Medical Biochemistry at a Glance

J. G. Salway

• Publisher: Wiley-Blackwell
• Year: 2011
• ISBN:
  9781118292402
• Edition: 3
• Topic: Medicine
• Subtopic: Medical Theory, Practice & Reference

Offering a concise, illustrated summary of biochemistry and its relevance to clinical medicine, Medical Biochemistry at a Glance is intended for students of medicine and the biomedical sciences such as nutrition, biochemistry, sports science, medical laboratory sciences, physiotherapy, pharmacy, physiology, pharmacology, genetics and veterinary science. It also provides a succinct review and reference for medical practitioners and biomedical scientists who need to quickly refresh their knowledge of medical biochemistry.

The book is designed as a revision guide for students preparing for examinations and contains topics that have been identified as 'high-yield' facts for the United States Medical Licensing Examination (USMLE), Step 1.

This third edition:

• Has been thoroughly revised and updated and is now in full colour throughout
• Is written by the author of the hugely successful Metabolism at a Glance (ISBN 9781405107167)
• Features updated and improved clinical correlates
• Expands its coverage with a new section on Molecular Biology
• Includes a brand new companion website of self-assessment questions and answers at www.ataglanceseries.com/medicalbiochemistry

Part 1: Acids, bases and pH

1

Acids, bases and hydrogen ions (protons)

Definition of PH

pH is defined as “the negative logarithm to the base 10 of the hydrogen ion concentration”,

For example, at pH 7.0, the hydrogen ion concentration is 0.000 000 1 mmoles/litre or 10⁻⁷ mmol/l.

Therefore, the negative log₁₀ is −(−7.0), i.e. +7.0 and hence the pH is 7.0.

What Is PH?

pH is “the “power of hydrogen”. It represents “the negative logarithm₁₀ of the hydrogen ion concentration”. So why make things so complicated: why not use the plain and simple “hydrogen ion concentration”? Well, the concept was invented by a chemist for chemists and has advantages in chemistry laboratories. In clinical practice we are concerned with arterial values between pH 6.9 and 7.9. However, chemists need to span the entire range of pH values from pH 1 to pH 14. Values in terms of pH enable a convenient compression of numbers compared with the alternative which would be extremely wide-ranging as shown in Fig. 1.3. Figure 1.6 shows the normal reference range for pH in blood and, in extremis, fatal ranges that may be seen in acidotic or alkalotic diseases.

Figure 1.1 Revision of logarithms.
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Figure 1.2 Examples of numbers and their logarithms.
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Figure 1.3 Understanding units.
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Figure 1.4 Brønsted and Lowry definition of acids and bases.
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Figure 1.5 pH and equivalent values.
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Figure 1.6 Examples of pH values seen in clinical practice.
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The PH Scale Is Not Linear

”The patient’s blood pH has changed by 0.3 pH unit” means it has doubled (or halved) in value.

It is sometimes stated that “the patient’s arterial blood pH has increased/decreased by, for example, 0.2 pH unit”. However, notice that because of the logarithmic scale, this can misrepresent the true change in traditional concentration units. For example, a fall of 0.2 pH units from pH 7.20 to pH 7.00 represents 37 nmol/l, whereas a decrease from pH 7.00 to pH 6.8 represents a change of 60 nmol/l.

Also note that because the log₁₀ of 2 = 0.3 (that is 2 = 10^0.3), a decrease in pH by 0.3, e.g. from pH 7.40 to pH 7.10, represents a two-fold increase in H⁺ concentration, i.e. from 40 nmol/l to 80 nmol/l. Similarly, an increase in pH from pH 7.40 to pH 7.70 represents a fall in H⁺ concentration from 40 nmol/l to 20 nmol/l.

The Henderson–Hasselbalch Equation

A weak acid dissociates as shown:

where HB is the weak acid that dissociates to a proton H⁺ and its conjugate base B⁻. NB Traditionally authors refer to the conjugate base as “A⁻”, i.e. the initial letter of acid, which is perhaps confusing.

Therefo...