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Introduction
The aim of this article is to provide the reader with a basic understanding of the physiology and biochemistry of acid base balance and its disturbances. This subject is often made unnecessarily complex and most disturbances of acid base control can be understood with the application of a few key principles.
The Hydrogen Ion and pH
The hydrogen ion consists of a single positively charged particle (the proton) that is not orbited by any electrons. The hydrogen ion is, therefore, the smallest ionic particle and is extremely reactive. It is this fact that accounts for its profound effect on the functioning of biological systems at very low concentrations.
In the environment hydrogen ion concentrations vary over an enormous scale (from less than 10-14 mol/l to more than 1mol/l).
The pH scale was developed in order to simplify (or perhaps further complicate!) the mathematics of handling such a large range of numbers. The pH is calculated by taking the negative logarithm of the hydrogen ion concentration, as shown below.
pH = -^ log10[H+]: where [H+] is the hydrogen ion concentration.
ACID BASE BALANCE
Dr Stephen Drage & Dr Douglas Wilkinson, Oxford, England
Key to terms used Nanomol (nmol) 1 x 10-9^ mol = 0.000000001 mol Ion Electrically charged particle formed when molecules are in solution Enzyme An organic substance which accelerates reactions Reduced state Some substances can combine reversibly with O 2 - the reduced state is when it is not combined with O 2 Aerobic metabolism Metabolic process using oxygen from the air Anaerobic metabolism Metabolic process without oxygen - often less efficient and used for short periods
Figure 1 gives an example of how pH is calculated.
Table 1 gives examples of pH values and corresponding hydrogen ion concentrations. It is important to note that an increase of one pH point results in a ten-fold decrease in hydrogen ion concentration.
Although pH terminology is widely used in textbooks and in biochemistry reports, it is important to realise that pH is merely a reflection of the hydrogen ion concentration. In the rest of this article both terms will be used to impress upon the reader that, essentially, they refer to the same thing.
Acids, Bases and Buffers
Acids: An acid is defined as any compound, which forms hydrogen ions in solution. For this reason acids are sometimes referred to as ìproton donorsî. To aid understanding of these concepts consider an imaginary acid with the chemical formula HA. In the first example in Figure 2, the acid dissociates (separates) into hydrogen ions and the conjugate base when in solution.
Bases: A base is a compound that combines with hydrogen ions in solution. Therefore, bases can be referred to as ìproton acceptorsî.
Strong Acids: A strong acid is a compound that ionizes completely in solution to form hydrogen ions and a base. Example 2 illustrates a strong acid in solution, where this dissociation is complete.
Weak Acids and Bases: these are compounds that are only partially ionised in solution. Example 3 shows a weak acid in solution with incomplete dissociation.
Buffers : A buffer is a compound that limits the change in hydrogen ion concentration (and so pH) when hydrogen ions are added or removed from the solution. It may be useful to think of the buffer as being like a
sponge. When hydrogen ions are in excess, the sponge mops up the extra ions. When in short supply the sponge can be squeezed out to release more hydrogen ions! All buffers are in fact weak acids or bases. Figure 3 shows how as hydrogen ions are added to a buffer solution they combine with A-^ (the conjugate base) and the reaction is pushed to the left. This creates more HA whilst removing the excess H +^ from the solution. Similarly, as hydrogen ions are removed from solution by addition of a strong base the reaction moves to the right restoring the hydrogen ion concentration and reducing the quantity of HA. The effects of buffers can also be illustrated graphically. If a strong acid is added slowly to a buffer solution and the hydrogen ion concentration [H+] is measured then a plot similar to the one in figure 4 will be generated. Notice that during the highlighted
Teaching point As a solution becomes more acidic or less alkaline, the pH falls (hydrogen ion concentration rises). The opposite happens when solutions become less acidic or more alkaline
Table 1: pH and Hydrogen ion concentration
pH [H +] nanomol/l
Production of Hydrogen Ions
The processes of metabolism generate hydrogen ions. Small amounts (40-80mmol/24h) are formed from the oxidation of amino acids and the anaerobic metabolism of glucose to lactic and pyruvic acid. Far more acid is produced as a result of carbon dioxide (CO 2 ) release from oxidative (aerobic) metabolism - 15,000mmol/ 24h (15x103 mmol/24h). Although CO 2 does not contain hydrogen ions it rapidly reacts with water to form carbonic acid (H 2 CO 3 ), which further dissociates into hydrogen and bicarbonate ions (HCO 3 - ). This reaction is shown below:
CO 2 + H 20 " H 2 CO 3 # HCO 3 -^ + H+
This reaction occurs throughout the body and in certain circumstances is speeded up by the enzyme carbonic anhydrase. Carbonic acid is a weak acid and with bicarbonate, its conjugate base, forms the most important buffering system in the body.
Acids or bases may also be ingested, however, it is uncommon for these to make a significant contribution to the bodyís hydrogen ion concentration, other than in deliberate overdose.
Control of Hydrogen Ion Concentration
With hydrogen ion concentration being so critical to enzyme function, the body has sophisticated mechanisms for ensuring pH remains in the normal range. Three systems are involved: blood and tissue buffering, excretion of CO 2 by the lungs and the renal excretion of H+^ and regeneration of HCO 3 -.
1. Buffers
As we have seen, buffers are able to limit changes in hydrogen ion concentration. This prevents the large quantities of hydrogen ions produced by metabolism resulting in dangerous changes in blood or tissue pH.
a) Bicarbonate This is the most important buffer system in the body. Although bicarbonate is not an efficient buffer at physiological pH its efficiency is improved because CO 2 is removed by the lungs and bicarbonate regenerated by the kidney. There are other buffers that act in a similar way to bicarbonate, for example: hydrogen phosphate (HPO 3 2-^ ), however, these are present in smaller concentrations in tissues and plasma.
b) Proteins As mentioned earlier many proteins, and notably albumin, contain weak acidic and basic groups within their structure. Therefore, plasma and other proteins form important buffering systems. Intracellular proteins limit pH changes within cells, whilst the protein matrix of bone can buffer large amounts of hydrogen ions in patients with chronic acidosis. c) Haemoglobin Haemoglobin (Hb) is not only important in the carriage of oxygen to the tissues but also in the transport of CO 2 and in buffering hydrogen ions (The physiology of oxygen delivery, Update in anaesthesia 1999; 10:8-14). Haemoglobin binds both CO 2 and H +^ and so is a powerful buffer. Deoxygenated haemoglobin has the strongest affinity for both CO 2 and H +^ ; thus, its buffering effect is strongest in the tissues. Little CO 2 is produced in red cells and so the CO 2 produced by the tissues passes easily into the cell down a concentration gradient. Carbon dioxide then either combines directly with haemoglobin or combines with water to form carbonic acid. The CO 2 that binds directly with haemoglobin combines reversibly with terminal amine groups on the haemoglobin molecule to form carbaminohaemoglobin. In the lungs the CO 2 is released and passes down its concentration gradient into the alveoli. The buffering of hydrogen ions formed from carbonic acid is more complicated. The chain of events that occurs within the red cell is most easily understood by referring to figure 6. In the tissues, dissolved CO 2 passes into the red blood cell down its concentration gradient where it combines with water to form carbonic acid. This reaction is catalysed by the enzyme carbonic anhydrase. Carbonic acid then dissociates into bicarbonate and hydrogen ions. The hydrogen ions bind to reduced haemoglobin to form HHb. Bicarbonate ions (HCO 3 -^ ) generated by this process pass back into the plasma in exchange for chloride ions (Cl-^ ). This ensures that there is no net loss or gain of negative ions by the red cell. In the lungs this process is reversed and hydrogen ions bound to haemoglobin recombine with bicarbonate to form CO 2 which passes into the alveoli. In addition, reduced Hb is reformed to return to the tissues.
2. Carbon Dioxide Elimination
As mentioned earlier CO 2 is responsible for the majority of hydrogen ions produced by metabolism. Therefore, the respiratory system forms the single most important organ system involved in the control of hydrogen ions. From respiratory physiology it should be remembered that the arterial partial pressure of CO 2 (PaCO 2 ) is inversely proportional to alveolar ventilation (ie: if alveolar ventilation falls the PaCO 2 rises). Therefore, relatively small changes in ventilation can have a profound effect on hydrogen ion concentration and pH. An acute rise in PCO 2 of 1Kilopascal (kPa) results in a 5.5nmol/l rise in the hydrogen ion concentration (resulting in a fall in plasma pH from 7.4 to 7.34).
The importance of PaCO 2 and hydrogen ion concentration is underlined by the fact that the control of ventilation is brought about by the effect of CO 2 on cerebrospinal fluid (CSF) pH. The detail of the control of breathing and elimination of CO 2 are beyond the remit of article but have been discussed in a previous issue of update (Update in anaesthesia 1999; 10:8-14)
3. Renal Handling of Bicarbonate and Hydrogen Ions
The kidneys not only secrete hydrogen ions but they also regenerate bicarbonate ions. The renal handling of electrolytes also influences acid base balance. All
aspects of renal involvement in acid base balance are interlinked, but for clarity are dealt with separately below. a) Regeneration of Bicarbonate: Bicarbonate ions are freely filtered by the glomerulus. The concentration of bicarbonate in the tubular fluid is equivalent to that of plasma. If bicarbonate were not reabsorbed the buffering capacity of the blood would rapidly be depleted.
The process of reabsorption of bicarbonate occurs mostly in the proximal convoluted tubule and is summarised in figure 7. Filtered bicarbonate combines with secreted hydrogen ions forming carbonic acid. Carbonic acid then dissociates to form CO 2 and water. This reaction is catalysed by carbonic anhydrase, which is present in the brush border of the renal tubular cells. This CO 2 readily crosses into the tubular cell down a concentration gradient. Inside the cell the CO 2 recombines with water, again under the influence of carbonic anhydrase, to form carbonic acid. The carbonic acid further dissociates to bicarbonate and hydrogen ions. The bicarbonate passes back into the blood stream whilst the H+^ passes back into the tubular fluid in exchange for sodium. In this way, virtually all the filtered bicarbonate is reabsorbed in the healthy individual.
(acidaemia) or too few hydrogen ions (alkalaemia). In other words, the pH is too low in acidaemia (less than 7.35) whilst in alkalaemia the pH is too high (more than 7.45). These disturbances may be due to respiratory causes (ie: changes in PaCO 2 ) or non- respiratory (metabolic) causes. When the cause of the acid base disturbance has been discovered, the words acidosis or alkalosis may be used in conjunction with the physiological cause of the disturbance (ie: respiratory acidosis, metabolic alkalosis etc). These are discussed in more detail below.
Respiratory Acidosis
This results when the PaCO 2 is above the upper limit of normal, >6kPa (45mmHg). The relationship between hydrogen ion concentration and CO 2 was discussed earlier (Production of Hydrogen Ions). Respiratory acidosis is most commonly due to decreased alveolar ventilation causing decreased excretion of CO 2. Less commonly it is due to excessive production of CO 2 by aerobic metabolism.
a) Inadequate CO 2 Excretion: the causes of decreased alveolar ventilation are numerous, they are summarised in Fig 10. Many of the
causes of decreased alveolar ventilation are of interest to theanaesthetist and many are under our control. b) Excess CO 2 Production: respiratory acidosis is rarely caused by excess production of CO 2. This may occur in syndromes such as malignant hyperpyrexia, though a metabolic acidosis usually predominates. More commonly, modest overproduction of CO 2 in the face of moderately depressed ventilation may result in acidosis. For example, in patients with severe lung disease a pyrexia or high carbohydrate diet may result in respiratory acidosis. Respiratory Alkalosis Results from the excessive excretion of CO 2 , and occurs when the PaCO 2 is less than 4.5kPa (34mmHg). This is commonly seen in hyperventilation due to anxiety states. In more serious disease states, such as severe asthma or moderate pulmonary embolism, respiratory alkalosis may occur. Here hypoxia, due to ventilation perfusion (V/Q) abnormalities, causes hyperventilation (in the spontaneously breathing patient). As V/Q abnormalities have little effect on the
excretion of CO 2 the patients tend to have a low arterial partial pressure of oxygen (PaO 2 ) and low PaCO 2.
Metabolic Acidosis
May result from either an excess of acid or reduced buffering capacity due to a low concentration of bicarbonate. Excess acid may occur due increased production of organic acids or, more rarely, ingestion of acidic compounds.
a) Excess H +^ Production: this is perhaps the commonest cause of metabolic acidosis and results from the excessive production of organic acids (usually lactic or pyruvic acid) as a result of anaerobic metabolism. This may result from local or global tissue hypoxia. Tissue hypoxia may occur in the following situations:
! Reduced arterial oxygen content: for example anaemia or reduced PaO 2.
! Hypoperfusion: this may be local or global. Any cause of reduced cardiac output may result in metabolic acidosis (eg: hypovolaemia, cardiogenic shock etc). Similarly, local hypoperfusion in conditions such as ischaemic bowel or an ischaemic limb may cause acidosis.
! Reduced ability to use oxygen as a substrate. In conditions such as severe sepsis and cyanide poisoning anaerobic metabolism occurs as a result of mitochondrial dysfunction.
Another form of metabolic acidosis is diabetic ketoacidosis. Cells are unable to use glucose to produce energy due to the lack of insulin. Fats form the major source of energy and result in the production of ketone bodies (aceto- acetate and 3- hydroxybutyrate) from acetyl coenzyme A. Hydrogen ions are released during the production of ketones resulting in the metabolic acidosis often observed.
b) Ingestion of Acids: this is an uncommon cause of metabolic acidosis and is usually the result of poisoning with agents such as ethylene glycol (antifreeze) or ammonium chloride. c) Inadequate Excretion of H+^ : this results from renal tubular dysfunction and usually occurs in conjunction with inadequate reabsorption of bicarbonate. Any form of renal failure may result in metabolic acidosis. There are also specific disorders of renal
hydrogen ion excretion known as the renal tubular acidoses. Some endocrine disturbance may also result in inadequate H+^ excretion e.g. hypoaldosteronism. Aldosterone regulates sodium reabsorption in the distal renal tubule. As sodium reabsorption and H+^ excretion are linked, a lack of aldosterone (eg: Addisonís disease) tends to result in reduced sodium reabsorption and, therefore, reduced ability to excrete H+^ into the tubule resulting in reduced H+^ loss. The potassium sparing diuretics may have a similar effect as they act as aldostrone antagonists. d) Excessive Loss of Bicarbonate: gastro- intestinal secretions are high in sodium bicarbonate. The loss of small bowel contents or excessive diarrhoea results in the loss of large amounts of bicarbonate resulting in metabolic acidosis. This may be seen in such conditions as Cholera or Crohnís disease. Acetazolamide, a carbonic anhydrase inhibitor, used in the treatment of acute mountain sickness and glaucoma, may cause excessive urinary bicarbonate losses. Inhibition of carbonic anhydrase slows the conversion of carbonic acid to CO 2 and water in the renal tubule. Thus, more carbonic acid is lost in the urine and bicarbonate is not reabsorbed. The importance of carbonic anhydrase in the reabsorption of bicarbonate was illustrated in Figure 7. Metabolic Alkalosis May result from the excessive loss of hydrogen ions, the excessive reabsorption of bicarbonate or the ingestion of alkalis. a) Excess H +^ loss: gastric secretions contain large quantities of hydrogen ions. Loss of gastric secretions, therefore, results in a metabolic alkalosis. This occurs in prolonged vomiting for example, pyloric stenosis or anorexia nervosa. b) Excessive Reabsorption of Bicarbonate: as discussed earlier bicarbonate and chloride concentrations are linked. If chloride concentration falls or chloride losses are excessive then bicarbonate will be reabsorbed to maintain electrical neutrality. Chloride may be lost from the gastro-intestinal tract, therefore, in prolonged vomiting it is not only the loss of hydrogen ions that results in the alkalosis but
PCO 2 alone causes an alkalaemia (high pH). The body is therefore using this mechanism to try to bring the low pH caused by the metabolic acidosis back towards normal.
By now the complexity of acid base disturbance should be clear!! As in many complex concepts examples may clarify matters. In the following examples work through the flow charts to interpret the data.
Example 1: A 70 year old man is admitted to the intensive car unit with acute pancreatitis. He is hypotensive, hypoxic and in acute renal failure. He has a respiratory rate of 50 breaths per minute. The following blood gas results are obtained:
pH 7.
PCO 2 3.0kPa (22mmHg)
BE -21.0mmol
From the flow charts: firstly, he has a severe acidaemia (pH 7.1). The PCO 2 is low, which does not account for the change in pH (a PCO 2 of 3.0 would tend to cause alkalaemia). Therefore, this cannot be a primary respiratory acidosis. The base excess of -21 confirms the diagnosis of a severe metabolic acidosis. The low PCO 2 indicates that there is a degree of respiratory
compensation due to hyperventilation. These results were to be expected given the history. Example 2: A 6 week old male child is admitted with a few days history of projectile vomiting. The following blood gases are obtained: pH 7. PCO 2 6.5kPa (48mmHg) BE +11.0mmol The history points to pyloric stenosis. There is an alkalaemia, which is not explained by the PCO 2. The positive base excess confirms the metabolic alkalosis. The raised PCO 2 indicates that there is some respiratory compensation. Summary This article has attempted to provide the reader with a grounding in acid base physiology and its importance in anaesthesia and intensive care. It is also hoped that the reader will be able to apply this knowledge in the interpretation of blood gas results. There is an excellent web page to learn more about blood gas results interpretation:- http://www.health.adelaide.edu.au/paed-anes/ javaman/Respiratory/AcidBase.html
Interpretation of Acid Base Disturbance
