Why is pH measurement important?

pH measurement is one of the most important and common tests in water chemistry. Almost every step in water supply and wastewater treatment (e.g., acid-base neutralization, water softening, sedimentation, coagulation, disinfection, and corrosion control) depends on pH. pH is used in the measurement of alkalinity and carbon dioxide and many other acid-base equilibria. At a given temperature, the strength of an acid or base activity of a solution is indicated by its pH, or hydrogen ion activity. Alkalinity and acidity are the capacity to neutralize the acid or base present in water and are usually expressed in milligrams of calcium carbonate per liter. Buffer capacity is the amount of strong acid or base (usually expressed in moles per liter) that is required to change the pH of a 1-liter sample by 1 unit. pH, as defined by Sorenson1, is -log [H+], which is a measure of the "strength" of acidity. Pure water is very slightly ionized and is in equilibrium with ionic products.

((1))

and

Where:

[H+] = activity of hydrogen ions, mol/L

[OH-] = activity of hydroxyl ions, mol/L and

Kw = ionic product of water.

Because of the ionic interactions in all very dilute solutions, it is necessary to use the “activity” of the ion rather than its molar concentration. In using the term pH, it is assumed that the activity of the hydrogen ion, aH+, is considered. Because of the approximate equivalence with molarity, [H+] is used instead of activity only in very dilute solutions (ionic strength < 0.1).

The logarithmic scale is suitable for expressing a wide range of ionic activities. Equation 1 is expressed in logarithmic form and modified to reflect activity:

((2))

or

Where:

pH = log10 aH +

and

pOH = log10 aOH-.

What is the pH at different temperatures?

Equation 2 states that as pH increases, pOH decreases inversely, and vice versa, because pKw is constant at a given temperature. At 25 °C, pH 7.0 is neutral, the activities of hydrogen and hydroxyl ions are equal, each corresponding to an activity of approximately 10-7 mol/L. The neutral point is temperature-dependent; pH 7.5 is neutral at 0 °C and pH 6.5 at 60 °C.

pH of a dilute solution

The pH of a very dilute solution is approximately equal to the negative logarithm of the hydrogen ion concentration. Natural waters usually have a pH in the range of 4 to 9, and most of them are alkaline due to the presence of bicarbonates and carbonates of alkali and alkaline earth metals.