A practical guide to pH measurement in freshwaters

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A practical guide to pH measurement in freshwaters W. Davison Ambleside, U.K. It is particularly dQj%uit to measure accurately the pZZof solutions of low-ionic strength, such as freshwaters, due largely to errors associated with the liquid junction of the reference electrode. Careful selection of appropriate electrodes and implementation of simple, but rigorous quality control procedures can circumvent most of the problems and enable the acquisition of high quality data.

Introduction It is important to be able to measure the pH of freshwaters precisely and accurately. Most processes in natural aqueous systems are pH dependent and pH is usually an essential parameter for defining thermodynamic and kinetic constants. Concern about the effects of acid precipitation has also highlighted the need to compare measurements of pH made at different sites by different workers. Many measurements made in the past are now regarded as suspect, and there is a need to ensure that there is sufficient confidence in the accuracy of current measurements for them to be useful in assessing future trends. Recently there have been several studieslm6 which have highlighted errors associated with measuring pH in low ionic strength solutions such as freshwaters. They have clearly demonstrated that most of the problems are associated with poor performance of the liquid junction of the reference electrode. Although there has been speculation regarding the role of the ceramic frit!, the precise cause of the error is unknown. It has been shown that junctions with relatively high flow-rates2y3 and free diffusion junctions, which are simply formed by bringing the solutions into sharp contact with one another in a tube7T8 are usually free from such problems. There is general agreement 4s that there are few problems associated with the glass electrode, but a recent study’ has suggested that low concentrations of cations may adversely affect the response time in dilute neutral solutions

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The errors associated with the liquid junction referred to above are not to be confused with the residual liquid junction potential which unavoidably arises when the standard buffer solution is different from the solution being measured”. If this residual liquid junction potential was zero pH could be accurately interpreted in terms of the activity of hydrogen ions: that k pH would f3qUal paH, where pa, = -logi& and an is the activity of the hydrogen ion. Although the standard buffer solutions have been selected to minimize the difference between pH and pa, there can be a difference for dilute solutions. Consequently, the accurately measured value of the operationally defined pH will differ from pan, according to the value of the residual liquid junction potential. So, for example, the operationally defined pH of 10m4M acid will be about 4.03, considerably higher than the pa, which can be calculated to be 4.005 (ref. 2). Unfortunately there is currently no consensus value for the operational pH of such dilute acids, even though they are used as standards2. Three different approaches for producing high quality measurements in low ionic strength solutions have emerged. One group, recognising that the operational definition of pH is lax in not defining how junctions should be formed, has recommended that measurements should be made using free diffusion junctions8’i1’i2, or that such junctions should at least be used in a reference method to check the performance of more conventional electrodes397. Whereas free diffusion junctions avoid many of the errors associated with ceramic frits, well designed systems that can form very reproducible junctions are cumbersome when compared with commercially available combination electrodes2T7Y13.Another group uses commercial electrodes, but performs the final calibration with either dilute acids or synthetic samples of freshwater or rainwater14’15. This procedure reduces the errors and also helps to eliminate the residual liquid junction potential, making the measurement interpretable in terms of the hydrogen ion activity. However, it relies on an implicit assumption that any measurement bias which may be present will be the same for the dilute calibrating solution 0

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and the sample. Comparisons of measurements made on synthetic and natural waters suggest that this assumption may be invalid3. The third approach is based on the observation that some commercial electrodes have junctions which perform wel1214,and so rigorous testing and screening of the electrodes, to validate their suitability for making measurements in low ionic strength solutions, will ensure that measurements made in the conventional manner using standard buffers will give reliable and accurate results. Performance tests on various electrodes 2,4 have shown that, if electrodes give the same pH in stirred and unstirred solutions and the correct pH for a solution of dilute acid, they will generally perform well. Until recently there were few such electrodes that were commercially availabe, but new electrodes which use frits with higher flow-rates have been tested and found to perform we112. We have adopted this last procedure in our laboratory and allied it to a system of quality control. The recommendations which are given in this article have been well tested and found to provide a practical, pragmatic approach which allows the pH of freshwaters to be measured with maximum accuracy and minimum fuss. A practical guide If these recommendations are followed the measured pH should be a good estimate of the true value. No procedure can guarantee the correct answer, and other equipment and procedures may produce equally good results. Selection of an appropriate reference electrodes, and adoption of a welldefined analytical procedure with built-in quality control, are the two most important considerations. The procedures given are appropriate for measurements performed in the laboratory. As far as possible measurements in the field should be performed with similar care and the buffers should be maintained at the same temperature as the sample. Collection and storage of samples

The pH of poorly buffered freshwater samples can change with time because of the following: the ionic equilibria in solution are temperature dependent the quantities of dissolved gases present change by re-equilibration with the atmosphere, photosynthesis, respiration or microbiological degradation processes reactions occur with suspended solids which are not in equilibrium with the water. Making measurements on samples as quickly as possible (hours), by quickly returning them to a labora-

tory or a suitably equipped vehicle, will overcome these problems. The pH of acid waters (pH <5.6) is not appreciably affected by the production or loss of carbon dioxide, so well washed plastic bottles can be used for sample collection and transport. The pH of neutral samples (pH >5.6) is very dependent on the carbon dioxide content and so gaseous exchange must be minimised or prevented. Samples should be collected in well-washed bottles with a closure designed to eliminate air bubbles. The bottles should be prerinsed with the sample and then completely filled so that when the cap is replaced it displaces water, preventing the entrainment of air. Bottles can be made of high density plastic or borosilicate glass, but metal, soda glass or low density, thin-walled, plastic containers should be avoided. Darkened bottles prevent photosynthesis which can raise the pH, but still permit respiration which can depress the pH. The pH of CO,-sensitive neutral samples should be measured by inserting the electrode into the bottle, so minimizing contact with the air. Ideally the container should be maintained at the in situ temperature of the natural water and the pH measured at the same temperature. If the measurement is made at some higher temperature it is possible, for neutral samples collected in sealed bottles, to use the temperature dependence of the carbonate stability constants to calculate the in sihc pH. The pH decreases by about 0.09 for each 10 “C increase in temperature. It is not simple to apply a temperature correction for acid samples, but it can be donei6. Acid samples can often be stored for several days or weeks without any appreciable change in pH. Electrodes

Electrodes must be appropriate for use in dilute solutions. Reference components which have fast flowing junctions (>20 ~1 h-l, most often a frit) usually work best, and those using a calomel internal element are less prone to clogging than those using silver-silver chloride. Russell CTL/LCW (Russell pH) and Amagruss CWL/LCW (Flowtech Instruments) combination electrodes have been found to perform well. There is no special merit in using separate glass and reference electrodes. Most manufacturers’ research grade glass electrodes are adequate, but a Corning 00311101 and a Radiometer G202C have been found to work well. Beckman 39416 calomel electrodes with quartz fibre junctions have performed very well as separate reference electrodes, but, because of their very high leakage rates (ca. 100 yl h-l), contamination by KC1 may be a problem for prolonged use in small volumes of sample. The filling solution of the reference electrode should be well

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maintained and the frit should never be allowed to dry out. Gel-filled and permanently sealed electrodes should be avoided. Meter

Most modern meters with a facility to read in mV to ~1 mV should be satisfactory. Beware of dampness affecting battery operated instruments, and inducing pH shifts and/or drift. Calibration and storage

Buffers, based on phthalate (pH 4) and phosphates (pH 6.8-7.6), purchased from a reputable manufacturer should be used. If they require preparation from a powder or tablet, care should be taken to avoid contamination, especially from acid-rinsed glassware. Buffers and unknowns should be at the same temperature ( f 2 “C) and the pH of the buffer, appropriate to that temperature, should be usedi’. Calibration should be performed daily with additional checks at the end of each series of measurements (hourly); with good electrodes there should be no change, to within 0.02 pH. Separate reference electrodes can be stored in 3.5 M KCl, and glass electrodes in pH 4 buffer. Combination electrodes should be stored in 3.5 M KC1 to which some pH 4 buffer is added (cu. 5%). Quality control

A well-defined routine for performing the measurements, using the same equipment each time, should be adopted. The electrode performance tests (l)-(4) should be applied to newly purchased electrodes, to electrodes which have not been used for more than a week, and to regularly used electrodes at monthly intervals. (1) Electrode sensitivity. Using the standard calibration buffers and the pH meter in mV mode the Nernstian response of the electrode can be simply checked: at 20 “C, 58 + 1 mV per pH increment, and at 25 “C, 59 + 1 mV per pH increment17. (2) Accuracy in dilute solutions. A solution of 10e4 M hydrochloric acid or 5.10e5 M sulphuric acid can be prepared accurately by weight, from a commercial volumetric stock. The pH should be between 4.00 and 4.05, irrespective of whether the solution is stirred or quiescent. If the acid is suspect an alternative check can be made using dilute buffers. pH 4 buffer is a solution of 0.05 mol kg-’ potassium hydrogen phthalate. When this is diluted 5-fold to 0.01 mol kg-’ it should have a pH of 4.11 + 0.01 at 20 “C and 4.12 +_0.01 at 25 “C (ref. 18). (3) Stirring shift. The shift in pH obtained upon stirring a formerly quiescent solution of dilute acid

should be SO.02 pH. (4) Response time. The reading obtained using dilute acid should be stable to within 0.01 pH within 2 min of immersion of the electrodes, irrespective of the solution used prior to washing. Natural waters with pH <5 will probably produce a stable reading in a similar time, but neutral dilute solutions (pH 6-8) may take up to 10 min to fully stabilize due to inherent problems with the glass electrode. When working with such solutions it is advisable to make some preliminary checks by measuring the pH at 2,5 and 10 min. If one of the shorter equilibration times gives the same result as the measurement made after 10 min, it may be used. Failure of tests (2) and (3) are usually associated with the reference electrodes. (1) and (4) could also be due to the glass electrode. Some attempt can be made to rejuvenate the defective component, after which it should be re-tested, but if this fails the electrode should be discarded. Rejuvenation may be attempted by (a) refilling the reference electrode with fresh KC1 filling solution and then leaving the junction to soak in concentrated KC1 for a day, or (b) soaking the glass electrode 5 times alternately in 0.1 M HCl and 0.1 M NaOH solutions for 5 min. Electrodes should be rinsed thoroughly before re-testing. If a conductivity meter is available the leakage rate of the junction can be measured by dipping the electrode into a beaker of distilled water while monitoring the conductivity. Calibration can be achieved by pipetting known aliquots of the filling solution into distilled water. The desired leakage will depend on the electrode, but should generally be >lO ~1 h-l (refs. 2 and 3).

Daily checks

To ensure fast response and well-mixed samples, measurements should usually be made whilst stirring the solution. A simple check of performance can be made by occasionally switching off the stirrer and checking that the pH of a low ionic strength sample does not change by more than 0.02 pH. As there should be virtually no difference between the stirred and unstirred values measurements on quiescent solutions are acceptable. For electrodes with an appreciable stirring shift, indicating a poorly performing junction2’18, the unstirred value is usually closer to the correct answe& lg . Electrodes should be washed between measurements with distilled or deionized water. The distilled (not deionized) water should have a pH of 5.64 + 0.1 if it is air-equilibrated, and, as a given laboratory source usually has a constant pH (ref. 2)) the rinse value can be used as a crude

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form of quality control. More rigorous quality control can be attained by including a sample of dilute acid in each batch. General Better pH measurements undoubtedly result if there is some understanding of how electrodes work (see Westcott2’ for a simple account and Bates17 for a more comprehensive treatment). Some manufacturers’ guides to pH measurement are excellent and recommended procedures for using individual electrodes should be followed. Electrodes should not be expected to last forever; normally they will give good service for one or two years, depending on use. Those that are used regularly (daily) normally perform and last better than those used intermittently. It is advisable to have two electrodes in continuous use to allow cross checks and as a guard against sudden failure or damage. References 1 2 3 4 5

R. C. Metcalf, Analyst, 112 (1987) 1573. W. Davison and T. R. Harbinson, Analyst, 113 (1988) 709. W. Davison and C. Woof, Anal. Chem., 57 (1985) 2567. D. Midgley, Atmos. Environ., 21(1987) 173. G. Marinenko, R. C. Paule, W. F. Koch and M. Knoerdal, J. Res. Natl. Bur. Stand., 91 (1986) 17. 6 D. P. Brezinski, Analyst, 108 (1983) 425.

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7 T. R. Harbinson and W. Davison, Anal. Chem., 59 (1987) 2450. 8 A. K. Covington, P. D. Whalley and W. Davison, Anal. Chim. Actu, 169 (1985) 221. 9 W. Davison and T. R. Harbinson, Analyst, 113 (1988) 1537. 10 A. K. Covington, R. G. Bates and R. A. Durst, Pure Applied Chem., 57 (1985) 33. 11 A. K. Covington, P. D. Whalley and W. Davison, Pure Applied Chem., 57 (1985) 877. 12 A. K. Covington, P. D. Whalley, W. Davison and M. Whitfield. In T. S. West and H. W. Nurnberg (Editors), The Determination of Trace Metals in Natural Waters, Blackwell, Oxford, 1988, p. 161. 13 C. Jones, D. W. Williams and F. Marsicano, Sci. Total Environ., 64 (1987) 211. 14 W. F. Koch, G. Marinenko and R. C. Paule, J. Res. Natl. Bur. Stand., 91(1986) 23. 15 W. F. Koch, G. Marinenko and R. C. Paule, J. Res. Natl. Bur. Stand., 91(1986) 33. 16 W. Davison, C. Woof and E. Tipping, Analyst, 114 (1989) 587. 17 R. G. Bates, Determination ofpH, Wiley, New York, 1964. 18 A. K. Covington, P. D. Whalley and W. Davison, Analyst, 108 (1983) 1528. 19 N. R. McQuaker, P. D. Kluckner and D. K. Sandberg, Environ. Sci. Technol., 17 (1983) 431. 20 C. C. Westcott, pH Measurement, Academic Press, New York, 1978. Dr. Davison is at the Institute of Freshwater Ecology, Windermere Laboratory, The Ferry House, Far Sawrey, Ambleside, Cumbria LA22 OLP, U.K.

observer

Industrial applications of vibrational spectroscopy Kathryn S. Kalasinsky Mississippi

State, MS, U.S.A.

Vibrational spectroscopy has played an ever increasing role in the analytical laboratory. As instrumentation and sampling accessories evolve, the applications to industrial an&ses broaden. Reported here is a brief discussion of the field of vibrational spectroscopy and its applications as reported in the literature as well as several specify sample analyses.

Analytical applications of vibrational spectroscopy have made their biggest impression on society through art work validation’ and forensic science2- . Although these represent only a fraction of the applications, they demonstrate the rapidly broadening 0165-9936/90/$03.00.

field of analyses suited for vibrational spectroscopy. Microscopy has played an important role in expanding the field of vibrational techniques into microspectroscop$-8. The ability to analyze a single fiber or minute paint chip, combined with the specificity and non-destructive nature of vibrational techniques, has made applications such as those mentioned above possible. Industrial concerns that need to closely monitor roducts such as plasticizers’ and P have found vibrational specpharmaceuticals lop1 troscopy to provide specific data rapidly. Biomedical research has also found the utility of vibrational spectroscopy12-14; the biggest problems of this field are not experimental design but interpretation of all the available data. @Elsevier

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