2564
Anal. Chem. 1985, 57, 2564-2567
Spectrophotometric Determination of Seawater pH Using Phenol Red Gillian L. Robert-Baldo, Michael J. Morris, and Robert H. Byme* Marine Science Department, University of South Florida, LI3t. Petersburg, Florida 33701
Phenol red absorptlon spectra, In conjunction with thermodynamic characterization of the equlllbrlum dye (yellow) += dye (red) 3- H", provldes the basts for spectrophotometric pH determlnatlons In haturaf seawater. Phenol red Is particularly well suited to pH determlnatlons In the cold seawater which comprises most of the ocean's volume. Within the bounds 33%0 IS I37%0 and 273 K IT I 303 K at 1 atrn total pressure, the phenol red equlllbrlum constant, K,, can be represented by the expression -log K, = (4834.00/T) 84.831 3- 30.758 log T f 0.004(35%0 S%o).
-
-
Seawater pH is an essential component in quantitative descriptions of ocean chemistry. Oceanic processes and properties which are strongly pH-dependent include mineral solubility ( I ) , dissolution kinetics (2),chemical speciation and adsorption ( 3 , 4 ) ,bioavailability (5),and redox kinetics and equilibria (3, 6). Buffering of seawater pH is, on a short geological time scale, dominated by the oceanic COz system. As a consequence of the direct relationship between pH and COZ partial pressure, examination of seawater pH using hydrogen electrodes is impractical, and such measurements are predominantly conducted with pH-sensitive glass electrodes. Potentiometric cells containing glass electrodes have, in general, proven satisfactory in measurements of solution pH. However, the behavior of potentiometric cells in seawater pH measurements often falls short of what is deemed acceptable. Even with the most careful treatment, the potential of cells containing glass electrodes often drifts slowly with time (C0.6 mV h-l) after such cells me placed in a new solution (7). Drift of cell potentials makes rapid measurement of pH quite difficult and is an especially severe problem in investigations dependent on precise observation of small pH differences. Measurements involving cells with liquid junctions are subject to further uncertainties due to the dependence of liquid junction potentials upon medium concentration and composition. Ideally, the change in liquid junction potmtial (residual liquid junction potential) between test solution and standardizing buffer should be small or at least highly reproducible. In practice, systematic errors between many pH and pK measurements suggest (a) that the reproducibility of the residual liquid junction potential is often poor and (b) that residual liquid junction potentials are dependent on the construction and/or history of the liquid junctions used in various investigations (7). As an alternative to potentiometric measurements, spectrophotometric techniques offer a variety of advantageous characteristics in determinations of seawater pH. Equilibration rates are dependent on aqueous molecular processes and are generally quite fast. Measurements can thereby be conducted rapidly, minimizing the contact time between seawater and surfaces susceptable to ion exchange. Spectrophotometric pH measurements in seawater are quite stable over periods of time much longer than those necessary for equilibration. This attribute, in conjunction with the sensitivity of spectrophotometric methods, allows precise exami-
nation of small differences in pH. Seawater is an ideal medium for the employment of pHsensitive spectrophotometric dyes. The physical/chemical characteriRtics of indicator dyes are expected to vary considerably with medium composition, concentration, and temperature (8). In seawater, however, medium composition is virtually invariant (9, l o ) ,concentration varies within a relatively small range (34 IS%O5 36 for 98% of the ocean's volume) ( l l ) and, , in much of the ocean, both salinity and temperature are substantially constant (0 5 t I6 "C and 34 IS%o I35 for 79% of the ocean's volume) (12). The uniform composition of ocean water (constancy of major component molar ratios) and very low concentrations of reactive trace elements (13, 14) allow characterizations of an indicator's physical/chemical behavior to be expressed in terms of only three variables: temperature, pressure, and salinity. In this work, we have quantitatively examined the behavior of phenol red in seawater a t 1 atm total pressure and temperatures between 5 "C and 30 "C. Phenol red was chosen for our investigation because its working range makes it especially suitable for use at the low temperatures typical of most of the ocean. Our characterizations of phenol red were engendered by our interest in the dissolution kinetics of marine aragonitic particulates (15, 16) and by our interest in the respiration rates of small individual zooplankton (17). Through the reactions CaCOJs)
+ H+ * Ca2+ + HC03-
(1)
and
COz + HzO + H2C03
H+ + HC0,-
(2)
seawater pH provides a sensitive indicator of carbonate dissolution and evolution of COP Use of pH-sensitive indicator dyes allow such processes to be conveniently examined in closed systems.
THEORY Conversion of the yellow, Dy, form of phenol red to the red, D R , form is accompanied by the appearance of a strong absorbance band centered at 558 nm. The absorbance changes which accompany the interconversion of yellow and red forms Dy
F+
DR
+ H+
(3)
can be quantitatively described by using the equation (18)
where hA is absorbance at wavelength A, DT is total dye concentration (DT= [Dy] + [DR]), is the absorbance per molar centimeter of DYat wavelength A, is the absorbance per molar centimeter of DR at wavelength A, K1 is the dissociation constant appropriate to reaction 3
0003-2700/85/0357-2564$01.50/00 1985 American Chemical Society
(5)
ANALYTICAL CHEMISTRY, VOL. 57, NO. 13, NOVEMBER 1985
1 is path length, and brackets denote the concentration of each chemical species. Equation 4 can be converted to a form in which pH rather than ha is the dependent variable
where pK1 = -log K1, pH = -log [H+], and we note that [H+] is the free hydrogen ion concentration in moles per kilogram of HzO. Equation 6 can be further simplified by observing that at a dye concentration DT !
where AA = AaDT1,hAMMax = h e ~ D ~and 1 , h A ~ i= n AEYDT~. Equation 7 is a somewhat more convenient formulation than eq 6 because direct measurement of the product heRDT1 obivates the need for independent measurements of DT and The product hAAILlan = heRDT1 is easily measured as the observed absorbance when, at high pH, DT = [DR].hAMin is measured as the observed absorbance when, a t low pH, DT = [DY]. Accordingly, in separate measurements of hA, h A k ,and xA-, eq 7 can be employed to determine pH by holding DT1 constant. A further convenience of this procedure is derived from the observation that in a constant diameter, variable path length cell (14-21) the application of pressure induces changes in both DT and 1 while maintaining a constant product D T ~ .
EXPERIMENTAL SECTION Our examinations of pK, as a function of temperature were conducted with seawater collected in the Gulf of Mexico. The salinities of our stock seawater media were determined with a Model 8400 Autosal salinometer. Phenol red stock solutions were prepared in deionized water using Eastman Kodak ACS phenol red in the sodium form. Phenol red stock solutions were added to each seawater medium at less than 1 to 600 mixing ratios. Each medium was further diluted with deionized water so that the final salinities of our t,est media were 35.0 f 0.1% with phenol red concentrations near 2 pmol/L. Test media were housed in an open top, 10-cm path length spectrophotometric cell which fit snugly in the thermostated ( f O . l "C) well of a Cary 17D spectrophotometer. Ports in the lid of the spectrophotometric cell permitted the simultaneous introduction of an Orion 8102 combination pH electrode (3 M KCl filling solution) and a mechanically driven stirring rod. This arrangement permitted the simultaneous measurement of both pH and absorbance. Measurements of pH were obtained with a Corning Model 130 pH meter and, through the use of Tris (0.0400 mol kg-l) seawater buffer (22),are reported on the free hydrogen ion concentration scale. The absorbances of phenol red in seawater were measured against 35.0% seawater reference solutions to which no phenol red had been added. HC1 (1 M) was added to the spectrophotometric titration vessel by use of 2 mL total capacity, Gilmont microburets. The series of titration points (558A,[H+])thereby obtajqed were fit by using eq 4 in a nonlinear least squares analysis (8, 23, 241, providing the phenol red dissociation constant, K1. Subsequent to completion of 21 experiments using the procedures outlined above, an alternative procedure for the determination of pK1 was developed. As a minor modification of the techniques described above, titration points were obtained at pH as high as 9.3. Our high pH data were obtained through addition of B(OH),O/NaB(OH), titrant solutions which were composed as NaOH-sodium borate decahydrate mixtures in deionized water. Our borate/boric acid mixtures were adjusted to produce upper bound pHs on the order of 9.4. This procedure obviated generation of brucite (Mg(OH),) in the high pH microenvironments which occur upon direct addition of NaOH. Finally, since the boric acid/borate titrant solutions raised the ionic strength of our solutions by about 1%,deionized water was used to obtain a formal ionic strength equivalent to 35%0seawater. The data
2565
Table I. Dissociation Constants of Phenol Red in Natural Seawatera as a Function of Temperature
4
"C 5.7 5.8 10.4 11.2 14.6 15.4 20.0 20.2 25.1 25.2 30.1 30.3
PK1 PK1 (data set l)b,c (data set 2)b*d
no. of experiments
7.724 f 0.002 7.716 f 0.018 7.653 f 0.003 7.645 7.605 f 0.003 7.587 f 0.003 7.546 f 0.002 7.543 7.495 f 0.006 7.484 & 0.002 7.451 f'0.006 7.437 f 0.006
2 3 3 1 2 3 3 1 6 2 3 2
"35%0 salinity and 1 atm total pressure. bThe range provided with each pK, indicates half the range of the individual observations at each temperature. CDataset 1 pK's were derived by use of eq 4. dData set 2 pK's were derived by use of eq 7.
obtained by use of these procedures were fit using eq 7, as follows: For each data set, 54k was initially estimated by using the 558A value obtained at the data set's highest pH. Awn was calculated by using the relationship
By use of titration data (&dvs. pH) and initial estimates of Ah and A&, pK, was calculated for each data point within the bounds 7.0 I pH I 8.3. The pK1 data so obtained were averaged and the result was used in conjunction with eq 7 to obtain refined estimates of AMm and AMin. This procedure was repeated until no further changes in pKl were observed.
RESULTS AND DISCUSSION The pK, results obtained in our investigations using eq 4 and eq 7 are shown in Table I. It is expected that the variation of pK with the thermodynamic temperature (2' = t + 273.15) can be represented by equations of the form (22)
Using eq 8 in a nonlinear least squares analysis, our entire data set was used to obtain the parameters A, B, and C. The results of this analysis are provided as eq 9
pK1= 4834'00 - 84.831 + 30.7580 log T ~
T
(9)
Equation 9 provides a good, smoothed description of our K1 results expressed in moles per kilogram of water. The pKl behavior predicted by using eq 9 is shown as the solid line in Figure 1. The maximum residual shown in the figure is 0.0066 and the average residual is 0.004. The parameter estimates shown in eq 9 can be used to provide estimates of the thermodynamic properties of phenol red in seawater. The standard changes of enthalpy (AH"), entropy (AS"), and heat capacity (AC,") for the dissociation process in natural 35% salinity seawater are given (22) as
AHo = R(A In 10 - C T )
ASo = -R(B In 10 + C ACpo = -RC
+ C In r )
(10) (11)
(12)
where R = 1.9872 cal mol-l K-l. Table I1 shows the AH", AS", and ACpo values obtained a t 25 "C and 0 "C using eq 9. The uncertainties shown in Table I1 encompass the range of values obtained when eq 8
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ANALYTICAL CHEMISTRY, VOL. 57, NO. 13, NOVEMBER 1985
1000/T ( K - l ) Figure 1. Temperature dependence of the phenol red dissociation constant: 0, data set one; A, data set two.
Table 11. Calculated AHo, A S o , and ACpoValues for Phenol Red at 0 and 25 OC
T,K
AHo, ea1 mol-'
ASo, cal K-' mol-'
AC,O
273.15 298.15
5423 f 130 3895 f 280
-15.9 f 0.8 -21.2 f 1.4
-61 f 16 -61 & 16
and 10-12 are applied to our first, second, and combined data sets. The uncertainties provided with the central estimates indicate that AH" and AS" are much better defined than
AC,'. In order to generalize eq 9 to an appropriate range of oceanic conditions, our characterization of phenol red in seawater was extended to include the influence of salinity on pK1. Our investigations of pKl vs. S%owere conducted at 25 "C and 13 "C over a salinity range between 35.9% and 31.7%. With eq 7, 55d and pH were simultaneously measured as salinity was reduced with deionized water. By accounting for the effect of dilution on A~ and A&, pKl was conveniently determined as a function of salinity. The result obtained in three experiments at each temperature, ApK1/AS%o= -0.004, indicated that a 1% salinity decrease results in a 0.004 increase in pK1. No significant difference in this result was observed as a function of temperature. Consequently, our analyses indicate that the behavior of phenol red in natural seawater can be reasonably 'described by using the equation PKl =
-4834'00 84.831 + 30.7580 log T + 0.004(35%0- S%o) T (13)
Although use of this equation at temperatures lower than 5.7 "C or salinities greater than 35.9% entails a slight extrapolation, the very small salinity dependence of pK1 and the slight curvature of pKl vs. T1 indicate that eq 13 can be reliably used between 0 "C and 35 OC and at salinities between 33%0and 37%. This contention is supported in part by the pKl determinations of Sendroy and Rodkey (25) in 0.15 M NaCl at 25 "C and 37 OC. Their result, ApK1 = pK1 (25 "C) - pK1 (37 "C) = 0.095 0.012, can be compared with our result, ApK1 = 0.100, calculated by extrapolating eq 9 to 37 OC. In spectrophotometric pH determinations using eq 7 and 13, it is important that the parameter *AMax is accurately assessed. Our analyses indicated that, in aqueous media, 558AMax exhibits very little, if any, dependence on medium
*
composition. Consequently, 558AMex can be conveniently determined a t a given temperature through measurements in either seawater or NaC1. Additions of phenol red stock medium to 0.7 M NaCl media at pH 11.0 allow direct determination of 55$iMax. As an alternative, phenol red can be added to seawater at high pH and eq 7 and 13 can be used to calculate 558AMax. In either case, it is, of course, important that dilutions of the phenol red stock medium are well defined. Our examination of 5&hIin supports the contention (25)that 558AMin can generally be set equal to zero in eq 7. Differences between pH determinations which use the approximation 558AMin 0 and determinations using the more refined estimate, 558AMin/558AMax = 1.2 X are quite small. Even at a pH as low as 7.2, well outside the pH range of normal seawater, the differences between pHs calculated using 55$iMin 0 and 5 5 8 A ~ i ~ / 5 5=d1.2 ~ ~X~ are on the order of 0.002 pH units. Due to the relatively small salinity dependence of pK1, spectrophotometric pH determinations employing eq 7 and 13 do not require detailed salinity measurements. However, in employing eq 13, it should be noted that measurements reported to fO.OO1 pH require temperature control to within f0.1 "C. Furthermore, although the temperature dependence of mA% is very much smaller than that of pK,, it is advisable to obtain 558AMax at the temperature used in determinations of 658A. In addition to the procedures outlined above which require absorbance measurements at a single wavelength, it should be noted that measurements at two or more wavelengths may offer significant advantages in some applications. The term (hA - hAMin)/(AAMax - hA) in eq 7 is equal to the concentration ratio, [DR]/ [DY]. Spectrophotometric measurements at two or more wavelengths permit direct calculation of the ratio [DR]/[DY] through examination of spectral shape (xA VI. wavelength) rather than comparison of ,,A and *Alllax. In a multiwavelength analysis, the concentrations [DR]and [DY] and their ratio can be calculated by using equations of the form at two or more wavelengths simultaneously (26). In the case that only two wavelengths are used in an analysis of phenol red spectra, particularly appropriate choices are 558 nm, the htR molar absorptivity maximum, and 433 nm, the hty molar absorptivity maximum. In this case, the ratio [DR]/[DY]is given as
where R = 658A/483A9 co = S E & Y / ~ ~ & , CI = x i & d m c ~ and cz = 433tR/43&. Since all of the terms on the right hand side of eq 15 involve absorbance ratios, or molar absorptivity ratios and red (XtR) forms of dye, meaappropriate to yellow surement of pH is rendered independent of dye concentration and path length. Spectrophotometric methods add considerable versatility to the determination of seawater pH. Convenience of measurement in closed systems, sensitivity, and rapid response are important general attributes of spectrophotometric pH determinations. Rapid response is of special significance when rapid data acquisition is important. The general attributes of spectrophotometric pH determinations appear particularly well suited to the in situ examination of fine structure in ocean pH profiles. Registry No. H20, 7732-18-5; phenol red, 143-74-8.
LITERATURE CITED (1) Garrels, R. M.; Christ, C. L. "Solutlons,Minerals and Equilibria":Harper and Row: New York, 1965.
Anal. Chem. 1985, 57, 2567-2570 (2) Morse, J. W.; Berner, R. C. I n “Chemical Modeling in Aqueous Systems”; Jenne, E. A., Ed.; American Chemical Society: Washington, DC, 1979; ACS Symposium Series No 93, Chapter 24. (3) Stumm, W.; Morgan, J. J. “Aquatic Chemistry”, 2nd ed.; Wiley: New York, 1981. (4) Baes, C. F., Jr.; Mesmer, R. E. “The Hydrolysis of Cations”; Wiley: New York, 1976. (5) Morel, F. M. M.; Morel-Laurens, N. M. L. I n “Trace Metals in Sea Water”; Wong, C. S., Boyie, E., Bruiand, K. W., Burton, J. D., Goldberg, E. D., Eds.; Plenum Press: New York, 1983; pp 841-869. (6) Kester, D. R.; Byrne, R. H.; Liang, Y. I n “Marine Chemistry in the Coastal Environment”; Church, T. M., Ed.; American Chemlcal Soclety: Washington, DC, 1975; ACS Symposium Serles No. 18, pp 50-79. (7) Culberson, C. H. I n “Marine Electrochemistry”; Whitfield, M., Jagner, D., Eds.; Wiiey: New York, 1981; pp 187-261. (8) Bates, R. G. “Determination of pH Theory and Practice”, 2nd ed.; Wiley: New York, 1973. (9) Sverdrup, H. U.; Johnson, M. W.; Fleming, R. H. “The Oceans: Their Physics, Chemistry and General Biology”; Prentlce-Hall: Englewood Cliffs, NJ, 1942. (10) Miliero, F. J. I n “The Sea”; Goldberg, E. D., Ed.; Wiley: New York, 1974; Vol. 5, Chapter 1. (11) Montgomery, R. B. Deep-sea Res. 1958, 5 , 134-148. (12) Plckard, G. L.; Emery, W. J. “Descriptive Physical Oceanography”, 4th enlarged ed.; Pergamon Press: Oxford, 1982. (13) Brewer, P. G. I n “Chemical Oceanography”, 2nd ed.; Riley, J. P., Sklrrow, G., Eds.; Academic Press: London, 1975; Vol. 1, Chapter 7. (14) Bruland, K. W. I n “Chemical Oceanography”; Riley, J. P., Chester, R., Eds.; Academic Press: London, 1983; Voi. 8, Chapter 45.
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(15) Byrne, R. H.; Acker, J. G.; Betzer, P. R.; Feeiy, R. A,; Cates, M. H. Nature (London) 1984, 372, 321-326. (16) Betzer, P. R.; Byrne, R. H.; Acker, J. G.; Lewis, C. S.; Joiiey, R. R.; Feely, R. A. Science 1984, 226, 1074-1077. (17) Morrls, M. J.; Byrne, R. H. Trans. Am. Geophys. Union 1982, 63 (3), 101. (18) Byrne, R. H.; Kester, D. R., J . Solution Chem. 1978, 7 , 373-383. (19) ie Noble, W. J.; Schiott, R. Rev. Sci. Instrum. 1976, 4 7 , 770-771. (20) Byrne, R. H. Rev. Sci. Instrum. 1984, 55, 131-132. (21) le Noble, W. J.; Schiott, R. Rev. Sci. Instrum. 1984, 55, 132. (22) Ramette, R.; Culberson, C. H.; Bates, R. G. Anal. Chem. 1977, 4 9 , 867-870. (23) Marquardt, D. W. J . SOC. Ind. Appl. Math. 1963, 7 1 , 431-441. (24) SAS Instltute, Inc. “SAS Users Guide: Basics”, 1962 ed.; SAS Institute Inc.: Cary, NC, 1962. (25) Sendroy, J., Jr.; Rodkey, F. L. Clin. Chem. (Winston-Salem, N.C.) 1981, 7, 646-654. (26) Byrne R. H.; Young, R. W.; Miller, W. L. J . Solution Chem. 1981, 10, 243-251.
RECEIVED for review February 11,1985. Accepted June 17, 1985. This research was supported by a University of South Florida faculty research grant, by a grant (NA80RAD00020) from the National Oceanic and Atmospheric Administration Air Resources Laboratory, and by the Gulf Oceanographic Charitable Trust (G. Robert-Baldo Graduate Fellowship).
Performance Tests for the Measurement of pH with Glass Electrodes in Low Ionic Strength Solutions Including Natural Waters ,William Davison* and Colin Woof Freshwater Biological Association, The Ferry House, Ambleside, Cumbria, United Kingdom
The performance of different commerclal reference eiectrodes, used to measure the pH of low Ionic strength natural waters, has been tested agalnst a renewable free diffusion Junction. Comparatlve measurements on synthetlc solutions were partially successful In predlctlng performance. Errors determlned In dilute buffers, dilute acids, and distilled water were simllar to those observed in natural waters, but the response In NBS buffers falled to reveal any problems. Poor electrodes depressed the pH from Its true value, more so in stlrred rather than quiescent solutions. The shlR In pH upon stirrlng was largest for the worst electrodes In the most dilute solutions. Better electrodes were characterized by high flow rates of Internal fllling solution through the Junction. As no one test could guarantee electrode performance, the adoption of a standard reference procedure to test electrodes was recommended.
Use of normal electrodes and established procedures for the measurement Of the pH Of low ionic does not guarantee the accuracy (1-6). The prime source of error is associated with the liquid junction of the reference electrode. Great care has been taken in establishing well-defined pH scales. If measurements are made with cell 1
I
reference KC1 unknown ( X ) or electrode > 3 . 5 mol dmW3 standard (S)
H+ion responsive electrode
,1 ,1 1,
they obey the conventional definition of pH and so they are
assumed to be “correct” (7). The measurement is calibrated by substituting a standard buffer solution ( S )for the unknown solution (X), where the pH of the standard buffer is conventionally assigned. This paper is concerned with the reproducibility of the pH measurement according to this convention. Interpretation in terms of the concentration or activity of H+ is left to others (8). Problems which arise are largely associated with the irreproducibility of the liquid junction (represented by the double line in cell 1). This is partly because definitions of pH do not explicitly state the type of junction which should be used.for practical measurements. However, as long ago as 1930 Guggenheim (9) tested various junctions and concluded that only a free diffusion junction (FDJ) could be relied on to give a reproducible potential approaching the ideal value. Free diffusion junctions are in fact incorporated into the pH scales. The BS scale uses such a junction for determining the pH of different standard buffers. The NBS scale uses it to check that the residual liquid junction potential between different buffers is minimal.. Thus a F D j is an obvious choice as a standard junction for precise measurements and it has been recently used as such by Culberson (IO), Illingworth ( I I ) , Brezinski (121,and Covington et al, ( I 3 ) . But what of other junctions? Most commercial electrodes use restrained junctions such as a ceramic or fiber plug or a sleeved joint. Although these have been shown to sometimes give very inaccurate results (11-14), some examples apparently perform well (4). At present there is no simple test which can be applied as a quality control procedure to assess the accuracy of a measurement made witha restrained junction. In this work a wide range of performance characteristics of a variety
0003-2700/85/0357-2567$01.50/0 0 1985 American Chemical Society
