Arrhenius Definition
An acid is a substance that increases the
hydrogen (hydronium) concentration in a
water solution.
 HCl(aq)
H+(aq) + Cl-(aq)
 HCl(aq) + H2O(l)
H3O+(aq) + Cl-(aq)
Either equation is acceptable and H+(aq) or
H3O+(aq) is a hydrated proton.
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H+ is very strongly hydrated in water because
of its small size and high positive charge
density.
 H+(aq) + H2O(l)
H3O+(aq)
Arrhenius definitions are limited to aqueous
solutions.
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A base is a substance that increases the
hydroxide ions in a water solution.
 NaOH(aq)
Na+(aq) + OH-(aq)
Remember that Arrhenius definitions are
limited to aqueous solutions.
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Bronsted-Lowry Theory (BLT)
An acid is a molecule or ion that donates a
proton.
 HCl(aq)
H+(aq) + Cl-(aq)
A base is a molecule or ion that accepts a
proton.
 NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
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The BLT of an acid and a base are not limited
to aqueous solutions.
When working with BLT, it is common to use
the terms conjugate acid and conjugate base.
 NH3(aq) + H2O(l)
base
acid
NH4+(aq) + OH-(aq)
conjugate conjugate
acid
base
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H2SO4(aq) + H2O(l)
acid
base
H3O+(aq) + HSO4-(aq)
conjugate conjugate
acid
base
HSO4-(aq) + H2O(l)
acid
base
H3O+(aq) + SO42-(aq)
conjugate conjugate
acid
base
HSO4- is called amphoteric or amphiprotic
because it can act as either a BLT acid or a
BLT base depending on its chemical
environment.
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Here HSO4- is acting as the conjugate base of
H2SO4.
H2SO4(aq) + H2O(l)
H3O+(aq) + HSO4-(aq)
Here HSO4- is acting as the conjugate acid of
SO42-.
HSO4-(aq) + H2O(l)
H3O+(aq) + SO42-(aq)
Each acid has one more proton than its
conjugate base.
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Each base has one less proton than its
conjugate acid.
Two important points to remember:
 The stronger the acid, the weaker its
conjugate base.
 The stronger the base, the weaker its
conjugate acid.
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Acid-Base Reactions
An acid-base reaction always proceeds
toward the weaker acid and weaker base.
HClO4(aq) + H2O(l)
stronger
acid
stronger
base
H3O+(aq) + ClO4-(aq)
weaker
acid
weaker
base
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When analyzing an acid-base reaction,
remember that you can’t have it both ways.
ClO4- is too weak a base to compete with the
stronger base, H2O, to acquire the proton.
 How do you know that ClO4- is such a
weak base?
Because HClO4 is one of the six strong
acids.
When you have a strong acid such as
HClO4, 100% ionization is assumed.
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If a molecule wants to completely ionize, why
would its anion want to undergo hydrolysis?
Similarly, H3O+ is too weak an acid to
compete with the stronger acid, HClO4, to
donate a proton.
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Lewis Acids and Bases
An acid is a substance that accepts an
electron pair.
 Al3+(aq) + 6H2O(l)
Al(H2O)63+(aq)
The Al3+ cation has the empty orbitals 3s, 3px,
3py, 3pz, as well as the size to accommodate
d-orbitals.
Also, the Al3+ has a large positive charge
density resulting in its interaction with water
molecules.
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The Al3+ cation acts as a Lewis acid and the
water with its two unshared pair of electrons
acts as a Lewis base (an electron pair donor).
The hydrated Al3+ cation, Al(H2O)63+, can now
behave as an Arrhenius acid or a
Bronsted-Lowry acid.
Al(H2O)63+(aq)
H+(aq) + Al(H2O)5(OH)2+(aq)
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A base is a substance that donates an
electron pair.
 Zn2+(aq) + 4OH-(aq)
Zn(OH)42-(aq)
The Lewis definition of acids and bases
expands the number of species that can be
acids.
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Strength of Binary Acids
The H – X bond strength is the most important
factor to consider when determining acid
strength in a group or family.
Consider the following bond enthalpies:
H–F
 H – Cl
 H – Br
H–I
567 kJ mol-1
431 kJ mol-1
366 kJ mol-1
299 kJ mol-1
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The bond enthalpies from the previous slide
indicate that the strength of the H – X bond
decreases as the atomic radii of the halogen
increases.
 Longer bonds are generally weaker or
less stable than shorter bonds.
Similarly, H2S is a stronger acid than H2O,
Ka(H2S) > Ka(H2O).
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The H – X bond polarity is the most important
factor to consider when determining acid
strength in a period or series.
 Because electronegativity increases
from left to right in a period, the acid
strength also increases proceeding from
left to right.
 The acid strength increases from left to
right in a period.
 Ka(HF) > Ka(H2O) > Ka(NH3) > Ka(CH4)
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Strength of Oxyacids
When comparing oxyacids, there are
additional factors to consider.
│
As the electronegativity of element
– X – O – H X increases, the stronger the acid.
│
When the electronegativity of X increases, the
polarizability of the O – H bond increases.
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As more O terminal atoms are added to the
central atom, X, the more the electron density
is pulled from the O – H bond.
By adding more electronegative atoms to X,
the acid strength is increased.
│
│
Cl – O – H < Cl – O – H < Cl – O – H <
│
│
O – Cl – O – H
│
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By adding more oxygens (the second most
electronegative element) to the central atom,
Cl, the electron density shifts more towards
the oxygens, making the O – H bond more
polarizable.
For oxoacids with the same number of O – H
bonds and the same number of oxygen atoms,
the acid strength will increase with an
increase of electronegativity of the central
atom.
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│
│
│
O – Cl – O – H > O – Br – O – H > O – I – O – H
│
│
│
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Oxoacid Wrap Up
For an oxoacid, the H atom that ionizes is
bonded to an O atom which in turn is bonded
to a nonmetal atom.
The strength of any acid depends on how
easily the O – H bond is broken.
 One deciding factor is the oxidation of
the central atom.
 The higher the oxidation number the
stronger the acid.
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To increase the ionization, the electron
density surrounding the O atom which is
bonded to the ionizable H, should be as low
as possible.
To decrease the electron density around the
O atom:
 Make the central atom more
electronegative.
 Add more O atoms to the central atom.
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A second deciding factor is the
electronegativity of the central atom.
 The more electronegative the central
atom, the stronger the acid.
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Extent of Hydrolysis
There are six strong acids that completely
ionize in water.
 HCl, HBr, HI, HNO3, HClO4, H2SO4
When representing the ionization of these
acids, a single arrow is used and 100%
ionization is assumed.
 HClO4(aq)
H+(aq) + ClO4-(aq)
The Ka of these acids is assumed to be
infinite.
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Most acids are weak and only partially ionize
in water.
 HC2H3O2 (aq) + 2H2O(l)
H3O+(aq) +
C2H3O2-(aq)
Or alternatively
 HC2H3O2(aq)
H+(aq) + C2H3O2-(aq)
The Ka of these acids is small and can be
looked up for each acid.
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Also note that with weak acids, a double
arrow is used and a dynamic equilibrium
results.
The most commonly encountered bases that
completely dissociate in water are:
 LiOH, NaOH, KOH, RbOH, CsOH,
Ca(OH)2, Sr(OH)2, Ba(OH)2
When representing the dissociation of these
bases, a single arrow is used and 100%
dissociation is assumed.
 LiOH(aq)
Li+(aq) + OH-(aq)
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The Kb of these bases is assumed to be
infinite.
Most bases are weak and only partial
ionization takes place.
 NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
The Kb of these bases is small and can be
looked up for each base.
Note that in the case of a weak base, water
must be explicitly written as a reactant unlike
the case of a weak acid.
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