Acid-base balance and buffer substances

Henderson (1913) worked through one case in The Fitness of the Environment. He took 1 kilogram of carbon dioxide in 100 liters of water and added sodium hydrate in portions. At the start the hydrogen ion concentration was about 0.0001 N, almost exactly one thousand times that at neutrality¹. He also noted that the chief actions of hydroxyl ions are catalytic, with an effect proportional to their concentration². Bayliss (1915) stated that the two concentrations multiply to a constant in water at one temperature³. He applied this to a bicarbonate and carbon dioxide mixture at 18° and 42°. At 42° the dissociation constant of water had become 3.76 x 10^-14, so the hydroxyl ion concentration rose while the hydrogen ion concentration rose little³.

Haldane (1917) proposed that dissolved carbon dioxide excites the respiratory centre through its acidity, and that other acids would help it. He wrote that no means then existed to measure the slight alkalinity of the blood accurately. Adding standard acid until an indicator changed colour was rough and wrong in principle. The blood is only very slightly alkaline, yet a large amount of acid can be added before it turns acid⁴. The acid needed therefore measures the buffer substances more than the alkalinity⁴. Haldane credited Henderson with the point that buffers allow much acid or alkali with no measurable change. He agreed, but argued that physiological disturbances are far more minute than chemical methods can detect⁵.

Loeb (1918) used hydrogen ion concentration to explain a muscle response that grows with repeated stimulation. He wrote that the first stimulations raise the concentration to a point where the effect becomes greater⁶. As the concentration rises further, he said, the response declines to zero. He offered this in place of an account of the muscle learning. Osterhout (1922) plotted the hydrogen ion concentration of sea water against added alkali. The concentration fell rapidly at first, then very slowly until the magnesium hydrate was all precipitated⁷. Crile (1926) wrote that the relation between hydrogen and hydroxyl ions stays within narrow limits throughout life. He judged it essential to his bipolar mechanism⁸.

Punjiya and colleagues (2019), including Levin, described pH change as an engineering problem in electroporation. In direct current electroporation, electrolysis occurs at the electrodes. Reduction of water at the cathode produces hydroxyl ions and a localized alkaline environment around the electrode⁹. In alternating current electroporation the electrodes swap roles as anode and cathode. This balances production of hydroxyl and hydrogen ions and eases any pH change at the electrodes⁹.