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An Introduction to Acid Base Physiology

Created: 20/2/2007
Updated: 27/2/2007

Resource: Metabolic

An Introduction to Acid Base Physiology

Dr Andy Walden PhD
Advanced Trainee in Intensive Care Medicine


A good understanding of acid-base physiology is integral to the optimal management of critically ill patients. Acid-base physiology constitutes a vital part of the diagnostic process in critically ill patients and also forms an important part of the assessment of response to treatment.

Focus on the definitions of acid and base

In the 18th century, Lavoisier first attempted to define an acid as “a substance containing oxygen”. This was because the majority of the substances he was working on were oxyacids (e.g. H2SO4 and HNO3), and the existence of the hydrohalic acids (e.g. HCl) was not known. This definition of acids was brought into question following the discovery of the halides by Humphrey Davy in 1810.

In 1887, Arrhenius formulated the model of an acid as “a substance that dissociated in water to free hydrogen ions” and a base as “a substance that dissociated in water to produce hydroxyl ions”. However, these definitions were not entirely encompassing because not all acids contain hydrogen ions and not all bases hydroxyl ions. Moreover, these definitions of acid and base are only applicable to aqueous models.

The theory moved on in 1923, when two separate teams attempted to refine Arrhenius’s original definition. The Brønsted-Lowry concept suggested that an acid was “a substance that donated hydrogen ions to a conjugate base”. The advantage of this definition was that it encompassed non-aqueous systems.

At approximately the same time, Lewis defined an electronic system, where “an acid is any compound that is able to accept an electron pair and a base is any compound that is a potential electron pair donor”. This model addressed the problem that not all acids contained hydrogen ions.

In 1939, the Russian chemist Usanovich proposed an even broader definition of acids and bases. According to Usanovich, “an acid is any substance that accepts negative species and donates positive ones” and “a base any substance that accepts positive species and donates negative ones”.

In the current field of acid-base physiology, a combination of these theories is often used interchangeably. A summary of the acid base theories is given in Table 1.

 Arrhenius  Aqueous solutions
Acid = H+ in solution
Base = OH- in solution
At neutrality [H+ ] = [OH- ]
 Bronsted-Lowry  Acid = H+ donor
Base = H+ acceptor
 Lewis  Acid = electron pair acceptor
Base = electron pair donor
 Usanovich  Acid = a substance that donates an anion or accepts a cation
Base = reverse of acid

Table 1. Summary of different theories of what constitutes an acid and base. [H+] is the hydrogen ion concentration. [OH-] is the hydroxyl ion concentration.

Focus on pH

The pH is a measurement without units and represents the activity of hydrogen ions within a given solution. The pH concept was brought about by SPL Sorensen in 1909. Although it lacks units, the pH scale it is not an arbitrary scale. It represents the reverse logarithmic scale of hydrogen ion activity. The fact that the scale is logarithmic (log10) means that for each reduction by one in pH, there is a 10-fold increase in the hydrogen ion activity:

pH = -log10 (aH+)

Where aH+ is the activity of the hydrogen ions and is dimensionless. It follows that, in theory, the pH could range from infinity to infinity. Indeed, concentrated hydrochloric acid used by chemists has a pH of –1.1. Within biological systems, the normal pH range is 0 to 14, and under these circumstances aH+ approximates to the hydrogen ion concentration ([H+ ]). Thus:

pH ≈ -log10 [H+]

Using this approximation, at a pH of 7 at 25ºC the solution is neutral or [H+] = [OH-].

Normal arterial pH runs between values of 7.35–7.45. However, the capacity of the human body is such that it can actually survive over a 10-fold range of pH activity from 6.8 to 7.8.

Focus on pH measurement

pH can be measured with litmus paper (e.g. historically used to confirm the acidity of stomach contents following nasogastric tubeinsertion) or in a more complex and quantitatively precise manner (e.g. blood gas analysis). The most commonly used pH meter is the glass electrode. This commonly takes the form of a glass tube ended with a small glass bubble containing a silver/silver chloride electrode and silver chloride solution (Figure 1). The straight part of the electrode is usually made of thick glass but the bubble is made as thin as is possible. When the solution is applied to both sides of the bubble, the H+ will collect on both sides of the glass, and the subsequent difference in charge between the two sides results in a potential difference. This potential difference is described by the Nernst equation and is directly proportional to the pH difference across the glass membrane.

Figure 1. A standard pH electrode utilising a silver electrode and silver chloride solution with a thin glass bubble at the end.

Blood gas analysers will analyse a sample at 37ºC, but pH is temperature dependent (pH rises as the temperature increases). The Rosenthal correction compensates for the patient’s temperature:

ΔpH = 0.015pH/Δt(ºC)

The change in pH (ΔpH) equals 0.015 times the pH per ºC change in temperature (t)

This clearly has implications in critically ill patients with abnormal temperature homeostasis. The pH in septic patients with apyrexia will be increased, whereas in a hypothermic patient it will appear decreased.

Focus on the physiological importance of pH

The pH has two major physiological effects. Firstly, the pH will affect the ionisation status of most molecules within the physiological range and is responsible for certain molecules remaining trapped within the cell. Specifically, phosphates, ammonium and carboxylic acid molecules are all trapped within the cell as a consequence of their ionisation status at physiological pH. This is known as the Davis hypothesis.

Secondly, pH has important effects on the body’s protein and enzyme systems. This is predominantly an intracellular function. The majority of the body’s enzyme systems work optimally at physiological pH (7.35-7.45). Figure 2 shows the standard effect of pH on an enzyme system. However, certain enzyme systems work optimally at non-physiological pH. For example, the digestive enzyme pepsin is secreted into the stomach and is optimal at pH 1.5-3.0. Amylase, which is secreted both by salivary glands and glands in the small bowel, has an optimal pH range 8-10. Clearly, these enzymes are designed to work outside the cell.

Figure 2. pH dependence of enzyme reaction. Every enzyme has an optimal pH, and activity falls off either side of this optimal value.

Key reference

Sirker AA, Rhodes A, Grounds RM, Bennett ED. Acid-base physiology: the 'traditional' and the 'modern' approaches. Anaesthesia 2002; 57: 348-356.

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