Unit 1 Mod 2 Buffers
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BUFFERS
A buffer solution is one which resists changes in pH when small quantities of an
acid or an alkali are added to it.
Acidic buffer solutions
An acidic buffer solution is simply one which has a pH less than 7. Acidic buffer
solutions are commonly made from a weak acid and a salt of the weak acid
usually the sodium salt.
CH3COOH
CH3COO- + H+
CH3COONa  CH3COO- + Na+
What happens when an acid is added to an acidic buffer solution
There would be an INCREASE in the concentration of hydrogen ions.
Via Le Chatelier’s Principle the system would then try to DECREASE the concentration
of H+ ions. Hydrogen ions would then combine with the ethanoate ions to make ethanoic
acid. This means the equilibrium reaction shifts to the LEFT.
NOTE
Although the reaction is reversible, since the ethanoic acid is a weak acid, most of the
hydrogen ions added from the acid are removed in this way. The pH of the buffer
solution will still drop but by a very small amount.
What happens when an alkali is added to an acidic buffer solution
The excess hydroxide ions will combine with the hydrogen ions and be removed from the
buffer. In essence by adding hydroxide ions, this causes a DECREASE in the hydrogen
ion concentration.
The system would then try to INCREASE the concentration of H+ ions.
More ethanoic acid molecules would dissociate producing more hydrogen ions. This
means the equilibrium reaction shifts to the RIGHT.
NOTE
Most of the hydrogen ions removed by the hydroxide ions are replaced The pH of the
buffer solution will still rise but by a very small amount.
Alkaline buffer solutions
An alkaline buffer solution has a pH greater than 7. Alkaline buffer solutions are
commonly made from a weak base and one of its salts.
NH3 + H2O O
NH4+ + OHNH4Cl  NH4+ +
ClWhat happens when an acid is added to an alkaline buffer solution
The hydrogen ions will combine with the hydroxide ions and remove them from the
buffer. In essence by adding hydrogen ions, this causes a DECREASE in the hydroxide
ion concentration.
Unit 1 Mod 2 Buffers
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The system would then try to INCREASE the concentration of OH- ions.
More ammonia molecules would dissociate producing more hydroxide ions. This means
the equilibrium reaction shifts to the RIGHT.
NOTE
Most of the hydroxide ions removed by the hydrogen ions are replaced The pH of the
buffer solution will still drop but by a very small amount.
Adding an alkali to an alkaline buffer solution
There would be an INCREASE in the concentration of hydroxide ions.
The system would then try to DECREASE the concentration of OH- ions.
Ammonium ions would then combine with the hydroxide ions to make ammonia. This
means the equilibrium reaction shifts to the LEFT.
NOTE
Although the reaction is reversible, since the ammonia is a weak base, most of the
hydroxide ions added are removed in this way. The pH of the buffer solution will still
rise but by a very small amount.
Checkpoint A
Tick the pairs of substances can act as buffers or cannot act as buffers. Of those pairs
which CAN act as buffers, write the letters “Ac” to represent acidic buffers or “Al” to
represent alkaline buffers,
a) HNO3 + NaOH b) CH3COOH + CH3COONa
c) H3PO4 + HCl d) NH3 + NH4Cl
How to calculate the pH of buffer solutions
Henderson-Hasselbalch equation
Where [A-] represents the MOLAR CONCENTRATION of the salt and [HA] represents
the MOLAR CONCENTRATION of the weak acid
Example 1
a) A 1.0 mol dm-3 buffer is prepared from 0.25 mol of ethanoic acid and 0.25 mol of
sodium ethanoate. pKa ethanoic acid = 4.7
pH = 4.7 + log (0.25/0.25) = 4.7 + log 1  = 4.7 + 0 = 4.7
b) Calculate the change in pH if 0.15 mol of OH- ions are added to the buffer. Note the
reaction of ethanoic acid and hydroxide ions react in a 1:1 ratio
CH3COOH + OH- 
CH3COO- + H2O
Therefore the # of moles of ethanoic acid would DECREASE by 0.15 and the # of moles
of ethanoate ions would INCREASE by 0.15.
Therefore the NEW molar concentrations are
[CH3COOH] = 0.25 – 0.15 = 0.1 and [CH3COO-] = 0.25 + 0.15 = 0.40
Hence the new pH = 4.7 + log (0.4/0.1)
= 4.7 + log 4 = 5.3
Unit 1 Mod 2 Buffers
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Example 2
A buffer solution was made by dissolving 10.0 grams of sodium ethanoate in 200.0 mL
of 1.00 M ethanoic acid. Calculate the pH of the ethanoic acid/sodium ethanoate buffer
solution. Ka acid is 1.7 x 10-5.
pKa = 4.7
[A-] = mass of salt / molar mass of salt = 10 / 82 = 0.122 x 5 = 0.611
[HA] = 1
Hence pH = 4.7 + log(0.61/1)
= 4.77 – 0.22
= 4.55
Checkpoint B
1. Calculate the pH of a buffer solution containing 0.12 mol dm-3 ethanoic acid and 0.34
mol dm-3 sodium ethanoate. The pKa of ethanoic acid is 4.75
2. 1 dm3 of a buffer contains 0.325 g of aspirin and 0.01M of the salt of aspirin
Ka aspirin = 3.3 x 10-4
a) Calculate the molar concentration of the aspirin (molar mass 180 g)
b) Calculate the pKa of aspirin ………………………………
c) Assuming that aspirin is a monobasic acid, write the Ka equilibrium
d) Using the answer from part (a) and the Ka value, calculate the [H+]
e) Calculate the pH of the buffer solution
The Carbonic-Acid-Bicarbonate Buffer in the Blood
By far the most important buffer for maintaining acid-base balance in the blood is the
carbonic-acid-bicarbonate buffer. The simultaneous equilibrium reactions of interest are
H+ (aq) + HCO3- (aq) ↔ H2CO3 ↔ H2O + CO2
Other buffers perform a more minor role like the phosphate buffer consists of phosphoric
acid (H3PO4) in equilibrium with dihydrogen phosphate ion (H2PO4-) and H+.
Haemoglobin also acts as a pH buffer in the blood.
Unit 1 Mod 2 Buffers
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How exercise affects the pH of blood
When we exercise, our heart rate etc increases and the body's metabolism becomes more
active, producing CO2 and H+ in the muscles. Eventually, with strenuous exercise, our
body's metabolism exceeds the oxygen supply and begins to use alternate biochemical
processes that do not require oxygen. The pH of the blood begins to decrease.
Due to Le Chatelier’s Principle, the addition of CO2 and H+ causes an increase in the
concentration of H2CO3 thereby decreasing the concentration of CO2 and H+. However
this occurs with mild exercise, with strenuous exercise, the buffer system cannot
adequately handle the drastic changes in pH that would occur. When there is
strenuous exercise, the lungs remove CO2, the kidneys remove HCO3- ions, but the high
heart rate can hinder CO2 removal.
Why would buffer systems in enzyme catalysed reactions and in food chemistry
industry be required?
Checkpoint C