Field Measurement of pH
Testing for pH is a fundamental part of any water treatment program. It is important that the testing procedure is suitable for the situation in which it is used and that it is properly performed. If neither of these conditions is met, the results can be inaccurate.Field Methods for Measuring pH
The determination of pH may be accomplished by two general methods: potentiometric or colorimetric. Selection of the most suitable method depends on several factors:
Composition of the aqueous solution to be tested.
Is it slightly or highly buffered?
Will dissolved solids concentrations vary over a wide range?
Are strong oxidizing agents and/or dissolved gases present?
Will the process yield a colloidal suspension, turbidity, or course particulate material?
Will the solution be colored?
Temperature range of the process.
pH range of the process:
How rapidly may extreme changes of pH occur?
What variation in pH may be tolerated before adverse results are encountered in the process?
Each method of pH measurement has its advantages and its limiting conditions. In general, the potentiometric method is the more reliable of the two methods. Objections to its use in the field are its cost and lack of portability.
The colorimetric method uses pH indicators that can be adversely affected by many substances. As a general rule, indicators are quite unreliable for measuring pH values below 3 and above 10. Determinations of pH in clear, colorless, well-buffered samples can be made quite accurately; but great care must be exercised to add the proper amount of indicator, and a reliable set of standards must be used. The colorimetric method can be tedious unless the approximate pH is known in advance.
Both methods require extraordinary precautions when measuring the pH of a slightly buffered solution or ultra-pure water. Because of its wide use in the field, the colorimetric procedure will be discussed first.
Colorimetric Determination of pH A practical colorimetric indicator must meet certain specifications. A definite, but gradual, color change over a narrow pH range must be obtained; substances other than hydrogen or hydroxyl ions should not affect the color changes, and the color formation must occur rapidly and produce a stable color. In general, these indicators are weak acids or bases that change color when changed from the neutral to the ionized form. These colors can be maintained with buffer solutions, the hydrogen ion concentration of which remains constant. Since the colors are characteristic of definite hydrogen ion concentrations, they can be used to estimate the hydrogen ion concentration by a system of comparison with a set of standards.
The colorimetric pH determination in buffered solutions is then simply a matter of knowing the approximate pH range needed and choosing an indicator that responds in this range. A series of buffer solutions with known pH values, containing equal amounts of indicator are prepared to produce a series of stable colored standards for visual comparison. These standards can be enclosed in glass ampoules or tubes, such as LaMotte or Taylor comparators; or duplicated on glass or stained gelatin, such as the Hellige disc comparators. A continuous spectrum may be applied to plastic, such as the Hach compactors. In every case, a blank tube containing sample only, must be used to partially compensate for sample color and hue. The amount of sample and indicator must not be altered. The procedure must be followed in every detail or unreliable data will result.
In no case should a color match with the first or last standard in any color standard set to be taken as an accurate determination. In such cases, the determination should be repeated with an indicator overlapping the range of the first indicator.
The sources of errors in the colorimetric determination of pH are many, especially when a slightly buffered or ultra-pure system is being measured. The errors encountered when dealing with buffered solutions are, for the most part, easily overcome. Avoiding them will produce accurate data. The errors and corrective measures are:
Maintaining clean glassware. Rinse after each use. If discoloration occurs, clean with a strong acid solution or strong alkaline solution and rinse thoroughly with distilled water. If scratches or fogging persists, use new tubes.
Do not allow standard tubes or discs to be exposed to prolonged light or high temperatures; this will cause deterioration with fading.
It is apparent that the color of a solution containing an indicator will be more intense as the dye concentration is increased. If you find that you must add more or less indicator to obtain the same result, something is wrong. Verify data with a pH meter. Check for deterioration of pH indicator solution and/or presence of adverse substances in the unknown solution. Regarding such substances, check the following questions:
Has a “redox” substance appeared? Generally, bleaching agents cause fading of the indicator, but halogenation can also intensify the color of certain indicators.
Has a dramatic change in ionic strength occurred? Has “salt” concentration dropped or increased? The ionization of weak acids is slightly, but noticeably, affected by the presence of neutral salts, and the optical properties of indicators are also altered by salts.
Is some organic compound affecting pH? Generally organic substances alter the indictor’s affinity by making the results unpredictable.
Finally, there may be no problems, but indicator solutions are behaving erratically. Again, verify with a pH meter. When in doubt, replace indicator solutions.Indicator Papers
Indicator test papers are available for pH determination. The paper is saturated with the appropriate pH indicator (short range) or a mixed indicator (wide range). By placing a drop of the test solution on the paper, the color is developed and compared with a chart. The use of pH paper is a very rough method for determining pH, and many problems are encountered. These papers are usually used only for pH adjustment. The following guidelines are helpful when using pH paper:
pH papers deteriorate rapidly. Avoid exposure to fumes and wrap dispenser and paper in foil when not in use.
Indicators are usually leached or washed out easily. Never leave indicator papers in the unknown solution when making a pH adjustment. Use a fresh strip of indicator paper and dip it momentarily into the unknown solution during pH adjustment, or add a drop of the unknown to a fresh piece of indicator paper. Never trust the comparison of the extremes on any color chart. A short-range paper is useless unless the unknown solution is within 0.5-pH unit of the paper. Indicator papers are available that allow for several color comparisons over a narrow pH range; that is, a match must occur over two or more colors to determine pH. These papers are usually more susceptible to running or leaching of the indicator. This problem has been avoided by use of a plastic coating on the paper. The added cost, coupled with the inherent inaccuracies of pH paper determination, does not warrant their use except under controlled conditions.
The leaching or running quality of indicator papers is used to advantage in pH measurement of many systems. Once a system is in control and limits are established, these papers offer a rough method of checking pH. The sample is placed in a standard sample tube. A definite length of pH paper is added, and the tube is stoppered and mixed. The indicator is leached from the paper, and the colored solution is now compared against a color chart. This technique can be highly successful if all precautions are observed and the procedure checked as frequently as needed with appropriate pH comparators or a pH meter.Potentiometric Determination of pH
The second and most accurate method that can be used for the determination of pH in a solution depends upon the measurement of the potential produced at an indictor electrode when it is dipped into the sample. The measurements relative to a reference electrode are generally carried out with some form of potentiometer, so as not to draw any appreciable amount of current from the reference cell, nor produce any polarization effect at the surface of the indicator electrode. Several types of electrodes have been suggested for the electrometric determination of pH. Although the hydrogen gas electrode is recognized as the primary standard, the glass electrode, in combination with the reference potential provided by a saturated calomel electrode, is most generally used. Developments in electrode technology and electronic circuitry have advanced the pH system from the hydrogen gas electrode with a potentiometer, to completely seal, combination electrodes with LED display. This system of determining pH is no longer confined to the laboratory. The small portable battery operated units are used extensively in the field. The glass electrodes capable of detecting such quantities as parts per billion and less are now called “probes”. The pH amplifier, which consistently amplifies currents lower than one millionth of an ampere, is now called a “pH meter”. The instrumentation is simple and trouble free, so much so, that maintenance is frequently forgotten. We forget basic pH theory and expect one probe to perform without error regardless of temperature, ionic strength, or sample peculiarities.
It is not our intent to dwell on the theory of pH electrodes or meters. Many questions can be answered by consulting the manual for the meter or probe. However, several basic rules should be followed on pH meter use and maintenance:
Always follow the manufacturer’s sequence of instructions for use of the meter.
Confirm all temperature readings with a thermometer.
After removing the electrode from buffer storage, be sure to rinse the electrode thoroughly with distilled or deionized water.
Blot the electrode bulb gently with paper tissue. Avoid rubbing, as this generates static electricity and can contribute to resistance.
Pure water carries high resistance, making pH measurements very sensitive to electrical interference (“noise”). Whenever possible, keep your meter away from motors and pumps and avoid changes of electrical field and static effects.
Standardize the instrument against a standard buffer solution with a pH approaching that of the sample, and then check the linearity of the electrode response against at least one additional standard buffer or a different pH. The readings with the additional buffers will afford a rough idea of the limits of accuracy to be expected of the instrument and the technique of operation.
Sluggish, erratic results arise from mechanical or electrical failure, such as:
Weak batteries – replace and check all leads.
Cracked electrodes – replace.
Plugged or fouled electrode – clean according to manufacturer’s instructions.
Organic material is usually removed by soaking in a slightly acidic alcohol solution overnight.
Inorganic precipitates are usually removed with quick dips in dilute solutions of acids or bases.
After any cleaning procedure, always rinse with distilled water and treat the electrode as if it were new. Most electrodes must be climatized before use by soaking the tip in a buffer solution.
Is level adequate? Check manufacturer’s specification.
Has electrode well become contaminated with sample? If so, clean thoroughly with distilled water, rinse with filling solution, and refill with fresh filling solution.
Are you using the correct filling solutions? Two combination pH electrodes are currently available. The calomel electrode requires saturated potassium chloride for the filling solution, while the silver/silver chloride electrode requires a mixture of saturated potassium chloride and silver chloride. Check the reference metal to distinguish between the two electrodes. The calomel electrode has a cylinder-shaped, two-phase, mercurous chloride coated mercury reference element. The silver/silver chloride electrode has a thin, silver chloride coated wire or a flat, paddle-shaped piece of silver metal coated with silver chloride as its reference cell. All pH electrodes have a small opening to permit a minute amount of filling solution (0.01 to 0.1 milliliters per day) to mix into the sample. The filling solution must flow or diffuse freely into the sample to maintain correct electrical resistance and assure accurate pH readings.
Combination electrodes are unique in that they contain both the reference and sensor in the same glass envelope, thus eliminating the cumbersome two-electrode system. The silver/silver chloride electrode has become very popular because of its high stability during temperature fluctuations. However, its use may have a definite drawback in water analysis: namely, pH measurements in ultra-pure, low ionic strength water (condensates). Silver chloride is much more soluble in potassium chloride than it is in pure water. When a silver/silver chloride filling solution is inserted in a low ionic strength water solution, and diffusion of the filling solutions begins, silver chloride may participate and clog the electrode opening. A very high electrical resistance results in sluggish meter response and pH measurement errors. In short, if slow meter response is your problem, you might invest in a calomel combination electrode and reserve it for this special use onlyElectrode Readiness:
If you use your pH meter frequently, we suggest you construct a permanent setup using a plastic sleeve, etc., and keep your combination electrode always immersed in buffer and ready to use. When an electrode dries out:
The diffusion hole is plugged and the electrode is inoperable. It has to be cleaned and reconditioned.
Permanent damage may occur, necessitating electrode replacement. Maintain your equipment and it will reward you with reliable data.
In summary, pH is an important and highly useful parameter. Establish the degree of sophistication needed on any particular job and decide which method to use. Once a system is in control, a comparator (colorimetric) method is a convenient tool for maintenance; but be prepared to back it up with pH meter analysis.