. Any stronger
base reacts with acetic acid solvent to form acetate ion.
OH- + HOAc-  H2O + OAc-
Leveling effect- in which the acids or bases are brought down to the limiting
conjugate acid or base of the solvent
Effect by which all acids stronger than the acid that is characteristic of the
solvent react with solvent to produce that acid; similar statement applies to
bases. The strongest acid (base) that can exist in a given solvent is the acid (base)
characteristic of the solvent.
In acetic acid, the acid strength is in the order HClO4 > HCl > H2SO4 > HNO3
Superacids – Acid solutions more acidic than sulfuric acid are called superacid.
For which George Olah won the Nobel prize in Chemistry in 1994. The acidity
is measured by the Hammett acidity functions Ho = pKBH+ - log[BH+]/[B].
Where B and BH+ are a nitroaniline indicator and its conjugate acid. The
stronger the acid, the more negative its Ho value.