Acid-Base Theories

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Acid-Base Theories                             
                            
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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)
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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 Cl- Acid 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.
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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 BrF4- SbF5
BrF3 ? BrF2 SbF6- SbF5 is an acid KF BrF3
? BrF4- K KF is a base
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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 ? F3BNH3
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Chapter 6 p166
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Chapter 6 p169
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Which theories can explain the following?
HI H2O ?     H3O I-  
HI    NH3 ? NH4 I-
I2 NH3 ? NH3I I-
I2    Cl-? ICl I-  
X- Y ? YX  
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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
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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 ?
H3NAgNH3
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Frontier orbitals and acid-base reactions
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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
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HOMO-LUMO interactions

• 2H2O Ca ? Ca2 2OH- H2
• water as oxidant

2. nH2O Cl-? Cl(H2O)n- water
as acid
3. 6H2O Mg2 ? Mg(H2O)62
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 bases
HOMO and the acids LUMO, the stronger the
interaction.
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Hydrogen bonding
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
4 nodes
3 nodes
2 nodes
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For unsymmetrical H-bonding B HA ? BHA, the
pattern is similar
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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
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• 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
• Solvation of the ions is a factor in these
  reactions, with fluoride ion being much strongly
  solvated than the other anions.
• 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.

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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 hard-hard
interactions.
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Color and low solubility typically go with
soft-soft interactions colorless compounds and
high solubility generally go with hard-hard
interactions, although some hard-hard combination
have low solubilities. LiBrgt LiCl gt LiI gt
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.
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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)
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• Four rules can be summarized
• For a given cation, covalent character increases
  with increase in size of the anion.
• For a given anion, covalent character increases
  with decrease in size of the cation.
• Covalent character increase with increasing
  charge on either ion.
• 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 O2-
Q2. Fe(OH)3 is much less soluble than Fe(OH)2
A2. Rule 3 Fe3 has a larger charge than Fe2
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• These rules are helpful in predicting behavior of
  specific cation-anion interaction, but not enough
• Li series does not fit
• Solubility MgCO3 gt CaCO3 gtSrCO3 gtBaCO3
• 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.

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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
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The class (b) ions form halides whose solubility
is in the order F- gt Cl- gt Br- gt 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
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Pearsons 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
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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 gt
Br gt Cl gt F then A is called a soft acid
If for B halide, the order of log K is I lt
Br lt Cl lt F then A is called a hard acid
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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
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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
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Chapter 6 p183
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Chapter 6 p184
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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.
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Quantitative measure Absolute hardness ? (I
A)/2 Mullikens definition of
electronegativity
? (I A)/2
Softness ? 1/?
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Chapter 6 p189
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Chapter 6 p189
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Chapter 6 p191
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• Measurement of Acid-Base Interactions
• Change in boiling points
• Direct calorimetric methods or temperature
  dependent of equivalent constants can be used to
  measure enthalpies and entropies
• Gas phase measurements of the formation of
  protonated species
• IR
• NMR
• UV-vis

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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. (1) HA
OH- ? A- H2O
?H1o (2) H3O OH- ? 2H2O
?H2o Ka (3) HA
H2O ? H3O A- Ka
?H3o ?H3o ?H1o - ?H2o ?S3o ?S1o - ?S2o ?G3o
-RTlnKa ?H3o - T ?S3o Ln Ka - ?H3o/RT
?S3o /R
On a plot of Ka vs 1/T, the slope is - ?H3o/R and
the intercept is ?S3o/R
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Chapter 6 p193
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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 lt NaOH lt KOH ltCsOH.
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,
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In binary acids, such as the hydrogen halides,
the strength of the acid is determined by the
strength of the HX bond. For a series such as
the hydrogen halides, the strength of the HX
bond decreases as the size of X increases. In
terms of acidic strength, HF lt HCl lt HBr lt 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 lt NH3 lt H2O lt HF
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Chapter 6 p195
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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).
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Inductive effects (electron releasing and
withdrawing) PF3 is a much weaker base than
PH3 Base strength NMe3 gt NHMe2 gt NH2Me gt NH3 But
the acid strength BF3 lt BCl3 BBr3 BF3 and BCl3
have significant ? bonding that increase the
electron density on B
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Inductive Effect
Effect Acid Base Examples
weaken
strengthen
Electron withdrawing by inductive effect (-I)
electron
releasing by inductive effect (I)
weaken
strengthen
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Strength of Oxyacids
 Many bases contain the OH- ion, but OH 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 OH 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.
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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.)
                                               
                                                  
                                       
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Oxyacid Strength

• More electronegative E, more ionic O-H bond,
  stronger acid
• H2SO4 gt H3PO4
• HNO3 gt H2CO3
• 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?

H-O-E-
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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
  attraction for H
• More oxygens, stronger acid

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Oxyacid Strength

• H2SO4 gt H2SO3
• HNO3 gt HNO2
• HClO4 gt HClO3 gt HClO2 gt HClO

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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.
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Solubility of the metal hydroxide is a measure of
cation acidity. The stronger the acid, the less
soluble the hydroxide
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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.
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Gas phase measurements of proton affinity Me3N gt
Me2NH gt MeNH2 gt 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)
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Solvation and acid-base strength In aqueous
solution Basicity Me2NH gt MeNH2 gt Me3N gt NH3

Et2NH gt EtNH2 Et3N gt NH3 Solvation energies
for the reaction RnH4-nN(g) H2O ?
RnH4-nN(aq) are in the order RNH3 gt R2NH2 gt
R3NH Solvation is dependent on the number of H
atoms available for H-bonding 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.
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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
HSO4- NH3 HOAc ? NH4 OAc- The 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-
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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 gt HCl gt H2SO4 gt HNO3
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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 - logBH/B.
Where B and BH are a nitroaniline indicator and
its conjugate acid. The stronger the acid, the
more negative its Ho value.