Strong acids and bases are compounds that are completely dissociated in water for practical purposes. Under normal circumstances, this means that the concentration of hydrogen ions in an acid solution may be equal to the concentration of acid. The pH value is then equal to minus the logarithm of the concentration value. Hydrochloric acid (HCl) is an example of a strong acid. The pH of a 0.01M-HCl solution is −log10(0.01), i.e. pH = 2. Sodium hydroxide, NaOH, is a strong base example. The p[OH] value of a 0.01M NaOH solution is −log10(0.01), i.e. p[OH] = 2. According to the definition of p[OH] in the pOH section above, this means that the pH is equal to about 12. For sodium hydroxide solutions at higher concentrations, the auto-ionization equilibrium must be taken into account.

This was Sørensen`s original definition in 1909,[16] which was replaced in 1924 in favor of pH. [H] is the concentration of hydrogen ions, called [H+] in modern chemistry, which appears to have concentration units. More precisely, the thermodynamic activity of H+ in dilute solution must be replaced by [H+]/c0, where the concentration in the standard state is c0 = 1 mol/L. This ratio is a pure number whose logarithm can be defined. This definition was adopted because the ion-selective electrodes used to measure pH respond to activity. Ideally, the electrode potential E follows the Nernst equation, which can be written for the hydrogen ion as follows: There are many uncertainties in the operational definition of soil pH in the above way. Since a difference in electrical potential is measured between the glass and the reference electrodes, the activity of H+ is actually quantified and not the concentration. H+ activity is sometimes referred to as “effective H+ concentration” and is directly related to the chemical potential of the proton and its ability to perform chemical and electrical work in soil solution in equilibrium with solid phases. [28] Clay and organic matter particles carry a negative charge on their surface, and H+ ions attracted to them are in equilibrium with H+ ions in the soil solution.

The measured pH is by definition quantified only in the aqueous phase, but the value obtained is affected by the presence and nature of soil colloids and the ionic strength of the aqueous phase. A change in the water-soil ratio in manure can alter pH by disturbing the water-colloidal balance, particularly ionic strength. Using 0.01 M CaCl2 instead of water avoids this water-to-soil ratio effect and gives a more consistent approximation of “soil pH”, which refers to plant root growth, rhizosphere and microbial activity, acidity of drainage water, and chemical processes in the soil. The use of 0.01 M CaCl2 brings all soluble ions in the aqueous phase closer to colloidal surfaces and makes it possible to measure the H+ activity as close as possible to them. The use of the 0.01 M CaCl2 solution thus allows a more consistent quantitative estimation of H+ activity, especially when different soil samples are compared spatially and temporally. As part of its operational definition of the pH scale, IUPAC defines a set of buffer solutions over a range of pH values (often referred to as NBS or NIST). These solutions have a relatively low ionic strength (≈0.1) compared to seawater (≈0.7) and are therefore not recommended for characterizing seawater pH, as differences in ionic strength cause changes in the electrode potential. To solve this problem, an alternative series of tampons based on artificial seawater has been developed.

[32] This new series solves the problem of ionic strength differences between samples and buffers, and the new pH scale is called the “global scale”, often referred to as pHT. The overall scale was defined using a medium containing sulphate ions. These ions undergo protonation, H++SO2−4 HSO−4↔, so the global scale includes the action of protons (free hydrogen ions) and hydrogen sulfate ions: In 2010, a new “uniform absolute pH scale” was proposed, which would allow different pH ranges in different solutions to use a common reference standard for protons. It was developed on the basis of the absolute chemical potential of the proton. This model uses Lewis` acid-base definition. This scale applies to liquids, gases and even solids. [24] The concept of pH was first introduced in 1909 by Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory[5] and revised to modern pH in 1924 to allow definitions and measurements related to electrochemical cells.