According to the Brönsted-Lowery classification an acid is a proton donor, HA gives H⁺+ A^-, and a base is a proton acceptor, B + H⁺ gives BH⁺. This classification makes no mention of solvent. However, in biological systems the medium is aqueous solution, and we will confine our attention to this condition.

Water can behave both as an acid and a base (amphiprotic character). This means that one H₂O molecule, acting as an acid, may donate a proton to an other H₂O molecule, acting as a base.

The hydrated proton (H₃O+) is termed hydronium ion and OH- is termed hydroxyl ion. A proton transfer equilibrium involving a single substance of the type just mentioned above is an example of an autoprotolysis equilibrium. The autoprotolysis constant of water (K_w) is defined by

K_w = [H₃O⁺] x [OH^-].

Taking the logaritm of both sides we obtain

log K_w = log [H₃O⁺] x [OH^-],

and changing the sign of both sides gives:

pK_w = -log [H₃O⁺] x [OH^-].

Because the [H₃O⁺] and [OH^-] vary over a wide range pH and pOH scaled have been introduced:

pH = -log [H₃O⁺]

and

pOH = -log [OH^-].

At 25°C the autoprotolysis constant of water K_w = 1x 10^-14. Further, in pure water [H₃O⁺] = [OH^-]. Accordingly, pOH = pH = 7 = 1/2 pK_w. If we put an acid into a solution to decrease th pH the pOH must increase to keep their sum equal to pK_w.

The equilibrium of an acid and a base in water can be expressed in terms of the acidity constant K_a and the base constant K_b, respectively:

K_a= [H₃O⁺][A^-]/[HA]

and

K_b= [OH^-][HA]/[A^-].

These constants vary over a wide range. It is therefore convenient to list them as their logarithms, and we introduce:

pK_a= -log K_a

and

pK_b= -log K_b.

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2. Buffers

Buffer action means the ability to oppose charges in pH when small amounts of acid or base are added to the solution. The mathematical basis of buffer action is given by the Henderson-Hasselbalch equation. If we plot the funktion we see that the graph is relatively flat near pH = pK_a.

The Henderson-Hasselbalch equation:

pH = pKa + log [A^-]/[HA].

The physical basis of buffer action is that an abundant supply of A^- ions can remove H⁺ ions brought by additional acid, and an abundant supply of HA molecules can supply H⁺ ions to react with any added base.

3. Fluid spaces in the human body

About 60% of the body consists of water. Roughly 2/3 of this water is confined to the intracellular compartment – intracellular fluid (ICF). The remaining 1/3 is regarded as extracellular fluid (ECF). From the acid-base point of view the erythrocytes can be considered part of the ECF, since the erythrocytes are very permeable to HCO₃^- ions. Any acute acid-base disturbances noted in the ECF, which are not a result of a general intracellular process, are initially buffered mainly by the ECF (including the erythrocytes). After a while the ECF will be buffered by bone and intracellular buffers too. When H+ ions move into cells other positive ions move out. The main intracellular. The main intracellular cation is K⁺. Accordingly, a decrease in extracellular pH tends to give rise to hyperkalemia.

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4. Acid-base balance

Protons are continuously produced in the body. Aside from crabonic acid the protons are mainly decived from ureagenesis, anaerobic glycolysis and ketogenesis. The pH in the ECF is maintained at about 7.4 by buffere in the ECF, ICF and bone, and by the lungs and the kidneys.

Nowadays, it is thought that one of the primary goals of the acid-base regulation is to preserve a relatively unchanged ionization of histidine, which is an amino acid that takess an aktive part in proton exchange within the physiological range of the pH scale. This is often regerred to as a–state-regulation, where a labels the ratio of ionized and unionized a-amino grops of histidine.

The range of pH in blood plasma which is compatible with life for more than a short period of time is about 6.8 - 7.7.

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4.1 Extracellular buffers

The most important buffers in the ECF (blood included) are the bicarbonate system, proteins (mainly hemoglobin), and the phosphate system. This system of buffers may be subdividend into the bicarbonate/carbonic acid buffer sytem, which is an open sytem (see below) and the other buffers, which are often (when considered together) regerred to as Hbuf (acid) and Buf^– (base). A buffer system where the concentration of a component can be changed is called an open sytem. A closed system, the converse, is a system where none of the operating components can enter or leave the sytem, i.e. the sum of acids and bases remins constant.

4.1a The bicarbonate/carbonic acid system

The bicarbonare/carbonic acid sytem has a gaseous component which can be released into the surrounding air. The Henderson-Hasselbalch equation, adapted to the bicarbonate/carbonic acid sytem, can be written as

pH = 6.1 + log [HCO₃^-]/[a Pco₂]

where a (soubility coefficient) is 0.226 mM/kPa and Pco₂ is expressed in kPa,

4.1b The phosphate/phosphoric acid buffer sytem

At physiological pH this system consists of two elements: H₂PO₄^- and HPO₄^2-. The pK_a of the phosphate system is 6.8, which means that the body fluids (especially ICF) this buffer system operates close to its maximum buffering capecity. The buffering capacity of this system in the ECF is far less than that of the bicarbonate system. However, in the tubular fluid of the kidneys and in the ICF the phosphate buffer is important. This is because the phosphate concentration is higher there than in the ECF. Further, the pH is lower in the ICF than in ECF as is usually the case in the urine as well. The Henderson-Hasselbalch equation, adapted to the phosphate/phosphoric acid system, can be written as:

pH = 6.8 + log [HPO₄^2-]/[H₂PO₄^-]

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4.1c. The hemoglobin system

Hemoglobin has three important funktions which are interrelated:

A. O₂ transport

B. Transport of CO₂

C. Buffering

A. At a Po₂ of 13.3 kPa hemoglobin, at a concentration of 150 g/l, permis the blood to carry about 65 times more oxygen than plasma would. The affinity for O₂ is reduced by an increase in [H⁺] (Bohr effekt), Pco₂ (mainly via increasing [H⁺]), 2,3-diphosphoglycerate and/or temperature.

B. About 20% of the CO₂ carried from the tissues to the lungs is combined with hemoglobin and plasmaproteins.

C. Although there is a relatively low concentration of hemoglobin in the blood it functions as an effective buffer because it has several buffering groups per molecule. The various buffering grops have different pK_a values (6.5-7.8), so buffering by hemoglobin cannot be characterized by a single Henderson-Hasselbalch equation. The numerus imidazole grops, with pK_a values of approximately 7, are responsible for much of the buffering carried out by hemoglobin in the blood. Oxygenated haemoglobin (H:Hb0₂) is a stronger acid than deoxygenated haemoglobin (H:Hb). This means that oxygenation of a haemoglobin solution will lead to a decrease in pH or it can be put the oppositeway - release of O₂ promotes uptake of protons (Haldane effect).

4.1d The role of the erythrocyte in the bicarbonate system

The C02 liberated from the aerobic catabolism may be transported by the blood from the tissues to the lungs in three different ways:

• Formation of HCOs' ions (70%).

• Forming carbamino compounds with plasma proteins (20%).

• Physically dissolved in plasma (10%).

The way via HCOs' is the most important, since the amount of dissolved C02 is small and the carbamino forming capacity of plasma proteins is low.

In the peripheral tissues the CO₂ diffuses into the red cell where the hydration is catalyzed by carbonic anhydrase. The concomitant dissociation of carbonic acid occurs in a fraktion of a sekond. Becausee of this reaction the concentration HCO₃^- in the erythrocyte is much higher than the concentration HCO₃^- in plasma. The protons are taken up by hemoglobin. Since the red cell membrane is highly anion permeable, due to an anion exchange mechanism, HCO₃^- starts to diffuse out of the cell. The exchange mechanism transports other negative ions into the cell. The natural ion for exchange is chloride because it is available at a relatively high concentration. In this way the intracellular hemoglobin buffers ECF. In the lungs H⁺and HCO₃^-will form CO₂ which diffuses out of the erythrocyte. The concentration of HCO₃^- in the erythrocyte beckmes lower than the HCO₃^- concentration in plasma. Thus, HCO₃^- ions are transportred into the erythrocyte in exchange for Cl^-. Within the span of one sekond (the approximale transit time of blood in the lung capillaries) the CO₂ carried in the blood must be released. Without carbonic anhydrase it would take more than a minute to liberale the CO₂ released during a single breath. Thus, the carbonic anhydrase plays an important role to the CO₂ transport.

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4.2 Intracellular and bone buffering

About 60% of the protons administered intravenously are thought to be buffered by bone and the intracellular compartment of various tissues. The immediately available buffers are those of the blood and extracellular fluid. However, after some time, intracellular buffers and bone add their contribution.

More generally, any process that reversibly consumes or releases protons contributes to pH buffering. Accordingly, the cytoplasmic pH can be buffered by three kinds of processes:

· Weak acids and bases (protein, bicarbonate and phosphate) - physicochemical buffering.

· Biochemical reactions - biochemical buffering.

· Transport of acid-base components across organellar membranes - organellar buffering.

The uptake of protons by bone can occur in exchange for surface Na⁺ and K⁺and by dissolution of bone mineral. However, the buffering is not only due to physicochemical processes. The activity of the osteoblasts and osteoclasts has also been shown to play a significant role in the buffering process.

4.3 The lungs

By changing the rate and depth of breathing, the concentration of CO₂ in the blood can be controlled. The respiration is governed by Pco₂ (H⁺) and Po₂. Under physiological conditions it is mainly Pco₂, via central chemoreceptors in the brain stem, that governs ventilation. A decrease in blood pH causes an increase in Pco₂.The CO₂ molecule diffuses easily through the blood-brain barrier (BBB). CO₂ is hydrated in the cerebral tissue, thereby forming H₂CO₃ which dissociated into H⁺ and HCO₃^-. The increase in [H⁺] stimulates the respiratory center. Acidemia is also a stimulus to peripheral chemoreceptors in the carotid body. The relative importance of these chemoreceptors in the respiratory reponse to metabolic acid.base disturbances is still controversial.

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4.3 The kidneys

The kidneys are responsible for the elimination of H⁺ and exert control over the concentration of HCO₃^-in blood. Based on the pH of the primary urine the kidneys can reclaim bicarbonate ions from the urine and return them to the blood, or they can allow them to leave the body. In the former case the excess protons leave the body together with another anion - usually phosphate. The phosphate/phosphoric acid pair is the main buffer in the urine. The kidneys are able to secrete H⁺ (by an energy dependent proton pump) against a concentration gradient of about 800:1 and consequently the urine pH may fall to 4.5. The kidney's ability to secrete an alkaline urine is very limited and the maximum pH is 7.8.

Although the ammonium/ammonia pair cannot function as a buffer in the conventional sense (pKa = 9.4), it plays an important role as a major vehicle in carrying protons in the urine. In a normal healthy adult about twice as much acid can be excreted at a urinary pH of 4.5 than would be the case if the ammonia mechanism were not present.

5 Basic concepts in acid-base physiology

pH: -log[H⁺] in blood plasma. Normal value is 7.4. Acidemia is defined as pH<7.36 and alkalemia as pH>7.44.

Pco₂: Partial pressure of carbon dioxide (mm Hg or kPa) in a gas phase in equilibrium with the blood - when there is no net exchange of carbondioxide molecules between the blood and the gas phase. Normal value is 4.6 – 6.0 kPa for healthy adults.

SHCO₃^-: Standard bicarbonate is defined as the [HCO₃^-] in plasma at a Pco₂ of 5.3 kPa and full oxygenation. Normal value is 22-27 mM for healthy adults.

BB: Buffer Base is the total concentration of buffering bases in whole blood (mM). Normal value is 48 ± 3 mM for healthy adults.

NBB: Normal Buffer Base refers to the BB value at normal pH and normal Pco₂. Thus, NBB is obtained by titrating the blood sample with strong acid or base to pH = 7.40 when the Pco₂ = 5.3 kPa. The total concentration of buffering bases in the titrated whole blood is the value of NBB. The value of NBB is dependent mainly on the concentration of HC0₃^-, hemoglobin and other proteins. For health adults the value is 48 ± 3 mM.

ABE: Actual Base Excess is defined as BB - NBB This quantity, often called Base deficit when the value is negative, measures the charge of bases in the sampled blood. Normal value for health adults is 0 ± 3 mM. ABE is an in vitro expression.

SBE: Standard Base Excess SBE estimates the change of bases in the whole extracellular compartment, the blood included. SBE is an in vivo expression.

Anion gap: This quantity is defined as [Na⁺ - ([Cl^-] + [HCO₃^-]). Normal value is 12 ± 2 mEq/l. Since the plasma proteins account for most of the missing anions, the anion gap must be adjusted according to the actual plasma albumin concentration. The approximate correction is a reduction in the anion gap of 0.25 mEq/l for every g/l decline in plasma albumin concentration. The anion gap is useful for classifying metabolic

acidosis. See Tab. 1.

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6 Investigation of acid-base balance

For the evaluation of the acid-base balance an arterial blood sample is needed, in which the acid-base parameters are to be measured. The pH value determines if there is a normal state, an acidemia or an alkalemia. Processes that tend to raise or lower the pH are called alkalosis and acidosis, respectively. To determine the primary cause of an acid-

base disturbance and if there is any accompanying physiologic compensatory response, some further measurements are needed. The Pco₂ measures the respiratory component of an acid-base disturbance. The most useful measure of the metabolic component is SBE. If SBE is not known ABE or SHCO₃^- can be used. Changes in Pco₂ are labeled respiratory and changes in BE and SHCO₃^- are labeled metabolic.

6.1 Respiratory acidosis

Respiratory acidosis is defined as an acidosis caused by an increase in Pco₂. When no significant compensatory response is present the condition is called acute respiratory acidosis and when there is a compensatory response, i.e. an increase in SBE, the condition is called chronic respiratory acidosis. Compensation is taken care of by the kidneys. They increase the bicarbonate reuptake to the blood and increase the proton liberation to the urine. A subsequent metabolic compensation takes several days to develop.

Among the causes are diseases of the respiratory muscles (Myasthenia, jmuscle fatigue or paralysis etc), lung diseases (edema, emphysema, bronchitis, severe asthma or pneumonia) and diseases causing inhibition of the medullary respiratory center (cardiac arrest, oxygen in chronic hypercapnia, drugs used in anaesthesia and central sleep apnea). |

Treatment is directed at the underlying cause. However, a plasma pH less than 7.20 or fatigue suggests respirator treatment, unless the cause can be rapidly deleted,

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6.2 Respiratory alkalosis

Respiratory alkalosis refers to alkalosis caused by a decrease in Pco₂. If no significant change in the metabolic component is present the condition is called acute alkalosis and if there is a decrease in SBE the condition is called chronic.

The kidneys respond to the alkalemia by decreasing the bicarbonate reabsorption and the proton excretion. Further, a moderate alkalemia will increase the lactic acid concentration in the plasma by some units. This is because of a reduced net uptake of lactic acid by the liver and an increased net production by other tissues.

Potential causes of respiratory alkalosis are hypoxemia, psychogenic or voluntary hyperventilation, neurologic disorders such as cerebrovascular accidents and pontine tumors, and salicylate intoxication. Respiratory alkalosis may also occur in patients with pulmonary diseases such as acute pneumonia, bronchial asthma etc., and almost certainly arises from stimulation of vagal afferent fibers within the lungs (animal studies have shown that vagal blockade can present hyperventilation) and from the concomitant hypoxia, which is a stimulus to peripheral and central chemoreceptors. However, the hypoxemia seen in patients with pneumonia, pulmonary oedema, asthma etc, seems to contribute very little to the excessive ventilation - clinical experience shows that administrating oxygen to these patients can correct the arterial hypoxemia but does not restore the Pco₂ to normal. On the other hand/ in the high altitude, hypoxia is indeed the primary stimulus to the hyperventilation.

Treatment is directed at the underlying cause.

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6.3 Metabolic acidosis

Metabolic acidosis is defined as an acidosis caused by reasons other than Pco₂, i.e. a process resulting in a decrease in SBE. Metabolic acidosis can be seen for instance in renal failure, diabetic ketoacidosis and severe tissue hypoxia. When no significant respiratory compensation is present the condition is called acute metabolic acidosis and if there is a

respiratory response the condition is called chronic metabolic acidosis.

Acidemia causes an increased ventilation because of the increased ammount of protons acting on the peripheral carotid and aortic chemoreceptors and on the central chemo-receptors. Respiratory compensation may be predicted by several empiric rules. The following equations may be used:

expected Pco₂ (kPa) = 0.2 x [HCO₃^-l] + 1 ± 0.3

expected Pco₂ (kPa) = 13 x (the decimal part of the pH value).

Maximal compensation will reduce Pco₂ to about 2 kPa. A Pco₂ that is higher than predicted indicates inadequate respiratory compensation. If a patient with metabolic acidosis has a Pco₂ that is lower than predicted, a primary disturbance causing hyperventilation is also present.

The anion gap and the chloride concentration are used to restrict the diagnostic possibilities. In Tab. 1 some causes of normal and elevated anion gap metabolic acidosis are listed.

Treatment is directed at the underlying cause. Base infusion may be appropriate in severe cases. The initial therapeutic goal is to raise the plasma pH to about 7.20, a level at which arrhythmias become less likely and cardiac contractility and responsiveness to catecholamines will be restored. The amount of base that should initially be given can be estimated by the formula below:

base to initially be given (mM) = Base deficit (mM) • 0.1 • body weight (kg).

When giving the base follow the electrolytes in plasma, particularly K⁺.

The most clinically common bases are bicarbonate, THAM, and Tribonate^®. Each has its advantages and drawbacks. Discussing this is beyond the scope of this text. However, it should be pointed out that a rapid infusion of bicarbonate i.v. will cause an increase in the Pco₂. Because carbon dioxide diffuses easily through cell membranes the acute effect will be an intracellular acidosis - although a base has been given. On the other hand if THAM is given rapidly, the Pco₂ will decrease, causing hypoventilation. Tribonate^®, which is a mixture of the sodium salts of bicarbonate, THAM, phosphate and acetate/ given to healthy spontaneously breathing persons does not significantly affect the Pco₂. However, when given to patients with inadequate respiration an increase in Pco₂ can be seen.

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Tab. 1. Some causes of metabolic acidosis.

Elevated anion gap acidosis	Normal anion gap acidosis
Ketoacidois	Renal tubular acidosis
Lactic acidosis	Minaralcorticoid deficiency
Salicylate intoxication	Diarrhoea
Methanol poisoning	Some cases of ketoacidosis, particulary duning insulin treatment
Renal failure

6.4 Metabolic alkalosis

Metabolic alkalosis is defined as an alkalosis due to causes other than Pco₂, i.e. a process resulting in an increase in SBE. Metabolic alkalosis can for instance be caused by:

1. Loss of protons from the stomach through vomiting or gastric suction.

2. Loss of K⁺. A low plasma [K⁺] causes a shift of protons into the intracellular compartment and increased release of protons by the kidneys.

3. Severe dehydration and thus hypovolemia. Hypovolemia is considered to produce an alkalosis only when the fluid lost contains an excess of protons or Cl^-in relation to HCO₃^-. Bleeding for instance will not cause any alkalosis, since the loss of Cl^- and HCO₃^- is in concentrations similar to those in plasma. Moreover, hypochloremia will increase the aldosterone production. The aldosterone effect upon the kidneys will preserve sodium and water but at the expense of potassium and protons.

4. Chloride depletion, often induced by diuretics or vomiting. A low plasma [Cl^-] promotes increased H⁺ secretion by the kidneys.

Depending on whether or not there is a signifikant compensation/ the alkalosis is labeled chronic or acute.

The respiratory response to metabolic alkalosis is variable and thus less accurately predictable than compensation for metabolic acidosis. However, Pco₂ should certainly be increased above the normal range. If Pco₂is normal or less than normal a respiratory alkalosis is also present Also

if the Pco₂ is above 7 kPa and certainly if greater than 8 kPa, there is a primary respiratory disturbance present (see next section).

Treatment is directed at the underlying cause. However, a plasma pH >7.55 is coupled to an increased mortality and should be counteracted. If the liver is functioning appropriately, ammonium chloride can be given i.v. To adults about 100 mmol is generally given the first 2 h. Follow the plasma electrolytes, particularly K⁺.

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6.5 Mixed acid-base disturbances

If there is a significant primary change in both the respiratory and the metabolic component, there is a mixed acid base disturbance For instance if a metabolic disturbance does not elicit appropriate respiratory compensation a mixed acid-base disturbance exists. Also if the respiratory response is greater than appropriate it can be deduced that there is a mixed disturbance too. The patient history and clinical status help the

clinician to differentiate between a mixed disturbance and a primary disturbance with compensation.

Treatment is directed at the underlying causes.