2496
Anal. Chern. 1986, 58,2496-2502
Measurement of pH in Solutions of Low Buffering Capacity and Low Ionic Strength by Optosensing Flow Injection Analysis B r u c e A. Woods, J a r o m i r Ruzicka, a n d Gary D. Christian* Center for Process Analytical Chemistry, Department of Chemistry, BG-IO, University of Washington, Seattle, Washington 98195 Robert J. Charlson Department of Civil Engineering, University of Washington, Seattle, Washington 98195
OptosenSing Wow injection analysls is investigated in this work as a technique for determining the pH of low ionic strength solutions. The optosensing technique involves reflectance spectroscopy of an lmmoMUred chemically selective reagent that results in the formetion of a colored indicator species. The MkaUng reactlon Is reversible. Optimal wavelengths for monitoring the color change and thkkness of the optosensing surface are determined. The durability of the immobilized indicator and the effect of different carrier buffer concentrations are investigated. The pK,’s and responses for three commercially available ImmoMiized indicators covering three different pH ranges are investigated. With inJectedsample volumes of 200-275 KL, approximately 30-35 determinatlons to per hour were performed, using solutions in the M hydrogen ion activity range to model the system. The relative standard deviation ranges from 0.02 to 0.17 pH units for the three Indicators investigated.
The determination of the pH of a solution is routinely done in nearly all laboratories. There are many possible ways for determining pH, such as the use of glass electrodes ( I ) , pH papers ( 2 ) , and indicators ( 3 ) . The glass electrode is probably the most widely used tool for the measurement of pH, because it offers good precision and measurements can be done rather rapidly. The glass electrode can be incorrectly used in solutions of low ionic strength, including rainwater, natural waters such as lakes or streams, boiler feedwater, and so forth. A recent paper by Covington et al. ( 4 ) reports on some of the problems encountered when determining the pH of low ionic strength solutions such as boiler feedwater and natural waters. Two recent papers, by McQuaker et al. (5) and by Sisterson et al. (6),discuss some of the possible errors associated with pH measurements of rainwater. The major errors for the glass electrode include (a) “streaming” potentials due to the liquid flowing past the electrode, (b) liquid junction potentials due to the differences in the ionic strength of the calibration buffers and samples, and (c) contamination of the poorly buffered samples by carryover of sample on the electrode. The total error associated with these can lead to reported pH values being incorrect by 0.1-0.15 pH units. Koch and Marinenko had proposed a protocol (7) for measurement of rainwater pH that recommends reporting pH values no more closely than 0.1 pH unit. The National Bureau of Standards has recently issued a Standard Reference for Rainwater (SRM 2694) with specific guidelines as to the electrometric measurement of pH of low ionic strength solutions. Galloway et al. (8) have proposed the use of dilute strong acids for calibration and the measurement of quiescent solutions to minimize liquid junction and “streaming” potential errors for pH measure-
ments of rainwater. In contrast, Covington et al. ( 4 ) recommend the use of a modified flow cell for determining the pH of low ionic strength samples. Another convenient and rapid method for determining pH is through the use of pH papers. While this method is less accurate than a glass electrode (0.2-0.3 pH unit accuracy for the indicator when using visual comparison with the color scale), the accuracy can be improved significantly by using a spectrophotometer to determine the color intensity change of the pH indicator pad. In early work done by Landsberg in 1952 (9), a large piece of indicator paper was used in place of a glass electrode to determine the pH of individual raindrops since they felt an accuracy of only 0.5 pH unit was needed. The problem in the use of early pH papers, though, was the bleeding of indicator dye from the pH paper into the solution being analyzed. This problem has been overcome with the introduction of nonbleeding p I j papers (2)in the last few years. These nonbleeding pH papers have the indicator dye covalently bound to a cellulose matrix. Indicators can also be used to determine the pH of solutions by preparing a series of buffered standards that cover the region of interest. An indicator is added to the buffered standards and the samples and a visual comparison is made to determine the pH of the test solution. This is not applicable to turbid or colored solutions and is a tedious process. Because of the uncertainties associated with low ionic strength determinations, we undertook to investigate a new approach of optosensing in flow injection analysis (FIA) for monitoring changes in pH indicator papers upon exposure to an injected sample plug. This method is based on reflectance spectroscopy of a chemically reactive surface (Figure 1) to determine a selected chemical species (IO). For our system, we chose to perform reflectance spectroscopy on a nonbleeding pH indicator pad. Advantages of using FIA include the short time interval between samples, allowing better determination of the rate of change of various chemical species during a short time interval, as in a rainstorm. The sample is protected from contamination by the environment because it is in a sealed flow path. The analytical system can be made very small and portable for field use. This allows the analytical system to be taken to the field site, which provides several advantages. Such an FIA system would offer the following: no contamination or change in the sample, as when collecting and transporting to the lab; the system is not as fragile as a glass electrode; the sensing surface is inexpensive; and the sytem can be set up to do essentially real time measurements like an on-line process analyzer. The sample necessary for a determination would be a few hundred microliters or less, compared to a few milliliters necessary for a semimicro combination glass electrode. The capability to do almost real time monitoring would also allow compliance with the latest U.S. EPA protocol ( I I ) that does not allow a “holding time” between when a sample is
0003-2700/86/0356-2496$01.50/00 1986 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 58, NO. 12. OCTOBER 1986 * 2497 LIGHT SOURCE
Figure 1. A general optosensing cell with a fibrous flow-through structure m which a chemically selective reagent is covalently bound: (A) bifurcated fiber optic for the light source and detector; (6) fibrous indicator pad; (C) reflector. taken and when the measurement of the pH is performed. This no holding time is usually considered to he less than 15 min. Covington e t al. ( 4 ) have also made a recommendation that pH determinations should he made shortly after sampling, because deterioration of the sample by bacteria present can cause pH changes in as little as 1-2 h. EXPERIMENTAL S E C T I O N Reagents and Solutions. AU the solutions were prepared from analytical grade reagents and deionized water ( R = 7 MR/cm). Dilute HCI standards were prepared from commercially available 0.1000 N HCI WWR Scientific). The dilute HCI standard concentrations were verified by titrating against a standard Tris solution (Gold Label, Aldrich). Sodium hydroxide solutions were prepared by dilution of a commercially available 1.0 M sodium hydroxide solution. The concentration of a 0.1 M sodium hydroxide stock solution was checked by titration against the 0.1oOO N HCI solution hy the method of Gran (12). All buffers prepared in this work were made following directions in the monograph by Perrin and Dempsey (13) or by modifying the directions in Pemn and Dempwy to prepare dilute buffers. Dilute buffers were prepared by diluting the acid salt in 900 mL of deionized water and then adjusting the pH by addition of acid; the buffer was then diluted to 1 L and the actual pH determined on a small aliquot. The final pH values of all buffers were determined with a pH meter using a glass electrode. The pH meter was calibrated with pH 4,7, or 10 buffers traceable to NBS standard buffers and using a two-point calibration. All pH papers used in this work were ColorpHast indicator papers (Merck). Equipment Used. A FiAtron SHS-300 microprocessor controlled solution handling system (FiAtron Systems, Inc.) was used in this work. The SHS-300combines a variable speed peristaltic pump, injection manifold, and standard selection manifold. Operating modes for the SHS-300allow far a fixed 25 pL volume injection, timed hased injections, single or repetitive injections, or standard addition injections. For most of these studies, a flow of 1mL/min was used. The light source was a small 20-W quartz halogen projector lamp, powered hy a homemade regulated voltage upp ply. The detector used was a Bauscb & Lomh Mini Spec 20 spectrophotometer that had the lamp removed, and the fiber optic hundle was inserted into the lamp compartment using a homemade adapter to align the fiber optic bundle. The Mini Spec 20 was connected to a chart recorder through a homemade logarithmic eonverter so that all chart recordings would he linear with absorbance readings of the spectrophotometer. A Corning Model 155 pH meter with a Corning X-EL semimicro combination glass electrode was used to check the pH of all solutions prepared in this work. A FlAstar 5023 spectrophotometer (Tecator) was used to obtain the reflectance spectra of the pH indicator pads. The FIAstar 5023 has a spectral bandwidth of 13-19 nm over the wavelength range of 400-700 nm. Construction and Modifications of the Flow Cell. Details for the construction of the optosensing flow cell have been presented hy Ruzicka and Hansen (IO). Our modified cell is shown in Figure 2. The original design inserted the fiber optic bundle in a poly(viny1 chloride) (PVC) plastic block to a recessed clear plastic window. The fiber optic bundle was sensitive to changes in alignment and movement and was immobilized in the present study to eliminate these limitations in the original design. The cell design also was modified by eliminating the clear plastic
DETECTOR
/ PH PAD
\
'PVCiILrn
DIFFUSE w n i i t REFLECTOR
Flgure 2. Construction of the mcdified flow cell. The flow path is 0.5 mm i.d. window where the fiber optic bundle was inserted in order to reduce light losses due to reflection of light at the surfaces of the plastic window. Instead, Tra-Bond 156 (Tracon, Inc.) fiher optic epoxy was used to pot the fiber optic bundle into the side of the 0.5 mm i.d. channel that had been embossed into the PVC plastic block. This is a black fiber optic epoxy that was used to fix the fiber optic bundle of Crofon 1610 (Du Pont, Inc.) into the 10 mm thick PVC block. Crofon 1610 is a poly(methy1 methacrylate) 64-strand fiber optic bundle. The epoxy was allowed to cure overnight and the fiber and epoxy sticking out above the channel were carefully ground down with a grinding wheel, then sanded with 320 grit sand paper, and finally polished with 5-pm polishing paper. This allowed the fiber optic bundle to serve as the cell window and fixed all the fibers at an angle perpendicular to the flow channel and eliminated problems of alignment and movement of the fiber optic bundle. Short pieces of 0.5-mm Microline tubing (ColeParmer) were glued into the PVC block to connect the pump and waste lines to the cell. A thin PVC film was then glued over the PVC block using a polymeric glue to seal the channel. A small hole was cut in this PVC film to provide a space for the placement of the pH pad over the end of the fiber optic bundle in the cell. A pH pad was punched out of a commercially available pH strip and the pad separated from the white plastic backing. The pH pad was placed over the fiber optic bundle in the cell and then the punched out white plastic was placed over the pH pad to act as a diffuse white reflector. This was then sealed with another layer of PVC film and polymeric glue to seal the hole where the pH pad was placed. The other end of the fiber was randomly hifuracted for connection of the light source and detector. These ends were epoxied together with Tra-Bond 110 (Tracon, Inc.), an optically transparent fiber optic epoxy, to make insertion and alignment in the light source and detector easier. RESULTS AND DISCUSSION Measurement of Reflected Spectra. The narrow range Merck pH strips have two identical pads of the indicator attached adjacently on the plastic dip strip, and the wide range Merck pH strips have up to four different adjacent indicator pads attached to the dip strip. Most of the pH pads have an initial yellow or orange "uncolored" form and a red, green, blue, or purple-violet "colored" form of the indicator. Reflectance spectra were run on each of these different pads, following a preliminary spectral scan, to determine an approximate wavelength of maximum reflectance of the "colored" form. The reflected signal intensities are reported as A?,which is meant to be the apparent absorbance reading on the spectrophotometer of the reflected signal from the opaque optosensing surface with a diffuse white reflector behind the opaque pH pad. If the acid form of the indicator was "colored", a pH 13 NaOH, KCI solution (13) was pumped through the cell and the detector was zeroed a t the approximate maximum A. wavelength of the 'colored" form and a stopped flow background spectrum was obtained. Then a pH
2498
ANALYTICAL CHEMISTRY, VOL. 58, NO. 12. OCTOBER 1986
Table I. Desoriptioa and Properties of Merck pA Papers
catalog
PH
pad no.
no.
range
from top
9580 9581 9582 9583 9585 9586
0-2.5 2.5-4.5 4-7 6.5-10 11-13 C-6
1 1
9588
5-10
9587
7.5-10
9590
0-14
color (acid-base)
1 2 3
violet-yellow red-orange yellow-blue yellow-blue yellow-violet yellow-blue violet-yellow redyellow
1
yellow-red
2 3 1 2 3 1 2
3
yellow-blue yellow-purple orangepurple orangered orangeviolet orange-purple orangered yellow-blue
4
purpleyellow
1 1 1
wavelength, nm (acid or base form)”
(max)
A 569 A 592 B 640 B 623 B 592 B 646 A 595 A 575 B 586 B 632 B 608 B 589 B 580 B 583 B 589 B 578 B 637 A 583
0.358 0.618 0.676 0.401 0.495 0.739 0.517 0.319 0.460 0.659 0.546 0.412 0.341 0.461 0.410 0.395 0.617 0.382
A,
‘A is the wavelength of maximum A. for the acid form of the indicator, B is the wavelength of maximum A, for the base form of the indieator.
A
1 HCI, KCI solution (13) was pumped through the cell until the A, a t the same wavelength reached a maximum value, the flow was stopped, and the spectrum of the ‘colored” form of the indicator was obtained. A similar procedure was used for the p H papers having a ‘colored” alkaline form, except that the background spectrum was obtained with the pH 1 solution. The “colored”spectrum of the alkaline paper was obtained by usingthe pH 13 solution. The background spectra were then subtracted from the ‘colored” spectra to obtain the maximum A, wavelength and A, maxima using the software routines built into the FIAstar 5023 spectrophotometer. The results are summarized in Table I. Typical background-subtracted spectra are shown in Figure 3 for the red acid form of a Merck 9581 indicator pad for the 2.5-4.5 p H range (Figure 3A) and for the red alkaline form of a Merck 9585 pH indicator for the 11 to 13 pH range (Figure 3B). Optimizing the pR P a d Thickness. The Kubelh-Munk function (15) for the reflectance of an opaque layer is given by
f(RJ = CEs-’
, ,
..
.
, “.
B
(1)
where R. is the refledance of an infinitely thick opaque layer, C is the concentration of the species t h a t is reflecting, E is the molar absorptivity, and s is the scattering coefficient. We utilized the maximum reflected signal to optimize the thickness of the pH pad. One, two, or three pH pads were placed in the flow cell and spectra were run in the same manner previously used to obtain the wavelength of maximum A, The reflected signal was expected to increase to a plateau a t infmite thickness as the thickness of the pad was increased. Instead, there was a slight decrease in A, with increasing pad thickness. We believe that the light scattering in the Kubelka-Munk function is the predominate term in this case. This is probably due to the inhomogeneous surface of the cellulose matrix of the pH pad, leading to increased light scattering rather than increased light reflectance. The above static stopped flow results give an indication of the maximum response we can obtain from the optosensing cell with increasing pad thickness and the signal loss due to light scattering, although the system is not operated in a stopped flow mode when in use. Runs were made in a continuously flowing system to determine how light scattering and dispersion would affect the response of the system. A single pad was initially used to
Figure 3. Typical backgramd substracted spectra for (A) Merck 9581 pH pad in its red acidic form and (6) Merck 0585 pH pad in its red
alkaline form. establish a response for comparison to responses when two or three pads were inserted into the optosensing cell. With two pads in the cell, the base line moved upscale (less light sensed). When three pads were inserted into the cell, the base line again shifted upscale from the single and double pad cases. The signal peak heights also decreased as the pad thickness
ANALYTICAL CHEMISTRY, VOL. 58, NO. 12, OCTOBER 1986 ~~
was increased and the peak widths were slightly broadened with increasing pad thickness; the peaks are due to a decrease in the light sensed relative to the base line. The upscale movement of the base line with the addition of each pad is probably due to increased light absorption and light scattering with the thicker pad in the cell, therefore resulting in less reflectance. The slight broadening of the peak with increasing pad thickness is probably due to increased dispersion of the sample plug as it passes through the tortuous fiber structure of the pH pad. The reduction in the peak height can be attributed to the increased amount of indicator dye, in the alkaline form, present in the cell as additional pads are added. These additional pads are the equivalent of adding more base to the system that must be converted to the acid form of the indicator before we are able to distinguish a color change in the indicator pad with increased thickness. Because of these effects, all subsequent work was conducted with a single pH pad in the optosensing cell. Durability of the Optosensing Surface. In their earlier work, Ruzicka and Hansen (10) reported that a pH pad could be used for several weeks or thousands of injections without any deterioration in response characteristics (i.e., slope of the calibration curve, speed of response). We investigated this in more detail because of possible hydrolysis of Remazol dyes (14). Using a Merck 9581 pad for the 2.5-4.5 pH range, we employed a 0.2 M acetate buffer of pH 5.41 as carrier and a 0.2 M acetate buffer of pH 3.38 as sample. Over a period of 3 days, a total of more than 2365 injections was made on a single pad a t a rate of one injection every 39 s. When the system was shut down a t night, the cell was left filled with the carrier. A slow decrease in signal peak height was observed over this 3-day time period, and when the system was restarted each morning, the signal peak heights initially increased and then stabilized for several hours before continuing the slow decrease in peak height. A Merck 9585 pH pad for the 11-13 pH range was also investigated to check for possible alkaline hydrolysis using solutions of pH 12 and 13 made from NaOH and KC1 (13), used as carrier and sample, respectively. This pad was subjected to a few hundred injections over a period of 3 days and was stored in pH 12 solution overnight to simulate adverse conditions during this durability test. A decrease in the signal peak height was again observed during this experiment. To determine if deterioration of the pH pad or deterioration of the alkaline solutions due to absorption of COz occurred, a new pH pad was inserted. This new pad was subjected to 30 injections per day for 3 days and only a slight change in peak heights was observed, in a positive direction; the pad was also left in pH 12 solution overnight between runs. The same pH 12 and 13 solutions were used with both of these Merck 9585 pH pads. We believe that most of the change in signal with pH pads of different ranges is due primarily to the shifting of the pad in the cell, where the pad is subjected to continuous flow, rather than due to deterioration of the pH indicator on the pad. Some hydrolysis of the indicator may contribute to signal change (see below, alkaline solution studies with the Merck 9585 and 9582 pads). Removing the pad from the cell and observing it under a dissecting microscope showed a structurally intact pad when compared to an unused pad. Continued modifications to the cell design should eliminate this problem. A t present, a calibration cycle every 30-40 injections is recommended; the pH pads are very durable for use as an optosensing surface. Determination of the pK, of the Indicator. To avoid possible errors in the application of these pH pads, it is necessary to determine if the indicator has more than one pKa value. If the indicator has more than one pKa value, it may
2499
~~
Table 11. pKa Values of Merck pH Papers catalog no. 9580
9581
solutions and buffers used" 6 N HCl pH 1.22 HCl, KCl pH 1.55 HCl, KC1 pH 1.85 HC1, KC1 0.01 M NaOH
av pK, (k2s)
1.68 1.82 1.90
1.80 (*0.22)
3.73 3.83 3.90
3.82 (k0.17)
5.29 5.28
5.28 (kO.01)
0.1 N HCl
pH 2.94 KHP buffer pH 3.28 KHP buffer pH 3.49 KHP buffer 0.01 M NaOH 9582
exptl pK,
0.01 N HC1
pH 5.21 KHP buffer pH 5.51 KHP buffer pH 5.79 KHP buffer 0.01 M NaOH
5.28
OThe solutions and buffers were prepared by following the directions in ref 13. The NaOH solutions were used to establish base line AI,,- and the HCl solutions were used to determine AHInin eq 2.
mean that there are different chromophores for each pK,, leading to nonlinear responses as the indicator chromophore changes with changing pH. The changing chromophore may result in changes in the A, wavelength maximum. The pK, values for the Merck 9580, 9581, and 9582 pH papers were determined by a method described by Bishop on pp 56-59 (3). The ratio of the undissociated and dissociated indicator is determined spectrophotometrically by using the equation
where AHh is the absorbance of the undissociated form of the indicator, Amiris the absorbance of the undissociated and dissociated mixture of the indicator forms, and AI,- is the absorbance of the dissociated form of the indicator. To determine the ratio in eq 2, we substituted A, in the place of A values. The pH values of the solutions used were determined with a pH meter and a glass electrode. The pK, was then calculated as pK, = pH
- log ([In-]/[HIn])
(3)
The solutions used were pumped through the cell until a steady-state signal was obtained. The solutions and buffers used and the calculated pK, values are summarized in Table 11. All the steady-state peak heights were measured relative to the base line established for the NaOH solutions listed in the table. According to the information available in the patents (16,17), the azo dyes employed as indicators may have up to four sulfonic and/or carboxylic acid groups on the dye molecule. Our experimental data and a list of the possible dye molecules in the patents lead us to conclude that the dyes we are examining have only a single acid functional group. This allows the use of these dyes without the worry of changing chromophores and nonlinear responses, as might occur with dyes having multiple pK, values. Note, however, the wide pH range of indicator 9583 suggests there may be multiple functional groups. The closeness of pK values and possible spectral overlap preclude the resolution of these by the present experiments. Optimizing the Carrier for Low Ionic Strength Samples. Different carrier buffers were investigated to optimize the system. Carriers of 1 X 1X and 1 X M phosphate buffers were tried in order to select a buffer of sufficiently low buffer capacity so that large sample volumes
2500
ANALYTICAL CHEMISTRY, VOL. 58,NO. 12,OCTOBER 1986 6 rnin H
Ar
b
A
C
0.05 - a
b
d
C
e
TIME TIME
A
r
a
b
c
d
e
B
Figure 4. Optimization of the carrier. The responses are recorded
for a Merck 9851 pH paper. Reflectance was monitored at 565 nm M HCI solution was used as sample with a sample and a 3.4 X injection volume of 3 mL. The carriers were phosphate buffers: (a) 1 X M, pH 7.0; (c) 1 X M, pH 6.3. M, pH 7.3; (b) 1 X
would not be required to wash the buffer out of the pad and provide a rapid and large sample signal, but that would still allow a reasonably rapid return to base line (Figure 4). A Merck 9581 pH paper was used for this investigation. The results of this experiment are presented in Figure 4. The peak heights decreased with decreasing pH and the time to return to base line increased as the buffer concentration was decreased. The time to return to base line was the more important factor for our investigation. We finally chose 5 x lo4 M buffer concentration as a good compromise for our two objectives. The pH of the buffer was typically chosen to be one pH unit higher than the maximum pH range of the pH paper. This means that for a pH paper covering the range of 2.5-4.5 pH, we would use a carrier of pH 5.5. Carriers that were on the acid side of the pH range of the pH paper were tried, but a peak did not result for our samples. This is believed to be due to a filling of all the indicator dye sites with hydrogen ions, as is discussed in more detail later. Monitoring of the Colored Form of the Indicator. As previously stated, most of the pH papers have a yellow or orange color that was taken to be the “uncolored” form of the indicator. The “colored” form of the indicator was usually red, blue, or green. When a Merck 9585 pH paper was investigated, in which the transition form colored alkaline form to uncolored acid form was monitored, we experienced significant base line drift and sometimes smaller signals. NaOH, KCl solutions of pH 12 and 13 were used as sample and carrier, respectively, in this case. When the sample and carrier solutions were reversed and the system was allowed to stabilize for a few minutes, the base line was much more stable and the peak heights were increased in size from the reverse case. The lamp voltage was regulated to within 0.01 V during these studies, so it is assumed that the drift is not caused by variations in the source intensity. Some drifting of the base line was also noticed when monitoring the alkaline “colored” form of a Merck 9582 indicator. The drifting base line observed may be due to the alkaline hydrolysis of the indicator while alkaline solutions are pumped through the cell. The alkaline hydrolysis would be expected to be more significant when highly alkaline solutions are used for carrier than when alkaline solutions are injected as short sample pulses. Non-Steady-State Injections of Dilute Acid. Five dilute to 1 X M HC1 solutions covering the range of 1 X were injected onto each of the Merck 9580, 9581, and 9582 indicator pads. Our objective was to keep injected sample volumes less than 1 mL. The FIAgrams for each of these indicators are presented in Figures 5A, 6A, and 7A. The Merck 9580 indicator shows low sensitivity for the dilute acid solutions, as seen from the small differences in the peak heights. A plot of hydrogen ion activity vs. peak height had a correlation coefficient of 0.99 for the solutions used, but the relative standard deviation was f0.17 pH units. The unex-
TIME Flgure 5. FIA and steady-state response for Merck
9580 pH indicator for the 0-2.5 pH range. (A) FIAgram for the dilute HCI solutions: (a) 1 X M; (b) 7 X M; (c) 5 X M; (d) 2 X M; (e) 1 X M. A sample volume of 200 pL was used with a 5 X loT4 M sodium acetate buffer adjusted to pH 3.81 as carrier. The wavelength monitored was 540 nm. (B) Steady-state signals for large injection volumes of up to 12 ml of the previously listed HCI solutions, illustrating all the dilute dilute acid solutions, reach nearly the same steady-state signal. Time scales are the same for A and B. 5 min b
A C
d
e
0.0 5
TIME
’
B
TIME Flgure 6. FIA and steady-state response for Merck
9581 pH indicator for the 2.5-4.5 pH range. (A) FIAgrams for the dilute acid solutions listed in Figure 5A. The injected sample volume is 200 pL of acid and M sodium acetate adjusted to pH 5.50. The the carrier is 5 X wavelength monitored is 565 nm. (B) Steady-state signals for the dilute HCI solutions listed in Figure 5A. Time scales are the same for A and B.
pected result was that we saw any response at all, because this indicator is made to cover the 0-2.5 pH range, while our test solutions are in the 4-5 pH range. The Merck 9851 indicator for the pH range of 2.5-4.5 demonstrates good sensitivity for differentiating the dilute
ANALYTICAL CHEMISTRY, VOL. 58,
b
d
C
e
TIME
B
Ar a
b
C
d
e
0.3
TIME
Flgure 7. F I A and steady-state response for Mer& 9582 pH indicator for the 4-7 pH range. (A) FIAgrams for the dilute HCI soiutions listed in Figure 5A. The injected sample volume is 275 pL of HCI and the carrier is 5 X lo-' M sodium hydrogen phosphate adjusted to pH 8.13. The wavelength monitored is 570 nm. Here the disappearance of the colored form of the indicator is monitored rather than the appearance of the colored form of the indicator. (E) Steady-state signals for the dilute HCI solutions listed in Flgure 5A. Time scales are the same for A and E.
acid solutions. We are monitoring the appearance of the red acidic form of the indicator in this case. A plot of hydrogen ion activity vs. peak height had a correlation coefficient of 0.99 and a relative standard deviation of k0.02 pH units for the smallest peaks in Figure 6A. The Merck 9582 indicator for the pH range of 4-7 shows good sensitivity for differentiating the dilute acid solutions, but in this case we are monitoring the signal as the disappearance of the blue alkaline color of the indicator. A slight negative drift of the base line can be seen for the indicator monitored in this manner. The problems of monitoring the disappearance of the colored form of the indicator have been discussed previously. Even in this case a correlation coefficient of 0.99 was obtained for a plot of hydrogen ion activity vs. peak height, with a relative standard deviation of k0.04 pH units. In this work all the indicators showed a high correlation for hydrogen ion activity vs. signal peak height, but the relative standard deviation for the Merck 9580 indicator is too large for an eventual attempt to apply this technique to environmental monitoring. The drifting base line of the Merck 9582 indicator is an undesirable correction. The best compromise is the Merck 9851 indicator and monitoring the appearance of its red acidic form of the indicator. For all the indicators tried we have maintained our objective of small sample volumes, using typically 200-275 pLfr. Steady-State Signals for Dilute Acid Solutions. After unexpected responses were observed for some of the indicators, an investigation of the steady-state signals of each of the selected indicators with five dilute HC1 solutions was undertaken. Large injection volumes of up to 12 mL were used to obtain the steady-state signals in each of these experiments. The results are presented in Figures 5B, 6B, and 7B. In almost all cases, the dilute acid solutions produced nearly the same steady-state signal for each of the individual indicators. The Merck 9580 pH indicator was observed visually
NO. 12,
OCTOBER 1986
2501
to turn to its violet acidic color when pumping samples of 1 X to 1 X M HC1. Note that while these solutions represent pH values of 4-5, this particular indicator (pH range of 0-2.5) does not start to turn violet in the batch mode unless the pH is below 1.0, according to the color comparison scale provided with the indicator strips. Similar phenomena were observed with the other two indicators examined in this work. Using a separate pump and cell with a Merck 9581 pH indicator (pH range of 2.5 to 4.5) installed and observing under the dissecting microscope, the indicator turned to its red acidic form while pumping 1 x 10" M HCl (pH 4). This indicator does not turn red in the batch mode until the pH is below 3.5. When a Merck 9581 indicator strip was placed in a closed M HC1 and left for over 3 amber bottle filled with 1 x months, no red acidic color appeared in this static system. The reason for the above behaviors is not known and appears to be unique to flow systems. Perhaps when the pad of immobilized indicator is placed in a flow cell and carriers of low ionic strength are used, the immobilized indicator is acting as an ion exchanger and "preconcentrating" hydrogen ions. An attempt was made to estimate the number of moles of indicator immobilized on the pH pads used in these studies. Information to perform this is scarce. By use of the limited information available in the patents (16,17),a general recipe for the preparation of the pads was found. From the weight of paper and the weight of dye used in the recipe and with a yield of 60-8070 immobilization of the dye, we estimate there would be approximately (5-7) X lo4 mol of indicator on a pad in our cell. Using the volume of acid pumped to first reach a steady state and the concentration of acid injected, we estimated the amount of indicator on the pad to be approximately 3 X lo-' mol. The experimentally calculated value is about 50 times larger than that derived from the patent information. This suggests about 2% maximum of the acid reacts with indicator under these conditions; these amounts reacting, or smaller, may be due to diffusion or ion exchange equilibrium kinetic factors, assuming a correct approximation of the amount of dye immobilized on the pad by the manufacturer, or simply a limited supply of the indicator. The steady-state signals attained also lead us to believe that we are observing an acid-base titration of the indicator in the flow cell. Because the immobilized indicator is, in effect, being titrated by the sample, it is expected that for buffered samples the magnitude or shape of the recorded signal will be a function of the buffering capacity (acid-base neutralizing capacity) as well as the pH, depending upon the carrier buffer concentration. Recognizing this, the presently described system should be limited to fixed or low buffer capacity systems. On the other hand, it should be possible to obtain buffer capacity information from the optosensing measurements and this possibility is being explored.
ACKNOWLEDGMENT Thanks are given to FiAtron Laboratory Systems, Inc., and Tecator, Inc., for providing us with the SHS-300 and FIAstar systems used in this work. LITERATURE CITED (1) Fisher, J. E., Am. Lab. (FairfieM, Conn.) 1985, (June), 54-60.
Neisius, K. H. Kontakte (Darmstadt) 1971, 2 / 7 1 , 15. Indicators; Bishop, E., Ed.; Pergamon Press: New York, 1972. Covinaton. A. K.: Whailev. P. D.:Davison. W. Pure ADD/. , . Chem. 1985. 57(6); 877-886. McQuaker, N. R.; Kluckner, P. D.; Sandberg, D. K. Environ. Sci. Techno/. 1983, 1 7 , 431-435. ( 6 ) Sisterson, D. L.; Wurfel, B. E. Int. J . Environ. Anal. Chem. 1984, 78, 143-165. (7) Koch, W. F.; Marinenko, G. "Simulated Precipitation Reference Materials: Measurement of pH and Acidity"; Sampling and Analysis of Rain: Campbell, S . A., Ed.; American Sociefy for Testing and Materi-
(2) (3) (4) . (5)
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Anal. Chem. 1986, 58,2502-2504
ais: Phlbdelphla, 1983; ASTM STP 623, pp 10-17. (6) Galloway, J. N.: Cosby, B. J.. Jr.: Likens, G. E. J. Oceanogr. 1079. 24, 1161-1165. (9) Landsberg, H. Arch. Meterol., Geophys. B/ok/imato/.,Ser. A 1054, 7 , 219-226 (10) Ruzlcka, J.: Hansen, E. H. Anal. C h h . Acta 1985, 173, PP 3-21. (11) Fed. Reg. 1984, 49(209), 11. (12) ~ossetti,E. J. ROSSO~~I, H. J . Chem. EdUc. 108s. 42, 375-378. (13) Perrin, D.: Dempsey, B. Buffers for pH and Metal Ion Control; Chapman and Hall: London, 1974. (14) Beech, W. F. Fibre-Reactive Dyes: Logos Press Limited: London, 1970: pp 218-225.
c.:
(15) Kubelka, P.; Munk, F. 2. Tech. Phys. 1031, 12, 593. (16) Nelsius, K. H. U S . Patent No. 4029597. (17) Neisius, K. H. US. Patent No. 4029598.
R ~ E for~review D January 3, 1986. Accepted M~~ 16, 1986. This work was supported in part by NSF Grant No. ATM8318028. We also wish to thank the University of Washington Graduate Research Fund for providing funds to purchase the pH meter used in this work.
Preconcentration of Copper in Water Samples with 2-Mercaptobenzothiazole on Naphthalene Masatada Satake* and Koichi Ishida
Faculty of Engineering, Fukui University, Fukui 910, Japan Bal Krishan Puri
Department of Chemistry, Indian Institute of Technology, Hauz Khas, New Delhi 110016, India Shiro Usami
Department of Industrial Chemistry, Faculty of Engineering, Toyama University, Toyama 930, Japan
A soltd chelating materlal, 2-mercaptobenrothlazde (P-MBT), on naphthalene provides a rapid and hlghly selective means of preconcentratlon of copper from natural water samples. Copper Is quantkathrdy retained on 2-MBT-napMhalene In the pH range 5.5-7.5 and at a flow rate of 5 mumin. The metal complex-naphthalene Is dissolved from the column wlth 10 mL of n -butyiamlne-dimethytformamlde (5100 v/v) and measured by an atomic absorption spectrophotometer at 324.7 nm. Beer's law is obeyed In the concentratlon range 2.0-80.0 pg of copper in 10 mL of the flnai solutlon. Ten replicate analyses of 40 pg of copper gave a mean absorbance of 0.190 wlth a relatlve standard devlatlon of 1.0%. The sensltlvlty for 1 % absorptlon Is 0.093 pg/mL (0.133 pg/mL for the dlrect AAS method from the aqueous solution). The method has been employed for the determlnatlon of copper in varlous standard reference materials and natural water samples.
the nonaqueous organic solvents (7, 8). The only difficulty is the filtration, i.e., handling of a small amount of naphthalene, which may result in an error in the determination. In the present communication, a chromatographic method has been developed for the selective preconcentration of copper from a large volume of the aqueous phase using 2mercaptobenzothiazole (2-MBT)-naphthalene as an adsorbent. The method is very convenient (no need to filter the metal complex-naphthalene), rapid, and sensitive. The adsorbed metal in the column is not eluted even on washing with water but can be dissolved with a suitable solvent like n-butylamine-dimethylformamide from the column along with naphthalene and can be determined directly by atomic adsorption spectrophotometry. Various parameters have been evaluated and the method has been employed for the determination of copper in different standard reference materials and natural water samples and may be employed for various environmental and biological samples.
The usual liquid-liquid extraction method cannot be employed directly for the extraction of metal ions that form complexes with the complexing reagent at a high temperature ( I , 2). This difficulty can be overcome with the method of solid-liquid separation after liquid-liquid extraction developed by Fujinaga and co-workers using naphthalene as an extractant ( 3 ) . Although this method has many advantages over the usual liquid-liquid extraction (4,5),it is inconvenient since the operation is carried out at a high temperature and cannot be applied for the extraction of metal ions that form thermally unstable complexes. In order to overcome these drawbacks a second method, solid-liquid separation after adsorption of metal chelates on microcrystalline naphthalene, was developed (6). The method is very convenient (carried out a t room temperature), rapid (no need of heating naphthalene), sensitive, and economical (hardly 0.4 g of naphthalene is needed), can be applied to many types of the complexes, and is especially useful for metal complexes that have low solubility in
EXPERIMENTAL SECTION Reagents. Copper sulfate solution was prepared in double distilled water from its analytical reagent grade sample and standardized (9). A more dilute solution ( 5 ppm) was prepared by taking the appropriate amount of the standard solution. Buffer solutions were prepared by mixing 1 M acetic acid and 1 M ammonium acetate solution for pH 3-6 and 1M aqueous ammonia and 1 M ammonium acetate solution for pH 8-11. N-Butylamine-dimethylformamide solution was prepared by mixing 5 mL of the amine with 100 mL of DMF. Naphthalene, 2mercaptobenzothiazole (2-MBT), DMF, and all other reagents were of analytical reagent grade. Preparation of Loaded 2-Mercaptobenzothiazole-Naphthalene Mixture. Naphthalene (20 g) and 2-MBT (3 g) were completely dissolved in 100 mL of acetone. This solution was transferred to 300 mL of water in a fast stream and stirring magnetic stirrer-hot plate arrangement at 50 "C. Naphthalene coprecipitated along with 2-MBT waa filtered through filter paper by suction, washed with water, dried in air for 2 days, and then stored in a bottle. Apparatus. A Toa-Dempa HM-5A pH meter and a PerkinElmer, 403 atomic absorption spectrophotometer were used. A
0003-2700/86/0358-2502$01.50/00 1986 American Chemical Society
