Acids and Bases
Ionization of water
Neutral water has a tendency to ionize
H 2O
H+ + OH-
The free proton is associated with a water molecule to form the
hydronium ion
H 3O+
2
Proton Jumping/Proton Hopping
H3O+ : 362.4 x 10-5 cm2•V-1•s-1
51.9 x 10-5 cm2•V-1•s-1
H+ appears to move extremely quickly in solution compared to K+ or
Na+ movement
Hydronium ion migration; hops by switching partners at 1012 per
second
Na+:
3
Dissociation of H2O
H 2O H OH
Water also dissociates
[H2O] = 55.5
[ H ][OH ]
K
[ H 2O]
Kw [ H ][OH ] 10
4
14
M
2
Ionization constant for water
What about the water
The concentration of H2O remains almost unchanged especially
in dilute acid solutions.
What is the concentration of H2O?
Remember the definition:
Moles per liter
1 mole of H2O = 18 g = 18 ml
1000 g/liter
5
1000 g
55.5M
18 g / mol
Equilibrium expression
H 2O H OH
Described by:
Keq = [H+] [OH-]/[H2O]
Keq = [H+] [OH-]/55.5M
Keq x 55.5M= [H+] [OH-] = Kw
Where Keq is the dissociation constant
Keq for pure water determined experimentally to be 1.8 x 10-16 M at 25ºC
Concentration of pure H2O = 55.5M (weight of water in 1 L (1000 g)
divided by mw of 18)
[H+] [OH-] = Keq x 55.5M
[H+] [OH-] = 1.8 x 10-16 x 55.5 = 1.0 X 10-14
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Equilibrium expression
[H+] = [OH-]
If X = [H+] = [OH-]
X2 = Kw = 1 x 10-14
1 x 10-7 = [H+] = [OH-]
This is the basis of the pH scale!
These numbers are very small and difficult to work with, so in
1909 Soren Sorenson introduced
the term pH to more conveniently express [H+].
Defined pH as the negative logarithm of the hydrogen ion
concentration:
pH = -log [H+] = log 1/[H+]
“p” is an operator – means to “take the negative log of”
Example: pOH = -log[OH-]; pH of pure water? [H+] = 1 x 10-7,
pH = -log (1 x 10-7) = 7
In a neutral solution:
7
pH Scale
Defined pH as the negative logarithm of the hydrogen ion
concentration:
pH = -log [H+] = log 1/[H+]
“p” is an operator – means to “take the negative log of”
Example: pOH = -log[OH-]; pH of pure water? [H+] = 1 x 10-7,
pH = -log (1 x 10-7) = 7
pH + pOH = 14
The pH scale ranges from 0 to 14
pH scale is LOG BASED! Used to keep track of large changes
important to acids and bases
Important to remember that the scale is exponential
One pH unit = 10 times more acidic or basic
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pH Scale
For pure water (neutral)
[H+] = [OH-] = (Kw)1/2 = 10-7 M
Acidic if [H+] > 10-7 M
9
Basic if [H+] < 10-7 M
Acids and bases
Applies to other acids and bases, not just water
Acid = Releases proton in water (proton donor)
Base = Accepts proton in water (proton acceptor)
Water can act as an acid and a base = amphiprotic
Strength of an acid is defined as its tendency to release a
proton (dissociate)
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Acid/conjugate base pairs
HA + H2O
HA
A- + H3O+
A- + H +
HA = acid ( donates H+)(Bronstad Acid)
A- = Conjugate base (accepts H+)(Bronstad Base)
Ka = [H+][A-]
[HA]
pKa = - log Ka
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Ka & pKa value describe tendency to loose H+
large Ka = stronger acid
small Ka = weaker acid
Acid/conjugate base pairs
Conjugate Base: base formed by the removal of a proton from an acid
The acid and conjugate base are complementary species – every acid
has a conjugate base
Should be able to identify acids and their conjugate bases:
Acetic acid/acetate pair:
12
pKa
Just as for water, we can write an equilibrium constant for the dissociation
of the acid (Ka)
We want Ka in convenient terms. Apply “p” rule and take the negative log
of Ka to get pKa.
pKa is a QUANTITATIVE measure of acid strength.
Smaller pKa
Stronger acid
Larger pKa
Weaker acid (stronger base)
Opposite of Ka where a LARGE number indicates strong acid
Large Ka means mostly dissociated into H+ and A-, not much HA left.
Ka large = pKa small
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Weak Acids
Weak acids do not completely dissociate:
They form an equilibrium:
If we ADD more H+, the equilibrium shifts to
form more HA using up A- that is present.
14
Relationship between pH and [H+] / [OH-]
concentration
15
16
Henderson - Hasselbalch equation
From
Rearrange
Take (-)Log of each
[ H ][ A ]
K
[ HA]
[ HA]
[H ] K
[A ]
[ A ]
pH log K log
[ HA]
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[A ]
pH pK log
[ HA]
What we learn from this class?
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