Hydrogen and hydroxyl ions (H+ and OH–) are extremely important in biology because they are the constituents of water and the basis of the properties of acidity and alkalinity.
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These influence all aspects of cell life. Degree of acidity is commonly expressed as pH, with a number representing the concentration of hydrogen ions, since it is ionized hydrogen that determines the immediate acidic activity of any solution. Alkalinity of solutions may be expressed as concentration of hydroxyl ions (pOH) (not hydroxyl groups attached to C).
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However, since concentrations of hydrogen ions and of hydroxyl ions are reciprocally related, as will be shown, degrees of either are commonly given in terms of pH only.
pH:
Since H ions determine acidity, acids or alkalies may be strong or weak, depending on their degree of dissociation. This is always a fixed value for any given electrolyte and is generally expressed as the dissociation constant.
Strong acids are those which, when dissolved in water, dissociate largely into positively charged hydrogen ions and negatively charged ions. For example, sulfuric acid dissociates into two hydrogen ions and a sulfate ion.
Weak acids like acetic or citric also dissociate, but to a lesser degree. The acidic activity of any acid solution depends upon the concentration of ions of hydrogen, and this is obviously dependent upon the ability of the acid to give them off into the solution or to dissociate.
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Thus, two acid solutions may be of the same concentration with respect to the total amounts of hydrogen available, yet have widely differing activity due to differences in the amount of active or ionized or dissociate hydrogen.
Hence we deal with a capacity effect, i.e., total available {dissociated plus undissociated) acid, as contrasted with an intensity or activity effect (dissociated acid or hydrogen ions alone).
As an example, let us compare acetic acid and hydrochloric acid. A liter of a normal solution of each contains exactly 1 gm of total available hydrogen, yet the activity of the N/1 acetic acid is slight while that of the N/1 hydrochloric acid is great.
Of the gram of available hydrogen in the acetic acid solution only 1.36 percent is in an ionized state, so that there is, in the liter of solution, only 0.0136 gm. of hydrogen ions. The gram of hydrogen in the liter of N/1 HC1 solution, therefore, is about 67 times as active or “strong” as the N/ 1 acetic acid.
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If one were to titrate the solutions, i.e., add N/1 NaOH solution until each became neutral, the total amount of alkali required would be the same in each case.
This is due to the fact that, as the alkali combines with the hydrogen ions, more hydrogen ions take their place from the undissociated acid, which strives to maintain a constant hydrogen ion concentration consistent with its dissociation constant.
Each acid finally give up all its available hydrogen and, since each of the solution by definition (N/1) contained exactly 1 gm. of available hydrogen to start with, each requires the same amount of alkali for its neutralization.
A measurement of hydrogen ion concentration differs from such a titration, in that the former determines the actual concentration of ionized hydrogen at the moment, without calling out any of the reserve, undissociated acid.
In acidimetric the term “normal” refers to the presence of 1 gm. of total available hydrogen per liter (dissociated plus undissociated). By contrast, a solution normal only with respect to ionized, hydrogen contains 1 gm. of hydrogen ions per liter.
This implies the presence of 1 gm. equivalent of a completely (100 percent) dissociated acid; an N/10 solution would contain 0.1 gm. equivalent of a completely dissociated acid, and so on.
If we were to express hydrogen ion concentrations or normality in terms of grams of hydrogen ions per liter we should have to deal with long words and long rows of zeros; a confusing and laborious system of nomenclature.
In 1909 Sorensen devised a simpler system based on the fact that water is itself a very weak electrolyte. As noted previously the extent of dissociation of any electrolyte is a physical constant (K) for that electrolyte under standard conditions.
A liter of pure, neutral water of 20°C always contains 0.0000001 gm. (1 X 107 moles) of OH–. In Sorensen’s system, the term “grams of hydrogen ions per liter” is replaced by the symbol pH, while the number of moles of H+ per liter (1 x 10-7 in the neutral water under discussion) is expressed as the logarithm of the reciprocal of the fraction, i.e., the positive number 7. The reaction of neutral water, and of any neutral solution, is therefore expressed as pH 7.
Now the product of the concentration of H+ and of OH– in neutral water is always 10-14 (H+ x OH– = Kw = 10-7 x 10-7 = 10-14). Since the product of the two is always the same (i.e., since the two are reciprocally related), the term pH is commonly used to express either.
For example, the pH of a solution containing 1 gm. of H+ per liter (i.e., normal [N/1] with respect to hydrogen ions or 1 gm. – equivalent of a completely dissociated acid) is 0 (log 1=0). Reciprocally, this is also pOH 14; the smallest fraction of a gram of OH+ per liter possible on the Sorensen scale. Similarly, pH 6 implies pOH 8; pH 2 implies pOH 12, and so on.
Since the number representing pH is derived from a fraction, the larger the fraction the smaller the pH number. Therefore pH numbers between 7 and 0 represent increasing degrees of acidity, and numbers between 7 and 14, increasing degrees of alkalinity. Unless one is familiar with the numbers they can at first be misleading.
For example, a change in pH from 7 to 6. represents a 10-fold increase in concentration of hydrogen ions since the 7 and 6 are logarithms; a change from pH 7.0 to 7.3 represents a 50 percent decrease in the concentration of hydrogen ions (1/2 x 10-7 = log 2 4- 7 = 0.3 + 7 = pH 7.3).