Environ. Sci. Technol. 1999, 33, 2607-2610
Role of Temperature and pH in Cu(OH)2 Solubility LOAY HIDMI AND MARC EDWARDS* Department of Civil Engineering, Virginia Polytechnic Institute, 407 NEB, Blacksburg, Virginia 24061-0246
The wide variation in reported cupric hydroxide solubility constants is attributed to the age of the solid, kinetic limitations in approaching equilibrium, and transient formation of basic copper nitrate solids. In situations similar to those where cupric hydroxide solubility constants were determined via precipitation in dilute nitrate solutions, copper nitrate solids (gerhadtite) form first followed by a rapid transition to cupric hydroxide and then a slower transition to tenorite. In solutions containing hydroxide and nitrate as the sole anionic constituents, fresh cupric hydroxide solids control soluble copper concentrations for minutes to days depending on pH and temperature. Thus, although a log K of 9.36 ( 0.02 was most consistent with the experimental data for Cu(OH)2 solubility, decreasing log K values from 9.36 [fresh Cu(OH)2] to 7.6 [CuO] can be expected with aging as controlled by temperature, pH, and other factors.
Introduction Understanding freshly precipitated copper hydroxide solubility has long been of key importance in the treatment of copper-containing wastewater (1, 2). Recently, copper hydroxide solubility has been implicated in controlling corrosion byproduct release to drinking water from relatively new copper plumbing (3); thus, the solubility product is of direct relevance to compliance with copper regulations in drinking water and in assessing human exposure. Corrosion of plumbing is also a key contributor to copper loading of wastewater treatment plants and the environment (4). Even under well-defined conditions, reported solubility products (5-21) for freshly precipitated Cu(OH)2 vary by nearly 3 orders of magnitude (Table 1). Of these constants, that determined by Schindler et al. (1965) is thought to be “the most reliable of those reported but it is classified as tentative because of the instability of the Cu(OH)2-alkali system and because there is no other work to substantiate it (7)”. The wide variation in reported constants has been attributed to surface area effects (6, 18) and/or aging (1). Our review indicates that reported solubility products can also be demarcated into groups of low (9.6) solubility, corresponding to experiments in which equilibrium was approached through dissolution or precipitation reactions, respectively (Table 1). This differentiation may also be important since equilibration times and experimental procedures are roughly similar within each type of experiment. That is, equilibration times were apparently always greater than 2.5 days when dissolution experiments were conducted, whereas a few hours or less is commonplace in precipitation experiments. At least three phenomena could produce the wide variation in copper solubility reported for the precipitation * Corresponding author; fax: 540 231-7916; e-mail: edwardsm@ vt.edu. 10.1021/es981121q CCC: $18.00 Published on Web 06/26/1999
1999 American Chemical Society
versus dissolution experiments. If differing solid age is dominant in controlling cupric hydroxide solubility, the shorter term precipitation experiments would produce solids with higher solubility (1, 18). Likewise, if a kinetic limitation prevented equilibrium from being achieved with a particular solid phase, an erroneously high solubility product would be estimated for the case of precipitation and a lower solubility would be expected for dissolution. Interestingly, only few researchers (12) have directly addressed the approach to equilibrium through both dissolution and precipitation reactions. Finally, since the type of solid that forms depends on the anion present, and chemical order of addition has been noted to influence results in the copper hydroxide system (1), unsuspected formation of basic copper nitrate solids could influence solubility under some circumstances. This work examines each of these possibilities by expanding the experimental scope and resulting interpretations of previous researchers (1). The pH range from 7 to 9, not directly examined in that work but of obvious importance in engineered systems, is a special focus of this investigation. Moreover, the practical impacts of temperature on Cu(OH)2 solubility and aging are established.
Materials and Methods All chemicals used in these experiments were analytical reagent grade. Solutions were prepared from deionized type III reagent grade water (Milli-Q water reagent system) filtered through a 0.2 µm pore size filter before use. Batch Tests. After preliminary experiments at pH 9 demonstrated identical results using either Teflon or glass beakers, 2 L glass beakers were used in all experiments. For precipitation experiments, 2 L of solution containing 0.5 mM of Cu(NO3)2 and 1 mM NaNO3 was titrated with 50 µL increments of 1 M freshly prepared NaOH dosed at 1 min intervals. Throughout the experiment, the system was mixed with a 5 cm magnetic stirrer at 200 rpm and monitored for time periods of up to 1 month. The final ionic strength was 2 mM as NaNO3 after all copper was precipitated. The pH was monitored using a Beckman Φ10 pH meter and a Corning combination electrode. During experiments over the course of a single day, the recalibration was every 2 h, whereas the meter was recalibrated immediately before measurements at longer time periods. The pH was maintained to within (0.15 pH units of the target value through addition of 1 N NaOH throughout the entire experiment. Temperature was measured using a digital thermometer and maintained to within (0.3 °C of the target value. CO2 was initially stripped from the water by bubbling the solution with N2, and during the remainder of the experiment, beakers were covered to the extent possible with a double layer of Parafilm. The total initial carbonate was 5 × 10-6 M in 0.1 M of NaNO3 precipitated by adding dilute NaOH. The value of log K of 10.3 was obtained after equilibration for 3-5 min. However, when that period was changed to 120 min the obtained value for log K was 10.0.
copper obtained after temperature-controlled centrifugation at 4000 rpm for 20 min. For the comparison, a 2 L solution with an initial copper concentration of 1 mM Cu(NO3)2 was adjusted to pH 6.0, 7.0., 8.0, and 9.0. Free copper was also determined in each solution using an Orion cupric ion specific electrode and a single junction reference electrode (1). The electrode was calibrated immediately before use with standards in the concentration range under consideration. Using this information, soluble copper could then be calculated with the electrode based on the measured free copper concentration, soluble pH, and reported cupric hydroxide complexation constants (20). Soluble copper in filtered or centrifuged samples was determined directly using a Varian Liberty axial inductive coupled plasma emission spectroscopy (ICP-ES) instrument according to standard method 3120 with a detection limit of (1 µg/L. At a pH of 8 or less, the difference in calculating free copper using the ISE and the ICP approach was always less than 9%, regardless of whether centrifugation or filtration was used for solid/liquid separation. For routine experiments, 0.45 µm filters were used for solid/liquid separation and the ICP-ES was used to quantify soluble copper. The exception was at pH 8.0-9.5 and at total soluble copper concentrations of e20 µg/L, in which case more than 20% of the copper was unavoidably lost to the filter. To obtain an accurate estimate of soluble copper at this higher pH range, a combination of centrifugation (providing an upper bound to actual solubility) and filtration (providing a lower bound to actual solubility) measurements was necessary. All solubility constants [K ) [Cu2+]/[H+]2] reported in this work are corrected to 298.1 K using reported enthalpy values (20) and to zero ionic strength using the extended DebyeHuckel approximation (22). Solid Characterization. Solids collected on the 0.45 µm pore size filter were rinsed with MQ reagent grade water three times and then dried at room temperature in a desiccator. These solids were then inspected by scanning electron microscopy (ISI-SX 30 KV), X-ray diffraction, and wet chemical analysis after sulfuric acid digestion.
Results and Discussion Initial experiments focused on the kinetics of equilibration for dissolution versus precipitation reactions. Thereafter, the effect of pH and temperature on solubility was examined followed by a determination of solids composition and morphology. 2608
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FIGURE 1. Effect of time on the soluble copper concentration as measured through dissolution and precipitation of Cu(OH)2 solids. Experimental conditions: pH 7.0, 23.4 °C, I ) 2 mM NaNO3 and 0.5 mM Cu2+total. Kinetics of Equilibration: Precipitation vs Dissolution. If the kinetic approach to solid/liquid equilibrium is insignificant when compared to the experimental timespan, the solubility product determined through precipitation or dissolution experiments should not differ. For precipitation of cupric ion at pH 7.0, the soluble copper concentration decreased to half its initial value in just 3 min, and then steadily decreased to a stable level of 6.2 ( 0.5 mg/L Cu between 2 and 23 h (Figure 1). During the period of relative stability, it was not necessary to add base to maintain a constant pH. After about 24 h, the soluble copper concentration began to decrease, corresponding to a time period in which base was once again required to maintain constant pH at 7.0. Between 24 and 720 h, the color of the solid gradually changed from blue to dark brown. For the dissolution experiment, a subsample of the solid formed during the precipitation experiment was collected by filtration. This solid was then immediately resuspended in a solution of identical pH and ionic strength but without copper. Soluble copper concentrations increased steadily during the first 6 h before plateauing at the same 6.2 ( 0.2 mg/L Cu level observed during the concurrent precipitation experiment (Figure 1). Thereafter, trends in solubility for the precipitation and dissolution experiments were nearly identical, although the concentration of copper in the precipitation experiment was slightly lower from that in the dissolution experiment between 30 and 92 h. After 720 h, soluble copper decreased to a relatively stable value of 0.054 mg/L in the precipitation experiment and 0.064 mg/L in the dissolution
TABLE 2. Particle Characterization for Solids Aged for 30 Minutes and 96 Hours pH 7.0
pH 8.0 pH 9.0
aging period
SEM
XRD
30 min 5h
no visible crystals tapelike and oval shaped solids.
gerhardtite (Cu2NO3(OH)3) gerhardtite (Cu2NO3(OH)3), likasite (Cu3NO3(OH)5‚2H2O), cupric hydroxide (Cu(OH)2) nitrate solids < 20% of copper solids likasite and cupric hydroxide nitrate solids < 15% of copper solids tenorite (CuO) gerhardtite and tenorite (CuO) tenorite (CuO) tenorite (CuO) gerhardtite, likasite, and tenorite tenorite (CuO) tenorite (CuO)
96 h 720 h 30 min 96 h 720 h 30 min 96 h 96 h
oval and slab shaped solids large oval shape small oval shape large oval shape large oval shape small oval shape large oval shape large oval shape
FIGURE 2. Effect of pH on cupric ion precipitation in 2 mM NaNO3 solutions at 23 °C. experiment. After considering the significant Cu(OH)+, Cu(OH)20, and Cu2(OH)2-2 complexes (20), the log K of the metastable equilibrium period between 6 and 23 h was 9.36 ( 0.02, whereas after 1 month of aging, log K was 7.35 ( 0.05. The calculated values of log K are highly dependent on the selected value of B2 for Cu(OH)2 formation, for which there is significant variation in the literature. If the lower range of B2 values (i.e., log B2 ) -16.24) is deemed more accurate (1), the above values for log K would increase to 9.69 and 7.75 for data at 6-23 h and 1 month of aging, respectively. Since the soluble copper concentrations obtained through precipitation and dissolution reactions were the same only after 6 h, a significant portion of the previously cited 2 order of magnitude differences in log K for cupric hydroxide solids may be attributed to kinetics of precipitation and dissolution. For example, if solubility had been determined after only 25 min in the precipitation and dissolution experiments for the exact same solid, log K’s of 9.6 and 8.45 would have been determined, respectively. Consistent with the observations of previous researchers (1, 6, 18), solid aging and/or surface area effects also seem dominant within the two types of experiments. That is, a metastable equilibrium condition is obtained for solids aged a few hours, which roughly corresponds to the log K reported for precipitation reactions, whereas another metastable equilibrium condition is achieved after days or weeks of aging which corresponds to solubility products that have been reported for dissolution reactions (Table 1). At intermediate timespans, log K’s between these extremes were valid. Solubility and pH. The precipitation experiment was repeated at a range of pHs (Figure 2). At pH 7.0, soluble copper concentrations at equilibrium were 6.41 mg/L, a level which was stable for 24 h. Thereafter, the soluble copper concentration started to decrease gradually until reaching a value of 0.056 mg/L after 1 month of aging. Interestingly, comparing results at pH 7.0 but repeated at different days (Figure 1 vs Figure 2), the duration of pseudo-equilibrium changed slightly but the average level of soluble copper was
the same at 6.3 ( 0.2 mg/L. This suggests that the period of metastability is dependent on very slight differences in chemical addition rate and other subtle factors not precisely controlled in experiments, a phenomena consistent with the instability of the system under consideration. Trends at pH 7.5 were similar to those at pH 7.0 (Figure 1), but the metastable copper level of 6.11 ( 0.2 mg/L persisted for just 4 h instead of the 24 h observed at pH 7.0. A somewhat different trend was observed for pHs higher than 8.0 in that an initial metastable period was not observed. Instead, within the first hour, the color of the solid immediately changed from blue to dark brown. As discussed before, samples of this solution were centrifuged and filtered to determine an upper and lower bound to the actual soluble copper concentration. Initially, the difference in the measured copper concentration between the two methods was approximately 0.4 mg/L for all tested pHs; however, that difference decreased with time. For example, at pH 8.0 and after 24 h of aging, soluble copper was 0.12 mg/L and 0.043 mg/L as determined by filtration and centrifugation, respectively. Thus, the solubility product log K was between 7.9 and 7.4. After 1 month, however, soluble copper levels were 0.043 ( 0.005 mg/L for both approaches. Solid Phases. Precipitated solids were collected after 30 min and 96 h of reaction time at pHs 7.0, 8.0, and 9.0. Another sample of solids aged at pH 7.0 for 5 h was collected to better characterize the solids present during the metastable equilibrium. In all cases, nitrate-containing copper solids (likasite or gerhardtite) were identified by XRD after 30 min reaction time (Table 2). After 5-96 h, crystalline cupric hydroxides or tenorite were also present, either exclusively or in addition to the nitrates. To better define the concentration of copper nitrate solids present, solids were dissolved in sulfuric acid and nitrate was determined by ion chromatography. After 5 h, copper nitrate solids always accounted for less than 20% of the total copper solids present, indicating that copper hydroxides were dominant. The relative concentration of copper nitrates tended to decrease with time, although they were always present through 96 h of aging. After 1 month of aging and at all pHs tenorite was the dominant solid and there were no traces of copper nitrate solids. Thus, nitrate-containing solids seem to be of low significance after 5 h aging time. On the other hand, at shorter time periods, the results strongly support the presence of basic cupric nitrate solids. Though the XRD data on likasite or gerhardtite might result from reaction of sorbed NO3- on amorphous Cu(OH)2 as the solids dried, this seems highly unlikely given the dilute solutions employed, the solid-rinsing procedure used, and the complete absence of these phases at later experimental times. Although the copper nitrate solids represented a relatively small fraction of the total copper solids present, the question remained as to whether these solids might be responsible for the initial metastable equilibrium period. To address this question, a pure copper electrode was immersed in a solution VOL. 33, NO. 15, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Effect of temperature on the precipitation of cupric hydroxide solids. I ) 2 mM NaNO3 and 0.5 mM Cu2+total. of deionized water open to atmospheric oxygen and galvanically oxidized at pH 5 according to procedures presented elsewhere (22). The stable reaction product was freshly precipitated cupric hydroxide, as predicted thermodynamically and measured by XRD. When this freshly precipitated solid was collected and placed into a container of deionized water maintained at pH 7.0 by NaOH addition, soluble copper concentrations leveled off at 5.76 mg/L after 4 h. After correcting for ionic strength effects, this translates to a log K of 9.35 for cupric hydroxide compared to the log K of 9.36 ( 0.02 determined for the mixed solids of age greater than 6 h in the presence of nitrate. Thus, at time periods greater than 6 h in the original experiment, solubility was controlled by Cu(OH)2 and was not influenced by nitrate. Temperature Effects. The role of temperature was examined in copper precipitation experiments at pH 7.0 using filtration to determine soluble copper. The general trends for changes in soluble copper with time were similar to those noticed at room temperature. However, differences in the soluble copper concentration and duration of the metastable equilibrium condition were noted (Figure 3) at 4.2, 12.5, and 23.6 °C. At 4.2 °C, the precipitate was always a light-blue color and a metastable copper concentration of 22.1 ( 0.4 mg/L was maintained for about 72 h. Thereafter, copper solubility rapidly decreased to 1.51 ( 0.1 mg/L after 1 month. At 12.5 °C, the initial duration of metastable equilibration was about the same as that at 4.2 °C, although the soluble copper concentration was 15 mg/L ( 0.15, eventually declining to 1.11 ( 0.12 mg/L after 1 month. The solids remained blue during the entire experiment. At 23.6 °C, the metastable equilibrium persisted for much less than 72 h as was observed previously. Assuming that the difference in copper levels during metastable equilibrium can be attributed to enthalpy considerations over the temperature range 5-35 °C using the Van’t Hoff relationship (23), ∆H for the Cu(OH)2 precipitation reaction was estimated as 14.5 kcal/mol using the data at 12.5 and 23.6 °C. This value is close to the 13.19 kcal/mol reported elsewhere for Cu(OH)2 solids in the literature (20). Final Comment on Sequence of Solid Formation. Recent studies of cupric ion precipitation by base addition in dilute cupric nitrate solutions have led to divergent conclusions regarding the likely sequence of solid formation. In one case (24), initial titration data were consistent with 2 mol of OHconsumption per mole of Cu2+ precipitated, followed by the appearance of a Cu(OH)1.5(NO3)0.5 solid as confirmed by X-ray
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diffraction and stoichiometry (if nitrate concentrations were above 0.01 M). The authors hypothesized that an amorphous cupric hydroxide solid phase formed first, followed by a gradual transition to basic cupric nitrate solids, although we note that the data are also qualitatively consistent with soluble Cu(OH)20 complex formation followed by precipitation of basic cupric nitrate solid. In contrast, interpretation of titration data by other investigators in 1 mM nitrate solutions suggested that a transient basic cupric nitrate phase might form first, followed by rapid conversion to cupric hydroxide (1). The X-ray and wet chemical analysis reported in this work are obviously more consistent with the latter hypothesis, although subtle differences in rates of base addition and nitrate concentrations could also partly explain differences in the results obtained.
Acknowledgments This work was supported by the National Science Foundation under Grant BES-9729008 and the American Water Works Association Research Foundation (AWWARF). The opinions, findings, conclusions, or recommendations are those of the authors and do not necessarily reflect the views of AWWARF or NSF.
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Received for review November 2, 1998. Revised manuscript received May 12, 1999. Accepted May 21, 1999. ES981121Q
