A buffer is a solution containing substances which have the ability to minimise changes in pH when an acid or base is added
to it. (Worthley, 1977).
A buffer typically consists of a solution which contains a weak acid HA mixed with the salt of that acid & a strong base
eg NaA. The principle is that the salt provides a reservoir of A- to replenish [A-] when A- is removed by reaction with H+.
2.2.2 Buffers in the Body
The body has a very large buffer capacity.
This can be illustrated by considering an old experiment (see below) where dilute hydrochloric acid was infused into a dog.
Before we proceed, lets just make sure we appreciate what
this experiment reveals. The dogs were infused with 14,000,000 nmoles/l of H+
but the plasma [H+] only changed by bit over 0.002%. By any analysis, this is a
system which powerfully resists change in [H+]. (My personal
analogy on appreciating the magnitude of this is to use the analogy of
depositing $14,000,000 in the bank, but then finding that after 'bank charges'
my account only went up by $36.)
Make no mistake: the body has:
a HUGE buffering capacity, and
this system is essentially IMMEDIATE in effect.
For these 2 reasons, physicochemical buffering provides a powerful first defence against acid-base perturbations.
Buffering hides from view the real change in H+ that occurs.
This huge buffer capacity has another not immediately obvious implication for how we think about the severity of an
acid-base disorder. You would think that the magnitude of an acid-base disturbance could be quantified merely by looking
at the change in [H+] - BUT this is not so.
Because of the large buffering capacity, the actual change in
[H+] is so small it can be ignored in any quantitative assessment, and instead, the magnitude of a disorder
has to be estimated indirectly from the decrease in the total concentration of the anions involved in the buffering.
The buffer anions, represented as A-, decrease because they combine stoichiometrically with H+
to produce HA. A decrease in A- by 1 mmol/l represents a 1,000,000 nano-mol/l amount of H+ that
is hidden from view and this is several orders of magnitude higher than the visible few nanomoles/l change in
[H+] that is visible.) - As noted above in the comments about the Swan & Pitts experiment,
13,999,994 out of 14,000,000 nano-moles/l of H+ were hidden on buffers and just to count the 36 that
were on view would give a false impression of the magnitude of the disorder.
The Major Body Buffer Systems
Site
Buffer System
Comment
ISF
Bicarbonate
For metabolic acids
Phosphate
Not important because concentration too low
Protein
Not important because concentration too low
Blood
Bicarbonate
Important for metabolic acids
Haemoglobin
Important for carbon dioxide
Plasma protein
Minor buffer
Phosphate
Concentration too low
ICF
Proteins
Important buffer
Phosphates
Important buffer
Urine
Phosphate
Responsible for most of 'Titratable Acidity'
Ammonia
Important - formation of NH4+
Bone
Ca carbonate
In prolonged metabolic acidosis
2.2.3 The Bicarbonate Buffer System
The major buffer system in the ECF is the CO2-bicarbonate buffer system. This is responsible for
about 80% of extracellular buffering. It is the most important ECF buffer for metabolic acids but
it cannot buffer respiratory acid-base disorders.
The components are easily measured and are related to each other by the Henderson-Hasselbalch equation.
Henderson-Hasselbalch Equation
pH = pK’a + log10 ( [HCO3] / 0.03 x pCO2)
The pK’a value is dependent on the temperature, [H,sup>+] and the ionic concentration of the solution. It has a
value of 6.099 at a temperature of 37C and a plasma pH of 7.4. At a temperature of 30C and pH of 7.0, it has
a value of 6.148. For practical purposes, a value of 6.1 is generally assumed and corrections for temperature,
pH of plasma and ionic strength are not used except in precise experimental work.
The pK'a is derived from the Ka value of the following reaction:
CO2 + H2O <=> H2CO3 <=> H+ + HCO3-
(where CO2 refers to dissolved CO2)
The concentration of carbonic acid is very low compared to the other components so the above equation is
usually simplified to:
CO2 + H2O <=> H+ + HCO3-
By the Law of Mass Action:
Ka = [H+] . [HCO3-] / [CO2] . [H20]
The concentration of H2O is so large (55.5M) compared to the other components, the small loss of water due to
this reaction changes its concentration by only an extremely small amount. This means that [H2O] is effectively
constant. This allows further simplification as the two constants (Ka and [H2O] ) can be combined into a new
constant K’a.
K’a = Ka x [H2O] = [H+] . [HCO3-] / [CO2]
Substituting:
K'a = 800 nmol/l (value for plasma at 37C)
[CO2] = 0.03 x pCO2 (by Henry’s Law) [where 0.03 is the solubility coefficient]
into the equation yields the Henderson Equation:
[H+] = (800 x 0.03) x pCO2 / [HCO3-] = 24 x pCO2 / [HCO3-] nmol/l
Taking the logs (to base 10) of both sides yields the Henderson-Hasselbalch equation:
pH = log10(800) - log (0.03 pCO2 / [HCO3-] )
pH = 6.1 + log ( [HCO3] / 0.03 pCO2 )
On chemical grounds, a substance with a pKa of 6.1 should not be a good buffer at a pH of 7.4 if it were a
simple buffer. The system is more complex as it is ‘open at both ends’ (meaning both [HCO3] and pCO2 can be
adjusted) and this greatly increases the buffering effectiveness of this system. The excretion of CO2 via the
lungs is particularly important because of the rapidity of the response. The adjustment of pCO2 by change in
alveolar ventilation has been referred to as physiological buffering.
The bicarbonate buffer system is an effective buffer system despite having a low pKa because the
body also controls pCO2