Chapter 70 – Acid–Base Physiology




Abstract




The word ‘acid’ is derived from the Latin acidus, meaning sour. Early chemists defined an acid as a chemical substance whose aqueous solution tastes sour, changes the colour of litmus paper to red and reacts with certain metals to produce the flammable gas, hydrogen. Likewise, a base is a chemical substance whose aqueous solution tastes bitter, changes the colour of litmus paper to blue and reacts with acids to produce a salt.





Chapter 70 Acid–Base Physiology




What is an acid?


The word ‘acid’ is derived from the Latin acidus, meaning sour. Early chemists defined an acid as a chemical substance whose aqueous solution tastes sour, changes the colour of litmus paper to red and reacts with certain metals to produce the flammable gas, hydrogen. Likewise, a base is a chemical substance whose aqueous solution tastes bitter, changes the colour of litmus paper to blue and reacts with acids to produce a salt.



What are the Brønsted–Lowry definitions of an acid and base?


Brønsted and Lowry independently recognised that acid–base reactions in aqueous solution involve the transfer of an H+ from one molecule to another, and they suggested the following definitions:




  • An acid is a proton donor.



  • A base is a proton acceptor.


The generic reaction between an acid and base is:



HA + B ⇌ BH+ + A
HA+B⇌BH++A–

where HA is a Brønsted–Lowry acid (as it donates H+), B is a Brønsted–Lowry base (as it accepts H+), BH+ is referred to as the conjugate acid and A‾ is referred to as the conjugate base.


Acids may be classified as being either strong or weak:




  • A strong acid is one that completely dissociates in solution.



  • A weak acid is one that only partially dissociates in solution.



What is pH?


pH is a measure of the acidity of an aqueous solution. pH is the negative decadic logarithm of the H+ ion concentration:




Key equation: pH




pH =  − log10[H+]
pH=−log10H+


where log10 is the logarithm (base 10) and [H+] is the molar concentration of H+ ions.


Note: pH is dimensionless; that is, it has no units.


Because the pH scale is logarithmic, a small change in pH represents a much larger change in [H+]:




  • The ‘normal’ body pH of 7.4 is equivalent to an H+ concentration of 40 nmol/L.



  • Acidaemia is defined as an arterial pH below 7.35.



  • Alkalaemia is defined as an arterial pH above 7.45.



  • A small reduction in pH from 7.4 to 7.0 represents more than a doubling of the H+ concentration, from 40 to 100 nmol/L.



What is Ka?


Ka is the ionisation constant for H+ from its acid in the equilibrium:


HAk1⇌k2H++A−

where k1 is the rate constant for the forward reaction and k2 is the rate constant for the backward reaction.


When the rate of the forward reaction equals the rate of the backward reaction, the reaction is said to be at equilibrium. The equilibrium constant Ka can then be written as:


Ka=k1k2=H+A−HA


What is pKa?


pKa is defined as the negative decadic logarithm of the ionisation constant (Ka) of an acid. It equals the pH value at which equal concentrations of the acid and conjugate base forms of a substance are present.




Key equation: pKa




pKa =  − log10Ka
pKa=−log10Ka


pKa is a measure of the strength of an acid. It is normally used to characterise weak acids.


From the equilibrium equation defining Ka above, it can be seen that:




  • A high Ka represents greater dissociation of HA into H+ and A‾ and therefore a greater concentration of free H+. A low pKa therefore corresponds to increased acidity.



  • A low Ka represents less dissociation of HA, resulting in a lower concentration of free H+. A high pKa therefore corresponds to reduced acidity.


Like pH, pKa is a logarithmic scale. Therefore, a small reduction in pKa represents a much larger increase in acidity.


The acidity of a substance in solution can be related to the pH in a more formal way. Rearranging the above equation:


H+=KaHAA−

As pH = –log10[H+]:


pH=−log10KaHAA−

Multiplying out the brackets:


pH=−log10Ka−log10HAA−

And as pKa = –log10Ka:




Key equation: the Henderson–Hasselbalch equation



pH=pKa+log10A−HA


or:


pH=pKa+log10Conjugate baseAcid


Illustrate these principles of acid–base balance in the HCO3‾/H2CO3 buffer system


The most important physiological buffering system is that of CO2, H2CO3 and HCO3‾, which follows the reaction:



CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3
CO2+H2O⇌H2CO3⇌H++HCO3−

H2CO3 is the Brønsted–Lowry acid, water is the Brønsted–Lowry base, HCO3‾ is the conjugate base and H3O+ is the conjugate acid.


Applying the Henderson–Hasselbalch equation to the HCO3‾/H2CO3 buffer system:


pH=pKa+log10HCO3−H2CO3

As the pKa of the HCO3‾/H2CO3 equilibrium is 6.1 and [H2CO3] can be related to the solubility and partial pressure of CO2 (PaCO2), this equation can be rewritten as:


pH=6.1+log10HCO3−0.23×PaCO2

where 0.23 is a solubility factor. PaCO2 is measured in kilopascals.


Normal plasma pH can therefore be predicted by inserting the ‘normal’ plasma values of [HCO3‾] = 24 mmol/L and PaCO2 = 5.3 kPa:


pH=6.1+log10240.23×5.3=7.4


How are disorders of acid–base balance classified?


Acid–base disturbance is traditionally classified by pH disturbance (i.e. acidosis or alkalosis) and by aetiology (i.e. whether it is of respiratory or metabolic origin). Acids of respiratory origin – namely CO2 – are known as volatile acids, as they may escape as a gas. Acids that are non-volatile (e.g. lactic acid) are known as fixed acids as they may not escape the system.


The four classes of acid–base disorders are:




  • Respiratory acidosis, in which decreased A results in pH < 7.35 and PaCO2 > 6.0 kPa. Hypoventilation may be due to:




    1. Depression of the respiratory centre; for example, due to opioids or obesity hypoventilation syndrome;



    2. Nerve or muscle disorders, such as Guillain–Barré syndrome and myasthenia gravis;



    3. Chest wall disease; for example, flail chest;



    4. Airway disease; for example, asthma and chronic obstructive pulmonary disease (COPD);



    5. Lung parenchymal disease; for example, acute respiratory distress syndrome (ARDS).

    Of particular relevance to anaesthesia, hypercapnoeic acidosis may also occur due to:


    1. Insufficient mechanical ventilation, which may of course be intentional; for example, permissive hypercapnoea in patients with ARDS;



    2. Increased CO2 production in malignant hyperpyrexia;



    3. Exogenous CO2 intake; for example, re-breathing CO2-containing exhaled gases or insufflation of CO2 in laparoscopic surgery.

    If a respiratory acidosis persists for a period of days, the kidneys increase HCO3‾ reabsorption; this is termed metabolic compensation. Raised plasma HCO3‾ concentration (>26 mmol/L) may be seen in patients with COPD, ARDS and obesity hypoventilation syndrome.



  • Metabolic acidosis, in which there is an increase in fixed acid, which may be endogenous (e.g. lactic acid) or exogenous (e.g. salicylate). As the increased fixed acid is buffered by HCO3‾, metabolic acidosis is characterised by low plasma HCO3‾ concentration (<22 mmol/L) and pH < 7.35. Identification of the cause of metabolic acidosis may be aided by the anion gap (see below). The respiratory system responds to a metabolic acidosis by rapidly increasing A, thereby reducing PaCO2; this is referred to as respiratory compensation.



  • Respiratory alkalosis, in which hyperventilation results in hypocapnoea (PaCO2 < 4.7 kPa) and alkalosis (pH > 7.45). Increased A may be the result of:




    1. Central causes; for example, head injury, pain, anxiety, progesterone (in pregnancy) and drugs (such as salicylate overdose).



    2. Hypoxaemia, in which afferent signals from peripheral chemoreceptors stimulate the respiratory centre. This may occur, for example, at high altitude.



    3. Activation of lung J‑receptors, as occurs in pulmonary embolus and pulmonary oedema.



    4. Excessive mechanical ventilation.




  • Metabolic alkalosis – the least common of the main acid–base disorders – in which plasma HCO3‾ exceeds 26 mmol/L in the absence of a primary respiratory acidosis. The more common causes of metabolic alkalosis are:




    1. Gain of exogenous alkali; for example, an infusion of sodium bicarbonate and massive transfusion, where citrate is metabolised to HCO3‾;



    2. Loss of endogenous acid; for example, from the stomach through severe vomiting or nasogastric drainage or from the kidney through the use of diuretics.


Only gold members can continue reading. Log In or Register to continue

Sep 27, 2020 | Posted by in ANESTHESIA | Comments Off on Chapter 70 – Acid–Base Physiology
Premium Wordpress Themes by UFO Themes