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Ind. Eng. Chem. Res. 2007, 46, 4734-4739
APPLIED CHEMISTRY pH Measurements in Monoethylene Glycol (MEG) + Water Solutions Kristian Sandengen,* Baard Kaasa, and Terje Østvold Department of Materials Science and Technology, Norwegian UniVersity of Science and Technology (NTNU), 7491 Trondheim, Norway
A convenient calibration method for pH measurements in monoethylene glycol (MEG) + water solvents is described. The calibration has to be performed once for each electrode type. pH measurements were performed in CO2(g, ∼1 bar)-saturated solutions of 60 and 90 wt % MEG at 4-80 °C with varying NaHCO3 and KHCO3 contents. Reproducible data were obtained using several different electrodes. Theoretical aspects are reviewed. Introduction Transportation of hydrocarbons and water in long subsea flow lines results in new challenges in the control of hydrate, corrosion, and scale formation. As the fluids cool, water condenses, and gas hydrates can form. Hydrate formation can be inhibited by an antifreeze agent such as monoethylene glycol (MEG, or 1,2-ethanediol). The presence of MEG changes the thermodynamics of the aqueous phase and lowers the solubility of most salts. MEG also has an influence on the pH, gas solubility, and CO2 acidity. During work with mineral precipitation or scale predictions in MEG + water mixtures, a need for accurate pH measurements in such solutions arose. In the literature, several reports have been published1-9 concerning the calibration, measurement, and interpretation of pH in solutions containing MEG. In the present work, the theoretical foundation for pH measurements in mixed solvents is reviewed. A method is suggested for how to treat pH in practical problems such as scale control. The aim is to minimize variations in the measured values when using different types of electrodes in water + MEG solutions. This work is based on calibration data reported by Mussini et al.1-3 Theoretical Aspects pH Electrodes. A typical combined glass electrode is shown in cell I, where the liquid junction is indicated by double vertical lines and the boundaries of the glass membrane are indicated by single vertical lines.
sensing electrode|test solution||reference electrode Ag, AgCl, H+, KCl (aq)|glass|test solution||KCl (3 M), AgCl, Ag (cell I) The sensing electrode consists of a pH-responsive glass with a buffer solution of constant pH on the inside. KCl (typically 3 M or saturated) solution in contact with silver/silver chloride is a common reference electrode. To create electrical contact it is necessary to introduce a liquid junction. This is typically made of a ceramic material that allows a small flow of KCl solution * To whom correspondence should be addressed. E-mail: krian@ statoil.com. Present address: Statoil Stjørdal, Strandveien 4, 7501 Stjørdal, Norway. Tel.: +47 74 86 20 00.
from the reference electrode into the test solution. This electrode yields the measured pH, pHmeas, from the measured potential, Emeas, according to the equation
pHmeas ) (Eo - Emeas)
F sR′T
(1)
where the subscript meas corresponds to the measured value and R′ ) R·ln 10. The two unknown parameters E°(standard potential) and s (sensitivity) can be found by measuring Emeas in at least two solutions with known pH values. Several calibration solutions are available that have been given a specified pHS (pH standard) value, as described in detail by Covington et al.10 and Buck et al.11 When pHmeas is obtained according to eq 1, E° is assumed to remain constant during calibration and measurement. This is an important factor for pH measurements in water + MEG solutions. It is useful to regard the standard potential E° as being composed of inner and outer contributions. The former corresponds to the inside of the electrode and depends on the type of sensing and reference electrodes, as well as the KCl concentration. The latter describes the contact between the interior of the electrode and the test solution, i.e., at the glass membrane and at the liquid junction. At constant temperature, the inner potential is truly constant. It will remain constant as long as evaporation or other processes do not alter the Clactivity in the reference electrode (see cell I). The outer potential, however, depends on the composition of the test solution. Hence, if the solution is much different than the calibration solution, the assumption of constant E° can lead to errors. If E° changes between calibration and measurement, the measured value of pHmeas will not equal the actual value pHactual. This can be expressed by the general relation
pHmeas ) pHactual + ∆pHERR
(2)
where ∆pHERR denotes the error caused by the change in E°. If, for example, the salinity of the test solution is much higher than that of the standard (pHS), E° changes as a result of the different potential across the liquid junction, ∆ELJ°, and possibly also over the glass membrane, ∆Eg°. Introducing these terms into eq 2 yields
10.1021/ie061305a CCC: $37.00 © 2007 American Chemical Society Published on Web 06/08/2007
Ind. Eng. Chem. Res., Vol. 46, No. 14, 2007 4735
pHmeas ) pHactual -
z ) T - θ/T,
F (∆Ego + ∆ELJo ) sR′T
) pHactual - (∆pHg + ∆pHLJ)
(3)
The ∆ term denotes a change relative to the standard solution. During practical use, it is impossible to separate the values arising at the glass membrane and liquid junction. Thus, the effects of the solution on the glass membrane, ∆pHg, and the liquid junction, ∆pHLJ, are included in a general ∆pH term, i.e.
pHmeas ) pHactual - ∆pHg - ∆pHLJ ) pHactual - ∆pH
(4)
where ∆pH describes the net effect, i.e., the difference between actual and measured pH values. Several factors can contribute to this error. A change in salinity and/or the presence of another solvent such as MEG are usually the most important. It is therefore valuable to divide the total ∆pH term into two other contributions, one due to the salinity of the test solution, ∆pHSalt, and one due to the MEG content, ∆pHMEG. Rearranging eq 4 and denoting the actual pH, pHactual, as pH gives
pH ) pHmeas + ∆pHSalt + ∆pHMEG
(5)
The effect of salt on the liquid junction potential (∆pHSalt) has been discussed by Bagg.12 Both Kan et al.4 and Sandengen13 measured ∆pHSalt in water and thereafter used the approximation that ∆pHSalt could be regarded as independent of the MEG concentration. In the present work, however, we focus only on the effect of MEG, i.e., ∆pHMEG. MEG Dependence. In a mixed solvent, measuring difficulties are caused by the outer potential, which does not remain constant when the solvent varies. If an electrode is calibrated in aqueous solutions and thereafter used for measurements in MEG + water solutions, the outer potential changes. If, however, the electrode is calibrated in a solution of the same MEG content (and ionic strength) as the test solution, the outer potential obviously remains constant between calibration and measurement. Buffer solutions of 0.05 m KHPh (potassium hydrogen phthalate) have been extensively studied and are designated as the reference value pH standard (RVS).11 Mussini et al.1-3 measured pHRVS for 0.05 m KHPh buffer in MEG + water solutions. Their experimental data1 were revised in a more recent publication3 and are the only standards available for MEG + water at present. A function for pHRVS must model these measured values, as well as the pHRVS values of aqueous 0.05 m KHPh in the desired temperature range. In addition, the function must have a reasonable extrapolation outside the experimental data (0-70 wt % MEG). The equation given by Mussini et al.,3 however, has an error in the temperature function and provides a somewhat strange extrapolation at high MEG concentrations. Thus, we fitted a new function from the aqueous data (0-95 °C) given by Covington et al.10 and the abovementioned mixed-solvent data.3 Equation 6 is of the same type as originally used by Mussini et al.1 and generally reproduces the data to within 0.005 pH units
pHRVS ) 4.00249 + 1.0907wG + 0.9679wG2 + 0.3430z + 0.03166wGz -0.8978wG2z + T T - z + 9.8795wG3 ln - z (6) 7.7821 ln θ θ
{ () }
where
{ () }
θ)298.15
and wG is the weight fraction of ethylene glycol in the salt-free solvent. The ∆pHMEG term is found by measuring pHmeas in this KHPh standard solution, which has the value of pHRVS given by eq 6
∆pHMEG ) pHRVS - pHmeas
(7)
pH Reference State. The pHRVS value in eq 6 actually has a different thermodynamic reference state for each MEG concentration. This implies that the H+ activity (pH) measured in a MEG-containing solution is not directly comparable to the value measured in water. Thorough descriptions of this topic can be found in the works of Bates et al.,7,9 Mussini and Mazza,5 and Sandengen.13 During practical work, however, these different references are usually not important; thus, further discussion will not be given herein. When measuring pH in MEG-containing solutions, the primary concern is to be able to reproduce the measurement using a different electrode system. This is achieved using the calibration method described in this work. Experimental Section MEG (p.a. 99.9%, Acros) is hydroscopic and was analyzed for moisture with a Methrom 831 KF Karl Fischer titration apparatus. It contained less than 500 ppm (0.05 wt %) water. Combined glass electrodes of the KCl (3 M) type (Mettler Toledo DG111-SC) and the KCl (saturated) type (Radiometer pHC2011-8) were used at 4-50 °C. At 50-80 °C, special hightemperature glass electrodes from Innovative Sensors (GT-DJ) and a KCl bridge glass electrode from Metrohm (Unitrode) were used. The pH electrodes were first calibrated with standard aqueous IUPAC solutions from Radiometer (pH ≈ 4, 7, and 10) at temperatures of 4, 25, 50, and 80 °C. Then, the electrodes were put into 0.05 m (moles per kilogram of solvent) solutions of potassium hydrogen phthalate (KHPh) with known MEG + water concentrations (0-100 wt %). These MEG-containing buffer solutions were prepared from dried (100 °C, 3 h) KHPh salt (p.a. 99.8% Merck). Application and Testing. A three-neck round-bottom flask (250-500 mL) was filled with a MEG + water solution of known composition (60 or 90 wt % MEG). The solution was magnetically stirred. A water bath was used for temperature control (4, 25, 50, 80 °C). CO2 was continuously bubbled through the solution with the flask being open to atmospheric pressure. When the pH reading became stable (∼1 h), known amounts of NaHCO3 or KHCO3 were added to the solution. pH was recorded after each addition, and the procedure was continued until the solubility limit was reached (4-48 h). CO2 was bubbled through the solutions at a flow rate of 20-50 mL/ min. The gas was passed through cleaning bottles with the same MEG + water content as the sample before the inlet of the reaction flask. This was done to minimize evaporation of solvent during the experiments. The CO2 partial pressure was calculated by subtracting the MEG + water vapor pressure14-15 from the atmospheric pressure. Literature data14,15 for vapor pressure are available only for salt-free solution. When salt was present, it was decided simply to calculate the water + MEG vapor pressure according to Raoult’s law, i.e., to assume that vapor pressure is proportional to the mole fraction of solvent, which is slightly reduced in the presence of salt. The solutions in this work were not highly saline, and the influence on the vapor
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Figure 1. pHmeas compared to actual pH (pHRVS, solid line) of the KHPh standard solution at 25 °C as a function of xMEG: ([) aqueous KCl bridge electrodes, (∆) low-temperature electrode from Jenway with an EtOH-based electrolyte.
seemed to be reasonable. Temperature is another concern. pH is usually measured at room temperature (∼20 °C). At higher temperatures, however, ∆pHMEG is different. Thus, for pH measurements far from room temperature, i.e., outside the range 15-30 °C, the calibration should be performed at the desired temperature. Mussini et al.1-3 did not perform measurements of pHRVS at concentrations higher than 70 wt % MEG. This, however, does not render the approach described in this work useless at higher concentrations. Equation 6 still gives a standard reference for the pHRVS value for any MEG concentration at any temperature. Reproducible results are obtained using any electrode as long as the same reference is used. If other values of the KHPh buffer solution are recognized as standard reference values (pHRVS), we can simply adjust the reference (eq 6). Calibration Results. ∆pHMEG. Each of the electrodes used during the experimental work was calibrated and given a ∆pHMEG function. Equation 9 was developed at room temperature (25 °C) for a Mettler Toledo 3 M KCl bridge electrode of the ceramic disk type
∆pHMEG ) 0.327w - 0.233w2 + 0.564w3 pressure was therefore generally negligible. The uncertainty in the CO2 pressure determination was estimated to be (0.01 bar. Results and Discussion When the pH meter is calibrated in aqueous IUPAC standard solutions, it measures the value pHmeas in MEG-containing solutions. Different electrode systems give different values of pHmeas under otherwise identical conditions, i.e., pHmeas is generally not identical for different electrodes in water + MEG solutions. Figure 1 shows a comparison of pHmeas and the actual pH (pHRVS) of the KHPh standard solution as function of MEG concentration. The measurements were performed with different electrodes in our laboratory and at the Institute for Energy Technology (IFE, Kjeller, Norway).16 pHmeas is clearly seen to vary between the different electrodes. The procedure below aims at minimizing this variation. This procedure needs to be performed once for each electrode. Calibration Procedure. To take care of changes in MEG content, the electrode is calibrated in the following way: (1) Calibrate the electrode in standard aqueous solutions (e.g., pH ≈ 4, 7, and 10). The pH electrode now gives pHmeas. (2) Obtain pHmeas in 0.05 m KHPh solutions of varying MEG concentration, e.g., 30, 50, 70, and 90 wt % MEG. (3) Calculate the pHRVS value of the KHPh buffer solutions from eq 6. (4) Calculate ∆pHMEG ) pHRVS - pHmeas for each of the measurements performed in the MEG-containing solutions. (5) Construct a function for ∆pHMEG in a manner similar to that given by eq 9. When this calibration has been performed once for a given electrode, pH can be found from pHmeas using the equation
pH ) pHMeas + ∆pHMEG
(8)
The main simplification is that the procedure uses a singlepoint calibration in MEG-containing solutions. Thus, it utilizes the same pH-E slope as obtained in water. This simplification can lead to errors when the actual pH is far from the pH of the buffer. The problem is solved if a second solution is defined in MEG + water solutions, but at present, no such standard reference is available. Several different electrodes were tested in our laboratory, and the assumption of constant pH-E slope
(9)
where w denotes the weight fraction of monoethylene glycol in the solvent. For other electrodes or at higher temperatures, the correlation will be slightly different. For practical use, it is interesting to see how much variation different electrodes actually give. Data from glass electrodes with a KCl bridge electrolyte are shown as solid diamonds in Figure 1, whereas the triangles are for a special low-temperature glass electrode with an EtOH electrolyte. The variations between the KCl bridge electrodes were not severe. A common ∆pHMEG function was therefore constructed for these electrodes at room temperature
∆pHMEG ) 0.416w - 0.393w2 + 0.606w3
(10)
If no calibration is available for a given combined glass electrode with a KCl bridge, eq 10 provides a convenient approximation. Using this equation together with an arbitrary KCl salt bridge glass electrode should yield an accuracy of (0.1 pH units. There is obviously no guarantee that any glass electrode will comply with this function, and certainly not any other type of pH electrode, e.g., a solid-state electrode. This is clearly seen from Figure 1 where the low-temperature electrode shows a very different behavior. Thus, for types other than the KCl bridge glass electrode, a full calibration in MEG-containing KHPh standards should always be performed. Reproducibility and Accuracy. When the calibration has been performed once for a given electrode, the actual pH can be obtained from the measured value pHmeas according to the equation
pH ) pHmeas + ∆pHMEG
(11)
In this work, we propose that it is sufficient to calibrate only once in MEG solutions (to get ∆pHMEG), while the common aqueous calibration is performed regularly. We have observed this to be a good procedure for electrodes that have been in use over a period of 2 years. The best practice would be to calibrate in a KHPh buffer with the same MEG content as in the sample to be investigated, as well as at the same temperature, shortly before use. This is obviously feasible only in research laboratories and not for practical use in, for example, oilfield operation. The response of a glass electrode is slower in water + MEG solution than in water, and the electrode should be left for at
Ind. Eng. Chem. Res., Vol. 46, No. 14, 2007 4737 Table 1. Measured pH and Calculateda pHcalc in 60 wt % MEGb at Temperatures of 4-25 °C with Varying NaHCO3/KHCO3 Contentsc 4 °C, PCO2 ) 1.00 bar
25 °C, PCO2 ) 0.98 bar
25 °C, PCO2 ) 0.98 bar
NaHCO3
pH
pHcalc
NaHCO3
pH
pHcalc
KHCO3
pH
pHcalc
0 0.06 1.9 9.4 44.7 93.1 158.8 227.7 280.9 328d 337.7 422.9
4.01 4.68 5.48 6.11 6.69 6.98 7.20 7.35 7.44
4.05 4.23 5.46 6.14 6.76 7.06 7.26 7.39 7.47 7.53 7.53 7.53
0 19.9 99.4 199.0 298.9 399.5 440d 599.0
4.13 6.58 7.21 7.49 7.65 7.77
4.14 6.61 7.26 7.53 7.68 7.79 7.83 7.83
0 0.2 1.3 11.4 42.1 78.2 152.8 236.1 338.2 451.1 671.2 994.3 1120d 1127.6 1459.7
4.22 4.55 5.46 6.40 6.94 7.19 7.46 7.64 7.79 7.90 8.06 8.20
4.14 4.71 5.47 6.38 6.90 7.15 7.42 7.59 7.73 7.84 7.96 8.11 8.14 8.14 8.14
7.52 7.52
7.80
8.27 8.32
a Calculations performed using the MultiMEGScale13,18 program package. b Weight percent in solvent. c Units of mmol/kg of solvent. d Calculated solubility limit
Table 2. Measured pH and Calculateda pHcalc in 90 wt % MEGb at Temperatures of 4-25 °C with Varying NaHCO3/KHCO3 Contentsc 4 °C, PCO2 ) 1.01 bar
25°C, PCO2 ) 0.99 bar
25°C, PCO2 ) 0.99 bar
NaHCO3
pH
pHcalc
NaHCO3
pH
pHcalc
KHCO3
pH
pHcalc
0.59 1.7 12.4 50.0 100.5 214.8 320d 452.4 897.5
5.60 5.97 6.78 7.33 7.59 7.87
5.41 5.84 6.68 7.24 7.51 7.81 7.96 7.96 7.96
0 1.0 9.0 47.0 104.3 164.1 216.9 270.2 313.1 390d 378.5 514.1
4.88 5.90 6.80 7.45 7.82 7.96 8.10 8.20 8.22
4.27 5.85 6.78 7.43 7.76 7.93 8.04 8.13 8.18 8.27 8.27 8.27
0 0.2 0.8 12.3 41.3 85.8 158.5 314.9 505.2 857.3 901d 1153.5 1725.9
4.96 5.50 5.92 7.03 7.49 7.76 7.99 8.26 8.45 8.69
4.27 5.15 5.75 6.90 7.39 7.67 7.91 8.16 8.35 8.54 8.55 8.55 8.55
8.03 8.03
8.29 8.30
8.74 8.74
a Calculations performed using the MultiMEGScale13,18 program package. b Weight percent in solvent. c Units of mmol/kg of solvent. d Calculated solubility limit
least 20-30 min in the solution to obtain stable readings. High salinity and high temperature make pH measurements more difficult to reproduce, but the calibration method proposed in this work should minimize the error. During practical operation, pH should generally be reproduced within to 0.1-0.15 pH units with any type of electrode. The best results are obtained at room temperature, but reproducibility of better than 0.05 pH units ((3 mV) seems unattainable. Application and Testing. MultiMEGScale13,18 is a computer model designed to predict mineral precipitation from water + MEG solutions. The model is based on the idea that the activity of a species, ai, can be expressed as
ai ) miγiSγiN
(12)
where m denotes the concentration (moles per kilogram of solvent) and γiS is the activity coefficient calculated from the Pitzer model17 as if the solVent were water, meaning that γiS has a numerical value independent of MEG concentration. The term γiN is introduced to compensate for the error incurred when incorrectly using the Pitzer model in a water + MEG solution. For each individual equilibrium reaction, the γiN term was empirically fitted to experimental data. This is virtually the same modeling approach as used by Kan et al.,4 the difference being that Kan et al.4 used units of moles per kilogram of water as the concentration measure and when calculating γiS. When pHdependent equilibria were investigated to create the MultiMEG-
Scale13,18 model, the electrodes were calibrated as presented in this work. An important feature of this scale model is calculations of carbonic acid equilibria (eqs 13-15) in water + MEG solutions.
CO2(g) ) CO2(aq)
(13)
CO2(aq) + H2O ) H + + HCO3-
(14)
HCO3- ) H + + CO32-
(15)
To verify that the MultiMEGScale model worked for two given salt systems, pH was measured in MEG + water solutions of known NaHCO3 and KHCO3 contents with continuous CO2 bubbling of the solutions (see Tables 1-3). These measurements were performed with different electrodes than used for construction of the model.13,18 Such solutions with high concentrations of bicarbonate are used to reduce corrosion in pipelines. The method is called pH stabilization.19 Figure 2 shows measured pH at seabed temperature (Tables 1 and 2) in 60 and 90 wt % MEG versus NaHCO3 concentration on a logarithmic scale. First, the measurements closely resemble the model for concentrations ranging from 1 mmol/kg to the solubility limit of sodium bicarbonate. Second, it is noted that the breakpoints in the pH curves coincide with the calculated solubility limit. Experiments were performed with NaHCO3 at 4, 25, 50, and 80 °C and with KHCO3 at 25 °C. The MEG
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Table 3. Measured pH and Calculateda pHcalc at Temperatures of 50 and 80 °C in 60 or 90 wt % MEGb with Varying NaHCO3 Contentsc 50°C, PCO2 ) 0.92 bar, 60 wt %
80°C, PCO2 ) 0.72 bar, 60 wt %
50°C, PCO2 ) 0.94 bar, 90 wt %
NaHCO3
pH
pHcalc
NaHCO3
pH
pHcalc
NaHCO3
pH
pHcalc
NaHCO3
pH
pHcalc
0 0.23 1.2 9.1 58.5 112.9 192.6 267.5 321.5 377.2 434.2 503.0 556.3 600d 609.4 672.4
4.39 4.5 5.62 6.54 7.31 7.56 7.79 7.94 7.98 8.04 8.1 8.15 8.18
4.27 5.00 5.72 6.54 7.29 7.54 7.75 7.87 7.94 8.00 8.05 8.10 8.14 8.19 8.19 8.19
0 3.8 12.8 60.1 114.3 273.2 394.9 584.2 751.9 930d 1005.5 1368.9
4.76 6.69 7.21 7.81 8.04 8.37 8.47 8.58 8.62
4.50 6.66 7.16 7.77 8.02 8.34 8.46 8.60 8.67 8.70 8.70 8.70
0 0.17 1.3 9.0 65.2 165.4 265.2 432.5 498d 644.2 964.4
5.10 5.56 6.32 7.12 7.9 8.27 8.5 8.63
4.33 5.39 6.26 7.07 7.86 8.22 8.40 8.58 8.64 8.64 8.64
0 0.96 8.8 46.1 92.9 180.6 297.2 452.2 651d 666.8 925.1
5.40 6.59 7.53 8.16 8.34 8.59 8.74 8.88
4.55 6.53 7.45 8.12 8.39 8.64 8.82 8.97 9.10 9.10 9.10
8.22 8.23
8.63 8.64
8.65 8.65
80°C, PCO2 ) 0.88, 90 wt %
8.95 8.95
a Calculations performed using the MultiMEGScale13,18 program package. b Weight percent in solvent. c Units of mmol/kg of solvent. d Calculated solubility limit.
calibrated in common aqueous standard solutions yields pHmeas. pHmeas is not reproducible by different electrode systems in water + MEG solutions. The actual pH can be found from pHmeas using the equation
pH ) pHmeas + ∆pHMEG ∆pHMEG is determined by calibration in MEG standards (0.05 mol KHPh/kg of solvent). This calibration has to be performed once for each electrode. For ordinary combined glass electrodes of the KCl bridge type, there is little variation between individual electrodes. Thus, for such electrodes, a convenient approximation (at room temperature) for ∆pHMEG is given by
∆pHMEG ) 0.416w - 0.393w2 + 0.606w3
Figure 2. pH in (]) 60 and (2) 90 wt % MEG + water solutions at a temperature of 4 °C as a function of NaHCO3 content (mmol/kg of solvent). Solutions saturated with CO2 at ∼1 atm total pressure Solid lines, calculated pH; dotted lines, calculated solubility limit. Calculations performed using the MultiMEGScale13,18 program package.
where w denotes the weight fraction of MEG in the water + MEG solvent. This ∆pHMEG function gives a slightly decreased accuracy, but enables a connection between pHmeas and pH when treating old data, i.e., when it is impossible to obtain a calibration of the actual electrode. Supporting Information Available: We thank Statoil and Norsk Hydro for financial support
concentrations used were 60 and 90 wt %, which are typical “lean” and “rich” MEG conditions, respectively, in oilfield applications. All results (see Tables 1-3) were generally within 0.1 pH units of the model. It can be concluded that the pH calibration method is well-suited for work in such saline solutions with high MEG contents. A thorough discussion of the results in Tables 1-3 can be found in Sandengen.13 If calibration values in MEG standard solutions are not available, it is possible to use eq 10. This is a less accurate approach but might be the only possible option in some cases, e.g., if old experimental data are to be treated. With such a common ∆pHMEG term, it is consequently possible to compare a pHmeas value with the output pH from the mentioned scale model.13 Conclusions For pH measurements in mixed-solvent systems, the pH electrode should be calibrated in common IUPAC aqueous standards as well as in mixed-solvent standards. An electrode
Literature Cited (1) Mussini, P. R.; Marcolungo, I.; Rondinini, S.; Longhi, P. Acidbase equilibria and acidity scales in ethylene glycol/water solvent mixtures: Recommended reference-value pH-metric standards and ionization constants for o-phthalic acid at normal and subzero temperatures. Chim. Ind. (Milan) 1991, 73, 262-268. (2) Mussini, P. R.; Mussini, T.; Rondinini, S. Thermodynamics of HCl and recommendations for the standard emf of the cell hydrogen/silver chloride in ethylene glycol/water mixtures from -40 to 50°C. Chim. Ind. (Milan) 1991, 73, 190-194. (3) Mussini, P. R.; Mussini, T.; Rondinini, S. Reference value standards for pH measurements in D2O and aqueous-organic solvent mixtures: New accessions and assessments. Pure Appl. Chem. 1997, 69 (5), 1007-1014. (4) Kan, A. T.; Fu, G.; Tomson, M. B. Effect of Methanol on Carbonate Equilibrium and Calcite Solubility in a Gas/Methanol/Water/Salt Mixed System. Langmuir 2002, 18 (25), 9713-9725. (5) Mussini, T.; Mazza, F. Medium effects as key electrochemical variables for corrosion and electrochemistry studies in non-aqueous and mixed solvents. Electrochim. Acta 1987, 32 (6), 855-862. (6) Mussini, T.; Covington, A. K.; Longhi, P.; Rondinini, S. Criteria for Standardization of pH Measurements in Organic Solvents and Water + Organic Solvent. Pure Appl. Chem. 1985, 57 (6), 865-876.
Ind. Eng. Chem. Res., Vol. 46, No. 14, 2007 4739 (7) Bates, R. G. Determination of pH. Theory and Practice, 2nd ed.; Wiley: New York, 1973. (8) Bates, R. G. The modern meaning of pH. Crit. ReV. Anal. Chem. 1981, 10 (3), 247-278. (9) Bates, R. G.; Paabo, M.; Robinson, R. A. Interpretation of pH measurements in alcohol-water solvents. J. Phys. Chem. 1963, 67, 18331838. (10) Covington, A. K.; Bates, R. G.; Durst, R. A. Definition of pH scales, standard reference values, measurement of pH and related terminology. Pure Appl. Chem. 1985, 57, 531-542. (11) Buck, R. P.; Rondinini, S.; Covington, A. K.; Baucke, F. G. K.; Brett, C. M. A.; Camoes, M. F.; Milton, M. J. T.; Mussini, T.; Naumann, R.; Pratt, K. W.; Spitzer, P.; Wilson, G. S. Measurement of pH. Definition, Standards, and Procedures. Pure Appl. Chem. 2002, 74 (11), 2169-2200. (12) Bagg, J. Computer calculation of liquid-junction potentialssI. Concentration profiles at junctions with concentrated potassium chloride. Electrochim. Acta 1990, 35 (2), 367. (13) Sandengen, K. Prediction of mineral scale formation in wet gas condensate pipelines and in MEG (monoethylene glycol) regeneration plants. Ph.D. Dissertation, Norwegian University of Science and Technology, Trondheim, Norway, 2006. (14) Trimble, H. M.; Potts, W. Glycol-water mixtures. Vapor pressureboiling point-composition relations. Ind. Eng. Chem. 1935, 27, 66-68.
(15) Villamanan, M. A.; Gonzalez, C.; Van Ness, H. C. Excess thermodynamic properties for water/ethylene glycol. J. Chem. Eng. Data 1984, 29 (4), 427-429. (16) Seiersten, M. Institute for Energy Technology (IFE), Kjeller, Norway. Personal communication, 2005 (17) Pitzer, K. S. Thermodynamics, 3rd ed; McGraw-Hill: New York, 1995. (18) Sandengen, K.; Kaasa, B.; Østvold, T. Prediction of Scale Potential in Ethylene Glycol Containing Solutions. Presented at the 17th International Oil Field Chemistry Symposium, Geilo, Norway, March 19-22, 2006. (19) Seiersten, M.; Dugstad, A.; Gulbrandsen, E. Conditions for scaling in pipelinesspH in glycol solutions. Presented at the 4th International Symposium on Oilfield Scale, Aberdeen, U.K., Jan 28-30, 2003; Paper SPE 80393
ReceiVed for reView October 12, 2006 ReVised manuscript receiVed April 18, 2007 Accepted May 8, 2007 IE061305A
