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2.2: Buffering

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    Definition of a Buffer

    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+.

    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.

    Swan & Pitts Experiment 2

    In this experiment, dogs received an infusion of 14 mmols H+ per litre of body water. This caused a drop in pH from 7.44\(([{H}^{+}]=36\:\text{nmoles/l})\) or a pH of 7.14 \(([{H}^{+}]=72\:\text{nmoles/l})\). That is, a rise in [H+] of only 36 nmoles/l.

    If you just looked at the change in [H+] then you would only notice an increase of 36 nmoles/l and you would have to wonder what had happened to the other 13,999,964 nmoles/l that were infused.

    Where did the missing H+ go?

    They were hidden on buffers and so these hydrogen ions were hidden from view.

    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.)

    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


    Buffer System




    For metabolic acids


    Not important because concentration too low


    Not important because concentration too low



    Important for metabolic acids


    Important for carbon dioxide

    Plasma protein

    Minor buffer


    Concentration too low



    Important buffer


    Important buffer



    Responsible for most of 'Titratable Acidity'


    Important - formation of NH4+


    Ca carbonate

    Important in prolonged metabolic acidosis

    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

    \[ pH=pKa + \log_{10}(\frac {[HCO_{3}]}{0.03}) \times \: pCO_{2} \]

    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.

    A note on terminology: 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:

    \[ CO_{2} + H_{2}O \Leftrightarrow H_{2}CO_{3} \Leftrightarrow {H}^{+} + HCO_{3}^{-}\]

    (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:

    \[ CO_{2} + H_{2}O \Leftrightarrow {H}^{+} + HCO_{3}^{-}\]

    By the Law of Mass Action:

    \[K_{a} =\frac {[{H}^{+}] \cdot [HCO_{3}^{-}]} {[CO_{2}] \cdot [H_{2}O]} \]

    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.

    \[ K' _{a} = K_{a} \times [H_{2}O] = [H^{+}] \cdot \frac {[HCO_{3}^{-}]} {[CO_{2}]} \]

    Substituting in the equation using:

    \[ K'_{a} = 800 nmol/L \: \text {(value for plasma at 37°C)}\]

    \[ [CO_{2}] = 0.03 \times pCO_{2} \text {(by Henry's Law) [where 0.03 is the solubility coefficient]}\]

    gives the form known as the Henderson Equation:

    \[ [H^{+}] = (800 \times 0.03) \cdot \frac {pCO_{2}} {HCO_{3}^{-}} \]

    \[ [H^{+}] = 24 \times \frac {pCO_{2}} {[HCO_{3}^{-}]} nmol/l \]

    Now this equation can be converted into another form using the following information (where [H+] is in mol/l ) and the standard rules of algebra and logs:

    \[ pH= \log_{10} [H^{+}] \]

    \[pK'_{a} = -\log_{10} K'_{a} = \log_{10} (800 \times 10^{-9}) = 6.1 \]

    Thus, the Henderson-Hasselbalch equation

    \[ -\log_{10}[H^{+}] = -\log_{10}(800 \times 10^{-9}) + \log \frac {[HCO_{3}^{-}]} {0.03 pCO_{2}} \]

    \[ pH = pK'_{a} + \log \frac {[HCO_{3}^{-}]} {0.03 pCO_{2}} \]

    \[pH = 6.1 + \log \frac {[HCO_{3}^{-}]}{0.03pCO_{2}}\]


    The distinction between pKa and p'Ka is usually forgotten and the Henderson-Hasselbalch equation is always written with pKa

    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 the key thing because of the rapidity of the response. The adjustment of pCO2 by change in alveolar ventilation has been referred to as physiological buffering.

    Note: This use of the word buffering is in the broader sense of something that resists change in a property, and is different from the definition of buffering (or 'physiological buffering') given at the top of this page. This shift in meaning of buffering can be confusing because the word buffering is mostly used in speech and in articles without the qualification of either 'physicochemical' or 'physiological' (or some qualifying word).

    The bicarbonate buffer system is an effective buffer system despite having a low pKa because the body also controls pCO2

    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.

    The phosphate buffer system is NOT an important blood buffer as its concentration is too low

    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:

    \(pK_{a1} = 2 \)

    \(pK_{a2} = 6.8 \)

    \( pK_{a3} = 12 \)


    \( \Leftrightarrow \)

    \( H^{+} + H_{2}PO_{4}^{-} \)

    \( \Leftrightarrow \)

    \( H^{+} + HPO_{4}^{2-} \)

    \(\Leftrightarrow \)

    \(H^{+} + PO_{4}^{3-} \)

    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.


    The 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 in these two sites is much higher than in extracellular fluid.

    Haemoglobin is an important blood buffer particularly for buffering CO2

    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 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+.

    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.

    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.

    Buffering in different sites

    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 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 CO2 because of its high concentration and its large number of histidine residues.

    Deoxyhaemoglobin is a better buffer than oxyhaemoglobin

    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.

    Link between Intracellular & Extracellular Compartments

    How are changes in [H+] communicated between the ICF and ECF?

    The two major processes involved are:

    • Transfer of CO2 across the cell membrane
    • Ionic shifts (ie proton-cation exchange mechanisms)

    Important points to note about CO2 are:

    • 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


    43% (by bicarbonate & protein buffers)


    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-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+.

    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 \( Ca_{10}(PO_{4})_{6}(OH)_{2}) \)] 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 Ca2+, 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 Ca2+ 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.


    • Bone is an important source of buffer in chronic metabolic acidosis (ie renal tubular acidosis & uraemic acidosis)
    • Bone is probably involved in providing some buffering (mostly by ionic exchange) in most acute acid-base disorders but this has been little studied.
    • Release of calcium carbonate from bone is the most important buffering mechanism involved in chronic metabolic acidosis.
    • Loss of bone crystal in uraemic acidosis is multifactorial and acidosis is only a minor factor
    • BOTH the acidosis and the vitamin D3 changes are responsible for the osteomalacia that occurs with renal tubular acidosis.


    1. Worthley LI. Hydrogen ion metabolism. Anaesth Intensive Care 1977 Nov; 5(4) 347-60. pmid:23014. PubMed
    2. Pitts RF. Mechanisms for stabilizing the alkaline reserves of the body. Harvey Lect 1952-1953; 48 172-209. PubMed
    3. Bernards WC Interpretation of Clinical Acid-Base Data. Regional Refresher Courses in Anesthesiology. 1973; 1: 17-26
    4. Bushinsky DA. Acidosis and bone. Miner Electrolyte Metab 1994; 20(1-2) 40-52. PubMed

    This page titled 2.2: Buffering is shared under a CC BY-NC-SA 2.0 license and was authored, remixed, and/or curated by Kerry Brandis via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request.