Recap – Last Lecture
Predicting Reactions
•  An acid is a proton donor
•  A base is a proton acceptor
•  A reaction between a strong acid and a strong base
will go to completion.
•  A conjugate pair differ by H+
•  A strong acid – weak base or weak acid – strong
base reaction will go to completion.
•  Strong A/B is completely dissociated
•  Weak A/B is in equilibrium
eg CH3COOH + OH– → CH3COO– + H2O
•  For a reaction between a weak acid and a weak base,
a comparison of pKa values enables us to determine
whether the reaction will occur.
•  The smaller pKa, the stronger the acid and the
weaker its conjugate base.
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Predicting Reactions
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Buffer
•  Example: Will hydrogencarbonate (HCO3–) react with
phenol (C6H5OH)?
C6H5OH + HCO3–
⇌ C6H5O– + H2CO3
pKa=10.0
pKa=6.35
• 
A buffer is a solution composed of moderate quantities
of both members of a conjugate acid-base pair
(e.g. CH3COOH and CH3COO– (Na+)).
• 
It maintains a solution at approximately constant pH
even when small quantities of H+ or OH– are added.
The answer is ‘no’. The equilibrium lies to the left.
•  But hydorgencarbonate will react with acetic acid:
CH3COOH + HCO3– ⟶ CH3COO– + H2CO3
pKa=4.7
pKa=6.35
Conjugate pair
p Ka
Optimum pH of buffer
CH3COOH / CH3COO-
4.76
H2PO4- / HPO42-
7.20
4.76
7.20
NH4+ / NH3
9.26
9.26
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Relationship between pKa and pH
•  The equilibrium between a conjugate acid – base pair
is affected by pH.
HA
⇌
H+ + A–
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Example
•  Which of CH3COOH / CH3COO– will dominate at a
physiological pH = 7.4 given CH3COOH pKa = 4.76?
•  At high [H+], (low pH), the equilibrium is towards the
left and visa versa.
•  Comparison of pH with pKa of the weak acid/base
system indicates in which direction the equilibrium lies.
•  If the pH is on the ‘acid side’ of the pKa, the conjugate
acid will predominate.
•  If the pH is on the ‘base side’ of the pKa, the
conjugate base will predominate.
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CH3COOH(aq)
⇌
CH3COO–(aq) + H+(aq)
Answer: pH > pKa
ie pH is on the ‘base
side’ of the pKa so
the conjugate base
will dominate:
CH3COO–
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Example
Amino Acids
•  A more complicated diagram results for a polyprotic
acid, eg H3PO4.
•  Note pKa1 < pKa2 < pKa3
• 
• 
An amino acid contains (at least) two groups that are
acid base active – a carboxylic acid and an amine.
Each group has a pKa associated with it.
e.g. glycine
pKa (H3PO4) = 2.2
pKa (H2PO4-) = 7.2
pKa (HPO42-) = 12.3
pK a = 9.60
(of conjugate acid)
H
H
N
C
H
H
O
OH
At physiological pH, this will exist in an ionic form:
H
H
H
N
C
H
H
O
C
O
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Applications – CO2 in blood
• 
H2CO3(aq) ⇌
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Applications – O2 in blood
Buffers are an important part of living systems. Human blood has a normal pH range
of 7.35-7.45. Any deviation from this can disrupt cell membranes, proteins and
enzyme activity. The major buffer solution that controls blood pH is
CO2(g) + H2O(l) ⇌
pK a = 2.34
C
H+(aq) + HCO3–(aq) … (1)
• 
Regulation of pH in the blood relates directly to effectiveness of oxygen transport.
The blood transports oxygen using haemoglobin (Hb) which reversibly binds H+ and
O2 in a competitive equilibrium:
HbH+ + O2
Notes:
•  Although carbonic acid (H2CO3)is diprotic, only the first ionization is important.
•  CO2 is a gas which provides the body with a mechanism to adjust the equilibrium.
Removal of CO2 by exhalation shifts the equilibrium to the left, consuming H+ ions.
•  Buffer solution in blood plasma has [HCO3–] ~ 0.024 M and [H2CO3] ~ 0.0012 M
(20/1 ratio). As a consequence the buffer solution has a high capacity to neutralize
additional acid, but low capacity to neutralize additional base.
•  Principle organs that regulate blood pH are the lungs and kidneys with input from
brain sensors.
•  When [CO2] rises, equilibrium shifts to the right which increases [H+] which, in turn,
triggers receptors that increase breathing rate and elimination of CO2 to restore pH.
•  Kidneys absorb & release H+ and HCO3–; much excess acid leaves the body in urine
(pH 5-7).
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⇌
HbO2 + H+
………(2)
Notes:
•  When the blood reaches the tissue where the [O2] is low, the equilibrium is shifted to
the left to release O2. An increase in [H+] and an increase in temperature also shifts
the equilibrium to the left (ΔH < 0, exothermic).
•  During exertion several factors work to ensure delivery O2 to active tissue.
•  As O2 is consumed the equilibrium 2 shifts to the left.
•  Exertion raises body temp which also shifts equilibrium 2 to the left.
•  Increased metabolic production of CO2 shifts equilibrium 1 to the right increasing [H+]
which also shifts equilibrium 2 to the left. Other acids such as lactic acid are
produced when tissue is starved of O2 which further shifts equilibrium 2 to the left.
(This extra production of acids may produce cramps.)
•  Increased [H+] stimulates increase rate of breathing, which provides more O2 and
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eliminates CO2.
Learning Outcomes:
•  By the end of this lecture, you should:
−  predict the outcome of an acid - base reaction.
−  understand the function and composition of a
buffer.
−  recognise the relationship between pH, pKa and
the form of the conjugate pair present.
−  be able to predict the charge state of an amino
acid at a particular pH.
−  be able to complete the worksheet (if you
haven’t already done so…)
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