Why is pH Important?
pH is one of the most common laboratory measurements because many chemical processes are dependant 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-availablity is dependent on pH. The physiological chemistry of living organisms usually has very specific pH boundaries. In our modern lives, virtually everything we use has been tested for pH at one time: from the tap water we brush our teeth with, the paper we write on, the food we eat, to the medicines we take, at some point a pH measurement was performed.
Basic pH Theory
The Term pH
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 a H+ 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 1 0 -11.5 .
Water (H20) dissociates into hydrogen ions (H+ ) and hydroxide ions (OH-) in an aqueous solution. The following equilibrium reaction is used to describe pH: 2H 2 O = H 3 O + + OH - or simply H 2 O = H + + OH - The dissociation constant, Kw, is the product of the hydrogen and hydroxide ion concentrations: Kw = [H+] [OH] = 1.0 x 10-16 M at 25 oC At 25 oC, Kw remains constant at 1 X l 0 -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 o 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 10 - 0 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 references 1 and 2.
1. 'Uses and Abuse of pH Measurements,' Feldman, lsaac, Analytical Chemistry, 28,1861 (1956).
2. "Determination of pH: Theory and Practice,' Bates, Roger, John Wiley and Sons, New York, 1973.
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, aiso 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:
E measured = EO + (2.3 RT/NF) log a H+
E rneasured is the measured potential from the sensing electrode, EO is related to the potential of the reference electrode, (2.3 RT/NF) is the Nernst factor, and log a H+ 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/E. 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 oC, the Nernst factor or slope is 59.16 mV/pH unit. All Thermo Electron 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. Thermo Electron 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.
pH and Temperature
The most common cause of error in pH measurements is temperature. There are at least five ways that temperature variations can affect pH:
Electrode Slope
pH Buffers
Samples
Reference Element Drift
Temperature Sensor Errors Electrode Slope Changes
The electrode slope will change with variations in temperature. Slope changes may be compensated manuallyor automatically with a temperature compensation probe (ATC) . Thermo Electron meters calculate the slope value based on the temperature input and automatically adjust the measured pH values. Figure 1 illustrates the change in electrode slope with temperature.
Figure 1 Change of pH Electrode Slope with Temperature
Buffer and Sample pH Changes
Buffer and sample pH values vary with temperature because of their temperature dependent chemical equilibria. The problem of differing pH values is easily solved by calibrating the electrode with characterized standard buffers whose true pH values versus temperature are known. pH buffer values at different temperatures are given in Table 1. Thermo Electron meters calibrate with the correct pH buffer values based on the manual, ATC or digital logr temperature value. The problem of the sample equilibrium varying with temperature in an uncharacterizable manner will always remain. Therefore, calibration and measurement should be performed at the same temperature and pH values should be reported along with temperature. For best results an ATC probe should be used.
Reference Element Drift
Drift can occur when the reference junction potential changes with time due to contamination of the junction material with sample. K-series electrodes use a patented internal reference gel that minimizes this effect.
Temperature Sensor Errors
When a pH and temperature probe are placed into a sample that varies significantly in temperature, the readings can drift for two reasons. First, the temperature response of the electrode and temperature probe may not be similar which prolongs equilibration and drift. Second, a sample may not have a uniform temperature, therefore the pH electrode and temperature probe are responding to different environments |
