Acid Base Balance (page 1)

Issue 13 (2001) Article 12: Page 1 of 3

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Acid Base Balance
Dr Stephen Drage & Dr Douglas Wilkinson,
Oxford, England

Key to Terms Used		Respiratory Acidosis
Introduction		Respiratory Alkalosis
The Hydrogen Ion and pH		Metabolic Acidosis
Acids, Bases and Buffers		Metabolic Alkalosis
The Importanceof Hydrogen Ion Concentration		Compensation
Production of Hydrogen Ions		Interpretation of Acid Base Disturbances in Blood Gas Results
Control of Hydrogen Ion Concentration		Summary
Disorders of Hydrogen Ion Homeostasis

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 O2 - the reduced state is when it is not combined with O2
Aerobic metabolism	Metabolic process using oxygen from the air
Anaerobic metabolism	Metabolic process without oxygen - often less efficient and used for short periods

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.

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.

Table 1: pH and Hydrogen ion concentration
pH	[H⁺] nanomol/l
6.0	1000
7.0	100
8.0	10
9.0	1

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

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 "proton donors". 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 "proton acceptors".

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 portion of the curve a large volume of acid is added with little change in [H⁺] or pH.

As we shall see later buffers are crucial in maintaining hydrogen ions within a narrow range concentrations in the body.

The Importance of Hydrogen Ion Concentration

Hydrogen ion concentration has a widespread effect on the function of the body's enzyme systems. The hydrogen ion is highly reactive and will combine with bases or negatively charged ions at very low concentrations. Proteins contain many negatively charged and basic groups within their structure. Thus, a change in pH will alter the degree ionization of a protein, which may in turn affect its functioning. At more extreme hydrogen ion concentrations a protein's structure may be completely disrupted (the protein is then said to be denatured).

Enzymes function optimally over a very narrow range of hydrogen ion concentrations. For most enzymes this optimum pH is close to the physiological range for plasma (pH= 7.35 to 7.45, or [H⁺]= 35 to 45nmol/l). Figure 5 shows a typical graph obtained when enzyme activity is plotted against pH. Notice that the curve is a narrow bell shape centred around physiological pH.

Although most enzymes function optimally around physiological pH it should be noted that a few enzymes function best at a much higher hydrogen ion concentration (ie: at a lower pH). The most notable of these enzymes is pepsin, which works best in the acid environment of the stomach - optimum pH 1.5-3 or [H⁺]= 3-30 million nanomol/l.

As enzymes have a huge number of functions around the body, an abnormal pH can result in disturbances in a wide range of body systems. Thus, disturbances in pH may result in abnormal respiratory and cardiac function, derangements in blood clotting and drug metabolism, to name but a few. From these few examples it is clear that the anaesthetist should strive to ensure that hydrogen ion concentration is maintained within the normal range.

(Continued ...)