Why pH is Important
pH is one of the most common laboratory measurements because many chemical processes are dependent on pH. The speed or rate of chemical reactions can often be significantly altered by changing the pH of the solution. The solubility of many chemicals in solution, and their bio-availability is dependent on pH. The physiological chemistry of living organisms usually has very specific pH boundaries. Virtually everything we use has been tested for pH at one time, tap water, food, medicines, at some point a pH measurement was performed.
Basic pH Theory
The term pH derives from a combination of “p” for the word power and ”H” for the symbol of the element Hydrogen. Together the meaning is the power or exponent of hydrogen.
The Chemical Equation for pH
pH is defined as the negative log of hydrogen ion activity, where the activity, aH+, describes the free hydrogen ion, or the “effective concentration”, in the presence of other ions:
pH = -log aH+ or H+ = 10-pH
Thus, a pH of 3 is equivalent to a hydrogen ion activity of 10-3 Molar(M), a pH of 11 is an activity of 10-11 M, and a pH of 11.5 would be a hydrogen ion activity of 10-11.5 or 3.2 x 10-12 M. Water (H2O) dissociates into hydrogen ions (H+) and hydroxide ions (OH-) in an aqueous solution. The following equilibrium reaction is used to describe pH:
2H2O = H30+ + OH- or simply H,O = H+ + OH-
The dissociation constant, Kw, is the product of the hydrogen and hydroxide ion concentrations:
Kw = [H+] [OH-] =1.0 x 10-14 M at 25 °C
At 25 °C, Kw remains constant at 1 x 10-14 M. Therefore, the concentration of hydrogen or hydroxide ions can be calculated if the other concentration is known. A pH of 7 is considered to be neutral at 25 °C because the activity of the hydrogen and hydroxide ions are both equal to 10-7 M. The activity range for the hydrogen ion, as defined by the dissociation product, is 100 to 10-14 M. The activity range of hydrogen ion relates to a pH scale of 0 to 14. Each unit on the pH scale represents a ten-fold change in the activity. For more information, see notes (1) and (2).
Practical Application of pH
pH serves as a convenient way to compare the relative acidity or alkalinity of a solution at a given temperature. As discussed, a pH of 7 describes a neutral solution because the activities of hydrogen and hydroxide ions are equal. When the pH is below 7, the solution is described as acidic because the activity of hydrogen ion is greater than that of hydroxide ion. A solution is more acidic as the hydrogen ion activity increases, therefore the pH decreases. Conversely, as the hydroxide ion activity increases, the solution becomes more alkaline, also referred to as basic, and the pH will increase. In practice, pH electrode measurements are made by comparing readings in a sample with readings in standards whose pH has been defined “buffers”. These measurements are relative rather than exact thermodynamic determinations of activity. pH electrode measurements can be used to detect a titration endpoint, which will give the acidity or alkalinity in terms of total concentration, rather than activity.
Basic pH Electrode Theory
pH electrodes measure the pH of a solution potentiometrically. A potentiometric measurement relies on an electrical signal. When a pH sensing electrode comes in contact with a sample, a potential develops across the sensing membrane surface. The membrane potential varies with the pH. Making a measurement requires a second unvarying potential to quantitatively compare the changes of the sensing membrane potential. A reference electrode provides this function. Electrode behavior is described by the Nernst equation:
Emeasured = E0 + (2.3 RT/nF) log aH+
Emeasured is the measured potential from the sensing electrode, E0 is related to the potential of the reference electrode, (2.3 RT/nF) is the Nernst factor, and log aH+, is the pH. The Nernst factor, 2.3 RT/nF, includes the Gas Law constant (R), Faraday’s constant (F), the temperature in degrees Kelvin (T) and the charge of the ion (n). For pH, where n = 1, the Nernst factor is 2.3 RT/F. Since R and F are constants, the factor and therefore electrode behavior is dependent on temperature. The electrode slope is a measure of the electrode response to the ion being detected and is equivalent to the Nernst factor. When the temperature is equal to 25 °C, the Nernst factor or slope is 59.16 mV/pH unit. Most Orion pH meters display the slope as a percentage of the theoretical value. For example, a 98.5% slope is equivalent to a slope of 58.27 mV/pH unit for a two-point calibration. When a pH meter detects the sensing membrane signal, reference signal and the temperature, the meter software calculates the pH using the Nernst equation. Orion microprocessor controlled pH meters contain pH versus temperature values for commonly used buffers. This allows the meter to recognize a particular pH buffer and calibrate with the correct value.
Notes
(1) “Uses and Abuse of pH Measurements”, Feldman, Issac, Analytical Chemistry, 28, 1861 11956)
(2) “Determination of pH: pH Theory and Practice”, Bates,Roger, John Wiley and Sons, New York, 19733
1/1/2025