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 eg NaA. The principle is that the salt provides a reservoir of A- to replenish [A-] when A- is removed by reaction with H+.
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+]. (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:
For these 2 reasons, physicochemical buffering provides a powerful first defence against acid-base perturbations.
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
For metabolic acids
Not important because concentration too low
Not important because concentration too low
Important for metabolic acids
Important for carbon dioxide
Concentration too low
Responsible for most of 'Titratable Acidity'
Important - formation of NH4+
In prolonged metabolic acidosis
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.
pH = pK’a + log10 ( [HCO3] / 0.03 x pCO2)
The pK’a 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.
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]
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 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.
The concentration of phosphate in the blood is so low that it is quantitatively unimportant. Phosphates are important buffers intracellularly and in urine where their 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
<= = = >
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. 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.
Note: The ‘true’ pKa2 value is actually 7.2 if measured at zero ionic strength but at the typical ionic strength found in the body its apparent value is 6.8. The other factor which makes phosphate a more effective buffer intracellularly and in urine is that its concentration is much higher here than in extracellular fluid.
Protein buffers in blood include haemoglobin (150g/l) and plasma proteins (70g/l). Buffering is by the imidazole group of the histidine residues which has a pKa of about 6.8. This is suitable for effective buffering at physiological pH. Haemoglobin is quantitatively about 6 times more important then the plasma proteins as it is present in about twice the concentration and contains 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+.
Deoxyhaemoglobin is a more effective buffer than oxyhaemoglobin and this change in buffer capacity contributes about 30% of the Haldane effect. The major factor accounting for the Haldane effect in CO2 transport is the much greater ability of deoxyhaemoglobin to form carbamino compounds.
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.
Respiratory disorders are predominantly buffered in the intracellular compartment. Metabolic disorders have a larger buffering contribution from the extracellular fluid (eg 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 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 CO2 because of its high concentration and its large number of histidine residues.
Another factor which makes haemoglobin an important buffer is the phenomemon of isohydric exchange. That is, the buffer system (HHbO2-HbO2-) is converted to another more effective buffer (HHb-Hb-) exactly at the site where an increased buffering capacity is required. More simply, this means that oxygen unloading increases the amount of deoxyhaemoglobin and this better buffer is produced at exactly the place where additional H+ are being produced because of bicarbonate production for CO2 transport in the red cells.
The two major processes involved are:
Important points to note about CO2 are:
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+. 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
43% (by bicarbonate & protein buffers)
57% (by protein phosphate and bicarbonate buffers) due to entry of H+ by:
(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-two percent (32%) of the buffering of a metabolic alkalosis occurs intracellularly and Na+-H+ exchange is responsible for most of the transfer of H+.
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.
Two processes are involved:
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:
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.