Wilfred F. Widdas
Great Lodnon, Milton Keynes

https://doi.org/10.36866/pn.43.11

Assuming that physiologists will continue to have a dominant role in the preclinical teaching of medical students it is important that in their teaching and in their own research they give a clear and understandable scientific description of pH concepts. The subject of pH and hydrogen ion control has grown through the twentieth century and many of the earlier simplifications are scientifically incorrect. They should now be more explicitly described, in this, the twenty first century, so that clinical and scientific aspects can be understood in simple terms without ambiguity.

Scientifically, H+ is the nomenclature of the proton or nucleus of the hydrogen atom. The hydrogen atom is the foundation stone of chemistry and chemical measurements. For convenience the lightest of the elements (H) is given the mass of one Dalton(] Da). With square brackets [H+], the same symbol has been used medically and in many papers in the biological sciences to mean _the hydrogen ion concentration in watery solutions or blood plasma. This is a misnomer, since free protons only exist in a rarefied gas in a high technology machine used in nuclear physics. This misnomer is further compromised and compounded by such phrases as “The Na+/H+ Exchanger”, the concept of which is of paramount importance in physiology and pH homeostasis.

The confusion in the phrase ” Na+/H+ Exchanger” lies in the conveyance of a hypothetical concept that the hydrogen ions (literally protons) have come from the cell cytoplasm beyond the membrane. This is often the firm understanding of junior physiologists although there is no experimental evidence for this interpretation. It is very simple to show that each hydrogen ion (H30+) in an average cell of ca 150 mM of monovalent ions such as KCl would be in a relevantly huge environment of positive and negative ions. In an acid solution of pH 6.2 each hydrogen ion would be surrounded by one million positively charged potassium ions and they could not migrate to the membrane protein “exchangers” through this maze of charged ionic species without contravening thermodynamic principles. The thermodynamic principle in this case is given in the second law of thermodynamics, which states that the entropy (random distribution) of an isolated system increases in a spontaneous change.

The earliest investigators of this sodium­ dependent alkalinisation of cells were careful to point out that the entry of hydroxyl ions would be equivalent to the removal of hydrogen ions from the cell interior. Thus, Murer et al. (1) in 1976 described their findings as a “sodium I proton antiport”. Terms such as NHE 2 or NHE 3 refer to a definite membrane protein, of which, the amino acid sequence has been determined and none of these terms create confusion. All that the experimental evidence clearly shows is that sodium ions enter the cell and that the external solution becomes acidic by an amount that exactly matches the sodium entry. This effect at the membrane surface would also apply if the membrane protein was able to split water into a hydroxyl and a hydrogen ion and that only neutral NaOH could penetrate the membrane and enter into the cell.

The random distribution of hydroxyl ions from the membrane proteins into the cytoplasm of a cell would produce an increase in chemical entropy and would not contravene thermodynamic principles. Further, the random distribution of hydroxyl ions throughout the cytoplasm would raise the pH value as measured by intracellular microelectrodes or fluorescent dyes. In short there is no magic in a sodium-dependent alkalinisation of cells. The ambiguity and confusion induced by the nomenclature is quite unnecessary.

There is still an ambiguity regarding the use of H+ for protons and [H+] for hydrogen ions. Both may exist in biological systems and it would be helpful if they were clearly differentiated in publications. The proton or hydrogen nucleus can be regarded as always being in a relatively much larger environment provided by a pair of electrons belonging to a larger element of the periodic table of chemical elements. Medical students would obtain an adequately understanding if taught that the two most important elements were oxygen and nitrogen when incorporated into carbon compounds of hydrophobic proteins or other carbon compounds soluble rn water. A simplistic view of considerable practical value would be to include a third category of such compounds which have an intermediate set of prope1ties, which are avid for seeking out the interface between planar water surfaces and hydrophobic proteins, including those proteins embedded in the hydrophobic lipid membranes surrounding cells. Many drugs and pharmacological agents seek the water/protein interfaces in such membrane proteins with clefts that can be filled by water.

Protons (H+) certainly exist in the electron she) Is of the elements oxygen and nitrogen as part of many hydrophobic proteins. The protons may travel by exchange (often via hydrogen bonds) through adjacent molecular groupings. Biochemical oxidation-reduction systems involve primarily the transfers of electrons (again often eventually in pairs) and protons can also transfer with the pairs. Where oxidative metabolism is driving the flow of electrons we have purely descriptive terms for hypothetical concepts like “proton pumps” proposed to explain gastric acid secretion. However, the acid in the stomach is an accompaniment of chloride and the hydrogen ion in water, as part of hydrochloric acid, is not the proton H+ but the hydronium ion H30+ so a further stage is required for a complete understanding of this mechanism.

In water the hydrogen ion, as understood by clinicians, is regarded as being hydrated but this concept does not clarify the fact that it is a tiny positive charge in a relatively enormous electron orbit of a larger element. In the case of water it is one of the oxygen atom’s lone pairs of electrons, which houses the third proton of the hydrogen ion. While the two occupied electron shells of water are roughly at right angles to one another (to give water its permanent dipole property) the third occupied electronic shell of the hydronium ion H30+ is probably at right angles to the other two. Thus, for a simple mental picture, the hydronium may be visualized as being like an old badminton shuttle with only three feathers left. However, the mass of the oxygen (inherently electro-negative) head will tend to lead the three hydrogen nuclei, which are sharing the full positive charge. Thus, when a strong acid reacts with bicarbonate, it is just as likely for the oxygen head to combine with the delta­electropositive central carbon of bicarbonate to form ortho-carbonic acid, as it is for one of the protons to form carbonic acid. There is no firm experimental evidence for either of these interpretations but the resolution of this ambiguity should be possible in future.

In 1909, Sorensen (2) advocated that the negative of the logarithm of the hydrogen ion concentration would be a more useful measure of the relative ionic dissociation of the water molecules in enzyme studies. For buffer systems the Henderson-Hasselbalch Equation has been very useful in medicine and physiology. It should, however, be realized that pH values have no arithmetical relation to the hydrogen ion concentration but the terminology has provided a convenient expression for logarithmic concentration terms normally in the range of 10(-4) 10(-10) (pH 4 to 10). In general, the logarithmic terms need careful scientific consideration; since they are not readily handled by the calculus or by simple polynomial numerical analysis. In 1981, Stewart (3) published a book on pH for biology and medicine. One of his main recommendations was that more advances in the understanding of pH problems in biological sciences would come from working directly rn hydrogen ion concentrations using nanomolar quantities as units. This is recommended to be the editorial policy for physiology in the twenty first century.

The same year, 1981, saw the publication of the Society’s excellent study guide on “Acid-base balance” edited by R. Hainsworth. In that guide, R. C. Thomas used nanomolar quantities when discussing buffering and pH control. Although up-to-date at the time of publication the subject has made several important advances since then. At recent Annual Meetings of the Society the revision of the nomenclature has been adumbrated. The Press Secretary of the Editorial Board has recently outlined the challenges and dangers facing the Journal in view of the spread of “free electronic access” ideas into scientific publications. Another possible danger lies in continuing to publish expositions and papers expressed in terms of outmoded science. The present object of achieving the highest quality of science may need the adoption of a terminology policy in pH homeostasis, which eliminates ambiguity and confusion.

Holding these views, I was asked to submit an article expressing them in the Magazine for consideration by the members. However, in drawing attention to a number of aspects at present considered unsatisfactory, or of controversial ambiguity, I recognise that I hold maverick views and have no wish to press them, other than for wider consideration. Perhaps a joint committee, with a wider coverage of chemists, biochemists and other biologists should be set up. Any such committee should include clinicians that are faced directly with such acid­base problems in treating patients and in teaching the concepts of pH control to junior doctors in teaching hospitals. Much of the teaching could be simplified in so far, as is necessary for medical scientists. The introductory teaching and clear explanation of the important acid-base concepts and their control may justifiably remain one of the basic responsibilities of physiologists.

References

1. Murer, H. et al., (1976) Biochemical J. 154 597-604.
2. Sorensen, S.P.L (1909) Bioch. Zs. 21 130-304.
3. Stewart, P.A. (1981) How lo understand acid-base: A quantitative acid-base primer for biology and medicine. Edward Arnold Ltd, London.