Buffers
Year 12 Chemistry
What is a buffer?
• A buffer is a solution that resists changes in pH
when small amounts of acid or base are added to
it
• There are 2 types:
▫ Acidic
▫ Alkaline
Acidic buffers
An acidic buffer has a pH less than 7
It is made from a equal molar concentrations of a
weak acid and it’s conjugate base
Example: ethanoic acid and ethanoate ion
CH3COOH  CH3COO- + H+
(weak acid)
(conj. base)
Acidic buffers
Take a 0.1M solution of ethanoic acid:
CH3COOH 
At equil:
(0.09583M)
What’s the pH of this solution?
CH3COO- + H+
(0.00417M)
(0.00417M)
-log [H+] = 2.38
Now, say you add HCl to this equilibrium. The acetate ion is the only
species available to reduce the added H+ and these are very low in
concentration, so the pH will drop dramatically.
What can you do to reduce this change? Add acetate ions (sodium acetate)
Acidic buffers – adding acid
CH3COOH 
-
CH3COO
+ H
+
If we buffer the solution by adding 1M sodium acetate, what will happen?
This will have the following effect:
1. Increase the amount of acetate ions, shifting the equilibrium to the left
2. Increase the pH to 4.76
3. When adding H+ ions, the extra acetate ions will react to reduce the [H+]
4. The solution resists changes to pH, meaning it is buffered
Acidic buffers – adding base
Add OH
-
CH3COOH 
-
CH3COO
+
+ H
Adding base results in more OH- being added to the equilibrium. This
extra amount of ions is removed by 2 processes. What are these?
1.
2.
CH3COOH + OH-  CH3COO- + H2O (ethanoic acid is removed
by reaction with undissociated ethanoic acid.
The small amount of H+ ions in the above reaction will also remove
OH- ions, shifting the equilibrium to the right to make more H+ ions
that further remove OH-.
pH of buffers depends on
concentration of conjugate pair
pH
vol. of 0.1M
acetic acid (ml)
vol. of 0.1M
sodium acetate (ml)
3
982.3
17.7
4
847.0
153.0
5
357.0
653.0
6
52.2
947.8
Note: you can also get various pH buffers by changing the acid/base pair.
Alkaline buffers
An alkaline buffer has a pH greater than 7
It is made from a equal molar concentrations of a
weak base and it’s conjugate acid
Example: ammonia and ammonium ion
+
NH3 + H2O   NH4 + OH
(weak base)
(conj. acid)
-
Alkaline buffers
+
-
NH3 + H2O   NH4 + OH
(weak base)
(conj. acid)
Due to the weak nature of ammonia, this equilibrium will be well to
the left. We can create a buffered solution by adding ammonium
chloride.
What will this do to the equilibrium?
Addition of ammonium ions will shift the equilibrium even further to
the left. The pH of this solution would be 9.25 for a 1M solution.
Alkaline buffers – adding acid
+
-
NH3 + H2O   NH4 + OH
What will happen if you add acid to this solution?
Two processes:
1.
NH3 + H+   NH4+ (removal by reaction with ammonia to
produce more ammonium ion)
2.
NH3 + H2O   NH4+ + OH- (removal by reaction with OHto produce water
Combines with
H+ to form water
Alkaline buffers – adding base
+
NH3 + H2O   NH4 + OH
-
What will happen to the equilibrium if you add base?
Adding base effectively adds OH-. This means:
The ammonium ion reacts with the OH- to shift the equilibrium to the
left, consuming most of the OH- ions.
NH4+ + OH-   NH3 + H2O
Summary
• Buffer solutions resist changes in pH when acids
and alkalis are added
• Buffers generally contain:
▫ Sufficient concentrations of a weak acid and it’s
conjugate base OR weak base and it’s conjugate
acid
• The pH of buffer solutions depend on the
concentrations and type of conjugate acid/base
pairs that are used.
Buffer applications
Blood
The pH of human blood must be maintained at 7.4 or serious health
consequences can result. The buffering system that maintains this pH is
H2CO3/HCO3-:
H2CO3  H+ + HCO3-
If blood increases in acidity, the
additional H+ ions will react
with the bicarbonate ions
If blood increases in alkalinity, the
H+ ions in the equilibrium will
react and shift to the right
Other buffers in the blood also help
to maintain the required pH.
Haemoglobin is itself a weak
base that helps in this process
Buffer applications
Body cells
The phosphate buffer system operates in the internal fluid of all cells. This
buffer system consists of dihydrogen phosphate ions (H2PO4-) as hydrogenion donor (acid) and hydrogen phosphate ions (HPO42-) as hydrogen-ion
acceptor (base).
These two ions are in equilibrium with each other as indicated by the chemical
equation below.
H2PO4-(aq) H+(aq) + HPO42-(aq)
Buffer applications
Swimming pools
The H2CO3/HCO3- buffer system that is used in the blood is also used
in swimming pools.
Sodium hydrogen carbonate is often added to swimming pools if it is
proving difficult to maintain the pH between 7.2 – 7.4.
The measurement “total alkalinity” in pools is a measure of OH- and
HCO3-. This is effectively a measure of the buffering capacity of the
water. If it is too low, then bicarbonate must be added.