Water ionizes
H3O+
H2O
OH-
pure water
(hydronium ion)
(hydroxide ion)
Water is an acid and a base
Hydronium ion
H-donor
Makes it an acid based on
the Bronsted-Lowry definition
Hydroxide ion
H-acceptor
Makes it a base based on
the Bronsted-Lowry definition
Remember, we stated that water can act as a solvent and as a reactant
This definition explains why we classified some aa as acids or bases
If water is an acid and a base, why don’t we get acid burns in the bathtub?
Actual:
Shorthand:
The concentrations of reactants and products are at equilibrium
Note: equil, not equal!!!
At equilibrium, things are NOT equal!!
Equilibrium indicates that things are not changing:
Amounts of reactants are not changing
Amounts of products are not changing
The reactions is still proceeding at the same rate
Chemistry C483
Fall 2009 Prof Jill Paterson
5-1
The concentrations of reactants and products are at equilibrium
At equilibrium, the ratio of products and reactants IS the equilibrium constant:
Keq =
Where Keq is expressed with the units Molar (M; moles/L)
Constants:
Concentration of pure H2O is 55 .5 M
Keq of pure H2O is 1.8 x 10-16 M
What is the significance of 1.0 x 10-14 M2 ?
Basis for the pH scale!
pH measures acidity of solution
pH=-log[H+]=log(1/[H+]) (because initial #s were too small)
Scale goes from 1-14
Kw = 1.0 x 10-14 M2 = [H+][OH-]
We can change the pH by changing the [H+] or [OH-]
Note: If we add H+, the [H+] increases, so [OH-] must
decrease because [H+]*[OH-] = 1.0 x 10-14 M2
What is the pH of a solution of 0.1 M HCl (a strong acid) in H2O?
Strong acids completely dissociate in water:
What is the pH of a solution of 0.1 M NaOH (a strong base) in H2O?
Strong bases completely dissociate in water:
Calculating the pH of a weak acid or weak base solution is not as straightforward…
What is the pH of a 0.1 M solution of formic acid?
Since acetic acid is a weak acid, it does not dissociate completely, so…
Therefore, even if we know the [CH3COOH], we do not know the concentration of [H+]
Many of the dissociation constants of common acids are known (See Table 2-4 of your text)
When dealing with weak bases we need to consider its conjugate acid
Conjugate acids/conjugate bases
Strong acids have weak conjugate bases
Weak bases have strong conjugate acids
Strong bases have weak conjugate acids
Weak acids have strong conjugate bases
All acids have a conjugate base (complementary)
All bases have a conjugate acid (complementary)
You should be able to identify an acid’s conjugate base (&vice versa) and
a base’s conjugate acid (& vice versa)
What is the pH of a 0.1 M solution of ammonia (weak base)?
The pKa of an acid
• pKa is the measure of the acidic strength of a compound
(Remember pH measures acidity of a solution)
• pKas are lower for stronger acids and higher for weaker acids (stronger bases). This is the opposite of
Ka where higher number indicates a stronger acid.
• The pKa is the point at which 50% of the compound is protonated, and 50% of the compound is
deprotonated
• A pKa can be determined easily experimentally, so is generally the value used in biochemistry
• The pKa is the inflection point of a titration curve
Titrations
How much base does it take to deprotonate CH3COOH?
Rxn:
CH3COOH + OH-
CH3COO- + H2O
1. At start of reaction, we have no OH- (0 equivalents of OH-). 100% of acetic acid is
protonated.
2. Because there is only 1 ionizable group on acetic acid (the carboxyl group), we will need only
1 equivalent of OH-.
3. The pH when exactly 50% of the acetic acid molecules are still protonated and 50% have
been deprotonated is the pKa.
Other acids have multiple ionizable groups
Termed polyprotic.
These acids release H+ sequentially, not randomly, because of differeing pKas
Knowing the pKa tells us about protonation states at other pHs
pH
[A-]
[HA]
ratio
pKa + 2
pKa + 1
pKa
pKa – 1
pKa – 2
We can determine the pH of a solution using a known pKa
Henderson-Hasselbalch equation:
pH = pKa + log
[A-]
[HA]
[A-] = [H+] in dissociation reactions
We use this to calculation the pH of weak acids and weak bases because these do not completely
dissociation and the pKa value allows us to account for this.
What is the pH of a 0.1 M solution of formic acid? (rework this problem with the H-H equation)
Chemistry C483
Fall 2009
Prof Jill Paterson
4-7
H-H equation can be used to make acid & base solutions
Problem: How many mL of 0.1 M formic acid and 0.1 M sodium formate are needed to make 100
mL of a 0.1 M solution having a pH of 4.8?
Buffers
When a small amount of acid or base is added to a solution, if the pH remains the same, the
solution is said to be “buffered”
Buffer range: linear portion of curve
Best buffering occurs when _________
General rule is buffers are good ____ (pKa is 6, range of buffer is 5-7)
Acid/conjugate base pairing makes a good buffer!
Buffers are necessary in living systems
• All biological reactions require a specific pH because biological molecules are charged differently
at different pHs
• Role of a buffer is to maintain equilibrium
One of the most important buffers is carbonic acid/bicarbonate pair found in blood and other fluids
physiological pH of blood is 7.4
Large amounts of acid can be added and the pH only drops to 7.2!!
How is this possible?
Chemistry C483
Fall 2009
Prof Jill Paterson
4-8
Blood buffering system
Blood is buffered by a network of reactions involving :
CO2
H2CO3 (carbonic acid)
HCO3- (bicarbonate, a conjugate base)
CO32- (carbonate, which is not relevant
because its pKa is 10.2)
Net reaction : CO2 + H2O
H+
pH
H+ + HCO3-
CO2(aq)
H2CO3
HCO3-
CO2 (g)
action
Increase
Decreases
Thus the system is easily regulated by breathing! (and some by kidneys (HCO3-)
This is needed- look at the pKa of bicarbonate.
Buffer range is 5.4-7.4
Blood is at 7.4- RIGHT AT THE CUSP OF THE BUFFER ZONE!!!
Chemistry C483
Fall 2009 Prof Jill Paterson
4-9