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
1.
A buffer typically consists of a solution which contains a weak acid HA mixed with the salt of that acid & a strong base e.g. 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 3.
The dogs were infused with 14,000,000 nmoles/l of H+
but the plasma [H+] only changed by a bit over 0.002%. By any analysis, this is a system which powerfully resists change in [H+].
The conclusion is that 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 buffering metabolic acids
Haemoglobin
Important for buffering carbon dioxide
Plasma protein
Minor buffer
Phosphate
Unimportant: Concentration is too low
ICF
Proteins
Important ICF
Inorganic phosphate
Important buffer
Carnosine
Important buffer, esp in skeletal muscle cells
Urine
Phosphate
Responsible for most of 'Titratable Acidity'
Ammonia
Important - formation of NH4+
Bone
Ca carbonate
Important 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 (for the bicarbonate buffer)
pH = pKa + log10 ( [HCO3] / 0.03 x pCO2)
The pKa value is dependent on the temperature, [H+] 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.
Derivation of the Henderson-Hasselbalch equation
Ka is the equilibrium constant for acid dissociation reaction. pKa is the negative log (to base 10) of Ka.
The pKa 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+] x [HCO3-]) / ([CO2] x [H20])
The concentration of H2O is so incredibly large (55.5M or 55,500 mmol/l) compared to the other components, the small loss of water due to
this reaction changes its concentration by only an extremely small amount. To get an idea of what this means, imagine that you have 100 million dollars
in the bank, and you give away $1. The amount your bank account has changed relative to the total amount is so incredibly small that you still have
a $100 million dollars in the bank. Thus back to the situation with water, the dissociation is so incredibly small 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+] x [HCO3-] / [CO2]
Substituting in the equation using:
K'a = 800 nmol/l (the value for plasma at 37C)
[CO2] = 0.03 x pCO2 (by Henry's Law) [wher
e 0.03 is the solubility coefficient when pCO2 is in mmHg, giving [CO2 in mmol/l]
gives the form known as the Henderson Equation:
[H+] = (800 x 0.03 x pCO2) / ([HCO3-]) nmol/l
[H+] = (24 x pCO2) / ([HCO3-] nmol/l
Now this equation can be converted into another form using the following information:
pH = negative log (to base 10) of the [H+] (using [H+] in mol/l)
pK'a = negative log (to base 10) of K'a = log10(800 x 10-9) = 6.1
(and the standard rules of algebra and logs).
Thus, the Henderson-Hasselbalch equation (for the bicarbonate buffer)
The distinction between pKa and p'Ka is now ignored and the Henderson-Hasselbalch equation is always written with pKa rather than p'Ka.
Note on Units: The above form of the equation using 0.03 mmol/l per mmHg for the CO2 solubility coefficient requires pCO2 to
have units of mmHg, and [HCO3-] to be in mmol/l.
If pCO2 is reported in kPa by the clinical laboratory in your region, then the solubility coefficient to use is 0.23mmol/l per kPa.
Closed versus open buffers: the pKa 6.1 issue
On chemical grounds, a buffer system with a pKa of 6.1 is not a good buffer at a pH of 7.4 as the the buffer's pKa is not within the required +/- 1 pH unit of the prevailing pH.
So how can the bicarbonate buffer be an effective buffer in the body?
The bicarbonate buffer system in the body is more complex and the usual phrase used to describe this is to say "the bicarbonate buffer is open at both ends". What this means is that
both [HCO3-] and pCO2 (the terms in the Henderson-Hasselbalch equation) can be adjusted in the body by physiological controls so the 'buffering'
is not just a physicochemical reaction as would occur in a test tube. This ability to physiologically adjust the terms independent of the chemical reaction greatly increases
the effectiveness of the bicarbonate buffer to resist changes in pH when an acid or base is added. This buffer is called an "open buffer". Other buffers in the body
(e.g., the phosphate buffer) cannot do this and are called "closed buffers". The excretion of CO2 via the lungs is the key thing because of the rapidity of the
response. The word buffering here is in a broader sense of something that resists change in pH even though for this buffer other factors such as change in ventilation are
required to make it buffer effective at pH 7.4. Thus:
The bicarbonate buffer system is an effective buffer system in the body at pH 7.4 despite having a low pKa because the body separately controls both pCO2 and [HCO3-]
2.2.4 Other Buffers
The other buffer systems in the blood are the protein and phosphate buffer systems.
These are the only blood buffer systems capable of buffering respiratory acid-base disturbances as the bicarbonate system is ineffective in buffering changes in H+ produced by itself.
Isohydric Principle
All buffer systems which participate in defence of acid-base changes are in equilibrium with each other. There
is after all only one value for [H+] at any moment. This is known as the Isohydric Principle.
It means that an assessment of the concentrations of any one acid-base pair can be utilised to provide a picture of overall
acid-base balance in the body. This is fortunate as the measurement of the concentrations of all the buffer
pairs in the solution would be difficult. Conventionally, the components of the bicarbonate system (ie [HCO3]
and pCO2) alone are measured. They are accessible and easy to determine. Blood gas machines measure pH
and pCO2 directly and the [HCO3] is then calculated using the Henderson-Hasselbalch equation.
Phosphate Buffer
The phosphate buffer system is NOT an important BLOOD buffer as its plasma concentration is too low (despite having a favourable pKa of 6.8)
Phosphate is an important buffer in urine where its concentration is higher.
Phosphoric acid is triprotic weak acid and has a pKa value for each of the three dissociations:
pKa1 = 2
pKa2 = 6.8
pKa3 = 12
H3PO4
<= = = >
H+ + H2PO4-
<= = =>
H+ + HPO4-2
< = = = >
PO4-3 + H+
The three pKa values are sufficiently different so that at any one pH only the members of a single conjugate pair are present in significant concentrations.
At the prevailing pH values in most biological systems, monohydrogen phosphate (HPO4-2) and dihydrogen phosphate
(H2PO4-) are the two species present. The pKa2 is 6.8 and this makes the closed phosphate buffer system a good buffer intracellularly and in
urine where its concentration is higher. The pH of glomerular ultrafiltrate is 7.4 and this means that phosphate will initially be predominantly in the monohydrogen form and
so can combine with more H+ in the renal tubules. This makes the phosphate buffer more effective in buffering against a drop in pH than a rise in pH.
The Henderson-Hasselbalch equation for the phosphate buffer in the body is: pH = 6.8 + log10 ([HPO4-2) / [H2PO4-] )
From that equation, a quick calculation reveals that at pH 7.4, 80% of phosphate is in the monohydrogen phosphate form.
Haemoglobin and Plasma Protein Buffers
Protein buffers in blood include haemoglobin (150g/l) and plasma proteins (70g/l).
Buffering in both is by the imidazole group of the histidine residues which have an effective pKa of about 6.8. This is suitable for effective
buffering at physiological pH.
Haemoglobin is about 6 times more important then the plasma
proteins (as a buffer) as it is present in about twice the concentration and has about three times the number of
histidine residues per molecule. For example if blood pH changed from 7.5 to 6.5, haemoglobin would buffer
27.5 mmol/l of H+ and total plasma protein buffering would account for only 4.2 mmol/l of H+.
The haemoglobin buffer is an important blood buffer for buffering H+ produced from the dissociation of H2CO3.
Although the discusson is always about 'the' haemoglobin buffer, this is actually 2 buffer systems (with different pKa values) with conversion between them by the addition or removal of O2:
OxyHb-H <=> OxyHb + H+
DeoxyHb-H <=> DeoxyHb + H+
The oxyHb buffer unloads O2 in peripheral capillaries and is converted to the more effective deoxyHb buffer
exactly where an increased buffering capacity is required so the change is physiologically advantageous.
The difference between the 2 Hb buffers is usually stated as:
Deoxyhaemoglobin is a better buffer than oxyhaemoglobin
This increase in buffer capacity in peripheral capillaries improves CO2 carriage: it contributes about 30% of the Haldane effect.
The factor accounting for 70% of the Haldane effect is the much greater ability of deoxyhaemoglobin to form carbamino compounds.
Buffering in different sites
Respiratory disorders are predominantly buffered (97%) in the intracellular compartment (e.g., in red cells).
Metabolic disorders have a larger buffering contribution from the extracellular fluid (e.g., ECF buffering of 40% for a metabolic acidosis and
70% for a metabolic alkalosis).
Various buffer systems exist in body fluids (see Table) to minimise the effects on pH of the addition or removal of acid from them.
In ECF, the bicarbonate system is quantitatively the most important for buffering metabolic acids. Its effectiveness is greatly increased by ventilatory changes which attempt to maintain a constant
pCO2 and by renal mechanisms which result in changes in plasma bicarbonate.
In blood, haemoglobin is the most important buffer for H+ produced from conversion of CO2 to bicarbonate because of its high concentration and its large number of histidine residues.
2.2.5 Link between Intracellular & Extracellular Compartments
How are changes in [H+] communicated between the ICF and ECF?
It is very lipid soluble and crosses cell membranes with ease causing acid-base changes due to formation of H+ and HCO3-. Because of this ease of movement, CO2 is not important in causing differences in pH on the two sides of the cell membrane.
Extracellular buffering of CO2 is limited by the inability of the major extracellular buffer (the bicarbonate system) to buffer changes in [H+] produced from the reaction between CO2 and water.
The result is that buffering for respiratory acid-base disorders is predominantly intracellular: 99% for respiratory acidosis and 97% for respiratory alkalosis.
The second major process which allows transfer of H+ ions intracellularly is entry of H+ in exchange for
either K+ or Na+. This ionic exchange is necessary to maintain electroneutrality. This cation exchange is the mechanism
which delivers H+ intracellularly for buffering of a metabolic disorder. In the cell, the protein and phosphates
(organic and inorganic) buffer the H+ delivered by this ion exchange mechanism.
Experiments in metabolic acidosis have shown that 57% of buffering occurs intracellularly and 43% occurs
extracellularly. The processes involved in this buffering are:
Processes involved in Buffering
ECF
43% (by bicarbonate & protein buffers)
ICF
57% (by protein phosphate and bicarbonate buffers) due to entry of H+ by:
Na+-H+ exchange 36%
K+-H+ exchange 15%
Other 6%
(see Section 10.6 for a chemical explanation of how an exchange of Na+ or K+ for H+ across a membrane can alter the pH by changing the strong ion difference or 'SID')
Thirty percent (30%) of the buffering of a metabolic alkalosis occurs intracellularly and Na+-H+ exchange is responsible for most of the transfer of H+.
2.2.6 Role of Bone Buffering
The carbonate and phosphate salts in bone act as a long term supply of buffer especially during prolonged metabolic acidosis.
The important role of bone buffers is often omitted from discussions of acid-base physiology4.
Bone consists of matrix within which specialised cells are dispersed. The matrix is composed of organic
[collagen and other proteins in ground substance] and inorganic [hydroxyapatite crystals: general formula
Ca10(PO4)6(OH)2] components. The hydroxyapatite crystals make up two-thirds of
the total bone volume but they are extremely small and consequently have a huge total surface area. The crystals
contain a large amount of carbonate (CO3-2) as this anion can be substituted for both phosphate
and hydroxyl in the apatite crystals. Bone is the major CO2 reservoir in the body and contains carbonate
and bicarbonate equivalent to 5 moles of CO2 out of a total body CO2 store of 6 moles. (Compare
this with the basal daily CO2 production of 12 moles/day)
CO2 in bone is in two forms: bicarbonate (HCO3-) and carbonate (CO3-2).
The bicarbonate makes up a readily exchangeable pool because it is present in the bone water which makes up the hydration shell
around each of the hydroxyapatite crystals. The carbonate is present in the crystals and its release requires
dissolution of the crystals. This is a much slower process but the amounts of buffer involved are much larger.
How does bone act as a buffer?
Two processes are involved:
Ionic exchange
Dissolution of bone crystal
Bone can take up H+ in exchange for Ca++, Na+ and K+ (ionic exchange) or release
of HCO3-, CO3- or HPO4-2. In acute metabolic acidosis uptake of H+
by bone in exchange for Na+ and K+ is involved in buffering
as this can occur rapidly without any bone breakdown. A part of the so called
'intracellular buffering'of acute metabolic disorders may represent some of this acute buffering by bone. In
chronic metabolic acidosis, the major buffering mechanism by far is release of calcium carbonate from bone.
The mechanism by which this dissolution of bone crystal occurs involves two processes:
direct physicochemical breakdown of crystals in response to [H+]
osteoclastic reabsorption of bone.
The involvement of these processes in buffering is independent of parathyroid hormone. Intracellular acidosis in osteoclasts results in a decrease in intracellular Ca++ and this stimulates these cells.
Bone is probably involved in providing some buffering for all acid-base disturbances. Little experimental evidence is available for respiratory disorders. Most research has been concerned with chronic metabolic
acidoses as these conditions are associated with significant loss of bone mineral (osteomalacia, osteoporosis).
In terms of duration only two types of metabolic acidosis are long-lasting enough to be associated with loss
of bone mineral: renal tubular acidosis (RTA) and uraemic acidosis. Bone is an important buffer in these two conditions.
In uraemia, additional factors are more significant in causing the renal osteodystrophy as the loss of bone mineral cannot be explained by the acidosis alone. Changes in vitamin D metabolism, phosphate
metabolism and secondary hyperparathyroidism are more important than the acidosis in causing loss of bone mineral in uraemic patients. The loss of bone mineral due to these other factors releases substantial amounts of buffer.
References
Worthley LI. Hydrogen ion metabolism. Anaesth Intensive Care 1977 Nov; 5(4) 347-60. PubMed
Pitts RF. Mechanisms for stabilizing the alkaline reserves of the body. Harvey Lect 1952-1953; 48 172-209. PubMed
Bernards WC Interpretation of Clinical Acid-Base Data. Regional Refresher Courses in Anesthesiology. 1973; 1: 17-26
Bushinsky DA. Acidosis and bone. Miner Electrolyte Metab 1994; 20(1-2) 40-52.PubMed