Calorimetric study of the digestion of gibbsite, A1(OH)3(cr), and thermodynamics of aqueous aluminate ion, A1(OH)4-(aq)

Transcription

1 Calorimetric study of the digestion of gibbsite, A1(OH)3(cr), and thermodynamics of aqueous aluminate ion, A1(OH)4-(aq) QIYUAN CHEN,' YUMING XU,~ AND LOREN G. HEPLER Department of Chemistry and Department of Chemical Engineering, University of Alberta, Edmonton, Alta., Canada T6G 2G2 Received January 29, 1991 This paper is dedicated to Dr. Richard Norman Jones QIYUAN CHEN, YUMING XU, and LOREN G. HEPLER. Can. J. Chem. 69, 1685 (1991). We have made calorimetric measurements of the enthalpies of solution of gibbsite, Al(OH)3(cr), in aqueous sodium hydroxide solutions at five temperatures from 100 to 150 C. Results of these measurements have been used to obtain the standard enthalpies of formation of Na+(aq) + A1(OH)4-(aq) at the experimental temperatures. These results have also led to values of AC,' for the reaction represented concisely by A1(OH)3(cr) + OHP(aq) = A1(OH)4-(aq), from which we have obtained standard state partial molar heat capacities of Na+(aq) + A1(OH)4-(aq). Combination of our results with those from earlier investigations has permitted calculation of thermodynamic properties of Na+(aq) + Al(OH)4-(aq) over a wide range of temperature and thence some generalizations about the usefulness of various equations for representing or predicting these thermodynamic properties. Key words: gibbsite, enthalpy of solution; sodium aluminate (aqueous), thermodynamic properties; heat capacities, Na+(aq) + A1(OH)4-(aq). QIYUAN CHEN, YUMING XU et LOREN G. HEPLER. Can. J. Chem. 69, 1685 (1991). Nous avons mesurt, par calorimttrie, les enthalpies de solution de la gibbsite, Al(OH)3(cr), dans des solutions aqueuses d'hydroxyde de sodium h 5 temp6ratures difftrentes allant de 100 h 150 C. Ces rtsultats ont permis d'obtenir les enthalpies standards de formation des ions Na+(aq) + A1(OH)4-(aq) h ces temp6ratures. Ces rtsultats ont Cgalement permis d'obtenir les valeurs de ACpO de la rtaction reprksentte de faqon concise par l'tquation : A1(OH)3(cr) + OH-(aq) = A1(OH)4-(aq), h partir de laquelle nous avons obtenu les capacitts molaires partielles de l'ttat standard des ions Na+(aq) + A1(OHkP(aq). En combinant nos rtsultats h ceux provenant d'ttudes prtctdentes, nous avons pu calculer les proprittts thermodynamiques des ions Na+(aq) + A1(OH)4-(aq) dans un large intervalle de temptratures et de faire une gtnhalisation relative h l'utilitt de diverses equations pour reprtsenter ces propri~t~s thermodynamiques. Mots clb : gibbsite, enthalpies de formation; NaA1(OH)4(aq), les proprittts thermodynamiques; capacitts calorifiques, Na+(aq) + A1(OHkP(aq). [Traduit par la rtdaction] Introduction The present calorimetric investigations of the enthalpy of solution of the mineral gibbsite, Al(OH)3, in aqueous sodium hydroxide (several concentrations and several-temperatures) have been carried out for two reasons. One specific reason has been to obtain thermochemical data that can be useful in connection with the Bayer process for treating gibbsite prior to processes leading directly to the production of metallic aluminum. A more general reason for the research described here has been to obtain thermodynamic data for the aqueous aluminate ion, Al(OH)4-, over a range of temperatures; these thermodynamic data can be useful in fields as varied as the geochemistry of aluminum-containing minerals and the corrosion of metallic aluminum. Results from these calorimetric measurements also permit testing earlier predictions of the temperature dependence of thermodynamic properties of A1(OH)4-(aq). We already have some thermodynamic data for Al(OH)3- (gibbsite, cr), Al(OH)4-(aq), ~l~+(aq), and related species such as OHP(aq) for 298 K. There are also some equilibrium data for higher temperatures. Most directly related to the results of 'Visiting Professor from Department of Chemistry, Central-South University of Technology, Changsha, Hunan , People's Republic of China. 2Present address: Alberta Research Council, Edmonton, Alta., Canada. this research are calorimetrically determined partial molar heat capacities for Na+(aq) + Al(OH)4-(aq) over the temperature ranges C (1) and C (2). In this work we have made calorimetric measurements of the enthalpy (AH) of the reaction represented by in which the non-reactive Na+(aq) is not shown. Experimental Synthetic gibbsite was prepared according to a procedure established by Yin et al. (3). The procedure of Yin et al. involved dissolving certified (Fisher, Lot ) aluminum hydroxide in concentrated aqueous sodium hydroxide to obtain a solution with molar ratio NaIAl = 1.7. This solution was diluted with distilled water until the concentration of NaOH was 3.5 mom. Then this solution was seeded with small crystals of gibbsite (confirmed by X-ray analysis) that were kindly supplied by Yin. The seeded solution, initially at 32 C was stirred slowly for 72 h while the temperature declined to 24 C. The resulting precipitate of gibbsite was washed with dilute sodium hydroxide and then with distilled water for several hours after the wash water became neutral. The washed Al(OH)3 was dried at 60 C and then at 105 C to constant weight. Finally, sieving led to a sample of synthetic gibbsite that was mesh to be used for analysis and calorimetric measurements. Heating to drive off all water led to f 0.05% A1203, as compared with 65.36% calculated from atomic masses.

2 1686 CAN. J. CHEM. VOL. 69, 1991 TABLE 1. Enthalpies of solution of gibbsite in aqueous sodium hydroxide (eq. [I]) NaAl(OH), in NaAl(OH)4 in final solution AH final solution AH (mol kg- ') (kl mo1-i) (mol kg-') (kj mo1- ') Ionic strength = 3.OO rnol kg-' Ionic strength = 1.00 rnol kg-' TABLE 1 (concluded) NaAl(OH), in NaAl(OH)4 in final solution AH final solution AH (mol kg-') (kj mo1-i) (mol kg- ) (kj mo1-i) Sodium hydroxide solutions were prepared from Assured Reagent (BDH, Lot ) NaOH and distilled water, with care being taken to exclude C02. Weight titrations with standard potassium hydrogen phthalate led to the temperature-independent molalities (mollkg) of the NaOH solutions. Separate titrations with hydrochloric acid solutions permitted establishment of the upper limit of carbonate contamination as 0.2% of the NaOH concentration. A Tian-Calvet heat-flow calorimeter (Setaram, Model C-80) was used for our calorimetric measurements. Calibration and testing of this calorimeter has been described in detail by Xu (4). On the basis of Xu's testing and our own experience, we estimate that accuracy of the AH values reported here is about 0.5%. Times required for complete reaction of gibbsite with excess NaOH ranged from about 45 to 75 min, depending on the temperature and the concentration of the NaOH solution. To minimize thermal effects due to corrosion of the cells, four preliminary calorimetric measurements were carried out with concentrated NaOH (5 molal) at 150 C. These preliminary experiments led to formation of a "tight" layer of corrosion product on the cell walls and effectively stopped further corrosion. Blank runs with NaOH solutions (no gibbsite) confirmed the absence of a further thermal effect due to corrosion. Results Results (AH values) of all the calorimetric measurements of heats of solution of gibbsite in aqueous sodium hydroxide are reported in Table 1. Because the reaction, as represented concisely by equation [ 1 ], is electrically symmetrical, the initial molality of NaOH is equal to the final molal ionic strength of the solution as listed in Table 1. It is possible to obtain AH0 and c,,' values corresponding to infinite dilution as follows. First we recognize that AH (at any one temperature) of an electrically symmetrical reaction is independent of ionic strength in the Debye-Hiickel range of concentration, and that initial dependence on ionic strength will be proportional to the ionic strength and also related to the composition of the constant ionic strength solution. The AH values in Table 1 are clearly dependent on temperature (AC, # 0), so we express the enthalpies of reaction as a quadratic function of temperature. All of this is summarized by the equation

3 CHEN ET AL Then we have and TABLE 2. Coefficients of eqs. [2]-[4] TABLE 3. AH' and Ac: of reaction [l] Values of the parameters in eqs. [2]-[4] are listed in Table 2. We have used the values of the parameters listed in Table 2 to calculate the values of AH corresponding to each calorimetric measurement. The standard deviation, S = [(yi- y)2/(n - 1)1'/~, associated with these calculated values is 0.14 kj mol-'. Derived values of AH0 and AC,' for the reaction represented by eq. [ ] are listed in Table 3. Thermodynamic calculations and discussion A principal aim of this research has been to obtain thermodynamic properties of aqueous sodium aluminate: Na+(aq) + A1(OH)4-(aq). To obtain these thermodynamic properties, we combine our experimental results with the already known thermodynamic properties of gibbsite and aqueous sodium hydroxide as follows. Hemingway et al. (5) have measured the heat capacities of gibbsite from 13 to 480 K. Their results lead to So = J K-' mol-i for Al(OH)3(cr,gibbsite) at T = K We have combined our enthalpies of reaction (Table 3) of gibbsite with NaOH(aq) with the enthalpies of formation of gibbsite (Table 4) and aqueous NaOH (Table 5) to obtain the enthalpies of formation of aqueous sodium aluminate, Na+(aq) + Al(OH)4-(aq), listed in the second column in Table 6. The values of Af(HoT - H ~ ~ listed ~ in ~ the, third ~ ~ column ) in Table 6 are calculated from heat capacities of the elements (7) and the partial molar heat capacities of NaAl(OH)4(aq) from this work as discussed in the following paragraphs. Finally, the fourth column in Table 6 lists our values for the standard state enthalpy of formation of aqueous sodium aluminate at T = K: the average is AfHO = f 0.24 kj mol-i, where the f indicates the standard deviation. The uncertainty in this AfHO is much larger than the standard deviation. Considerations of uncertainties in auxiliary data (other standard enthalpies of formation and heat capacities) and our measurements lead us to estimate that the total uncertainty in our AfHO values in Table 6 is about f 1.5 kj mol-'. To obtain values for the standard state partial molar heat capacities of aqueous sodium aluminate, we start with the bcp0 values listed in Table 3 for the reaction represented by eq. [I]. Combinations of these ACP0 values with heat capacities at the same temperatures for gibbsite (Table 4) and aqueous NaOH (Table 5) lead to the standard standard state heat capacities of NaA1(OH)4(aq) that are list in the last column of Table 6. In addition to our heat capacities of NaAl(OH)4(aq) for five temperatures (1W150 C) listed in Table 6, we have experimental heat capacities at four temperatures (10-55 C) from Hovey et al. (I) and at five temperatures ( C) from Caiani et al. (2). As our first step toward assessing and then using these standard state heat capacities for NaAl(OH)4(aq) at 14 temperatures from 10 to 245"C, we have made a graph of CP0 against T (Fig. 1) and observed that 13 of the results fall on a reasonably smooth curve. We have therefore proceeded with analysis, omitting the CP0 from Caiani et al. (2) for 53"C, which seems to be inconsistent with the other 13 heat capacities from three independent investigations. An equation proposed by Holmes and Mesmer (14) has been used by them and others for convenient representation of standard state heat capacities of aqueous electrolytes. We have fitted their equation to the results of Hovey et al. (I), Caiani et al. (2), and the present results to obtain and also to values of (HOT - H ~ ~ for ~ this ~ substance., ~ ~ ) for NaAl(OH)4(aq). For this fitting we have weighted the dir- Hemingway and Robie (6) have applied hydrofluoric acid cal- ect heat capacities (1, 2) twice as heavily as we have our own orimetry to determine AfHO = f 1.19 kj mol-l heat capacities (obtained by differentiation of AH0 values). for gibbsite at T = K. We have combined all of these The standard deviation of the fit of eq. [5] is 12 J K-I mol-', properties with those of the elements (7) to obtain the thermo- which is very good. We have also used the same heat capadynamic properties of gibbsite that are listed in Table 4. cities of NaAl(OH)4(aq) for fitting the HKF equation (15) For NaOH(aq) at T = K we select AfHO = and find that the standard deviation of the fit is twice as large f kj mol-i and So = * 0.25 J K-' mol-' from as the 12 J K-I mol-i found above for the Holmes-Mesmer the CODATA critical compilation (8). Standard state partial molar heat capacities of NaOH(aq) have been selected from results of five calorimetric investigations (9-13). All of these results fit together smoothly and apparently reliably over the temperature range of interest to us. We have combined these properties with those of the elements (7) to obtain the thermodynamic properties of NaOH(aq) that are listed in Table 5. equation. The Holmes-Mesmer equation fits the experimental data better than does the HKF equation, but it should be noted that the former equation has more adjustable parameters than does the latter. Because the fit of the Holmes-Mesmer equation is better, we use this equation in some further calculations. Although we are using the Holmes-Mesmer equation here, let us point out that the HKF equation has a substantial value.

4 CAN. J. CHEM. VOL. 69, 1991 TABLE 4. Thermodynamic properties of gibbsite, Al(OH)? t GO SO HOT - ~ ~ A~HO ("c) (J K-' mol-i) (J K-' mol-') (kj mol-') (kj mol-i) TABLE 5. Thermodynamic properties of NaOH(aq) t GO So HOT - ~' A~HO ("C) (J K-' mo1-') (J K-' mo1-i) (kj mo1-') (kj mo1-i) TABLE 6. Enthalpies of formation and partial molar heat capacities of NaAl(OH),(aq) T Af~OT AfffOT - AfH AfH * ("C) (kj mo1-' ) (kj mo1-') (kj mo1-') c," (J K-' mol-') Hovey et al. (1) used their heat capacities over the relatively narrow range from 10 to 55 C to evaluate the parameters in the HKF equation and thence estimated standard heat capacities of NaA1(OH)4(aq) at much higher temperatures. Although the experimental results from the present investigation and from Caiani et al. (2) for high temperatures are to be preferred to the values estimated by Hovey et al. (I), it should be recognized that the estimated values are close enough to the better experimental values to be useful for a variety of thermodynamic calculations. On the basis of the calculations summarized above and a few similar calculations for other electrolytes, we offer the following generalizations. (i) The Holmes-Mesmer equation is usually better than the HKF equation for accurate represen-

5 CHEN ET AL. TABLE 7. Thermodynamic properties of NaA1(OH)4(aq) t C," SO H0 - ~ ~29a.15 AfHO A~GO ("C) (J K-' mol-i) (J K-' mol-i) (kj mol-i) (kj mol-i) (lhmol-i) TABLE 8. Thermodynamic properties of A1(OH)4-(aq) t GO so H0 - ~ ~ AfHO A~GO ("C) (J K-I mol-i) (J K-' mol-i) (kj mol-i ) (kj mol-i) (kj mol-i) - calculated A Hovey et al(1) x Caiani at al(2) this woh FIG. 1. Standard state partial molar heat capacities of N&~(oH),- (aq) from various sources. tation of experimental heat capacities of aqueous electrolytes over a wide range of temperatures. (ii) The HKF equation is most useful in the common case where one has heat capacities for a narrow range of temperatures (say C) and wants to estimate heat capacities for higher temperatures. In Table 7 we summarize the thermodynamic properties of NaA1(OH)4(aq). Calculations leading to the tabulated quantities have made use of properties already cited and the smoothed heat capacities obtained from the Holmes-Mesmer equation [5]. Next we use the thermodynamic properties for NaAl(OH)4- (aq) that are listed in Table 7 with the thermodynamic properties of HCl(aq) and NaCl(aq) to calculate the conventional thermodynamic properties of Al(OH)4-(aq) that are listed in Table 8. All of these tabulated properties are based on the commonly used convention that A~GO, A~H', SO, and cp0 of H+(aq) are zero at every temeprature. For these calculations we have used thermodynamic properties from this investigation and from refs. 8, Thermodynamic properties listed in Tables 7 and 8 are related to the earlier work by Chang (19, 20), Kuyunko et al. (21), and Apps et al. (22, 23). It is a convincing indication of thermodynamic consistency of a variety of experimental I. -C results that the Gibbs energies of A1(OH)4-(aq) reported by Apps et al. (22, 23) are in good agreement with our values that are based on calorimetric measurements. It appears (24) that

6 1690 CAN. J. CHEM. VOL. 69, 1991 consideration of all of these results together leads to selection of "best" A~HO and SO values that are closer to those we have selected here than to those selected earlier (23) by Apps et al. Acknowledgements We thank the Natural Sciences and Engineering Research Council of Canada for support of this and related research. We also thank A. Yin for providing seed crystals of gibbsite and advice about chemical preparation. Finally, we thank J. K. Hovey for useful advice on our calculations. 1. J. K. HOVEY, L. G. HEPLER, and P. R. TREMAINE. J. Phys. Chem. 92, 1323 (1988). 2. P. CAIANI, G. CONTI, P. GIANNI, and E. MAT~EOLI. J. Solution Chem. 18, 447 (1989). 3. A. YIN et al. To be published. 4. Y. XU. Thesis, University of Alberta B. S. HEMINGWAY, R. A. ROBIE, J. R. FISHER, and W. H. WILSON. J. Res. U.S. Geol. Survey, 5, 797 (1977). 6. B. S. HEMINGWAY and R. A. ROBE. J. Res. U.S. Geol. Survey, 5, 413 (1977). 7. M. W. CHASE et al. JANAF thermochemical tables, 3rd ed. (J. Phys. Chem. Ref. Data, 14, Supplement No. 1 (1985)) J. D. Cox, D. D. WAGMAN, and V. A. MEDVEDEV. CODATA key values for thermodynamics. Hemisphere Publishing Corporation, New York P. P. SINGH, E. M. WOOLLEY, K. G. MCCURDY, and L. G. HEPLER. Can. J. Chem. 54, 3315 (1976). 10. G. C. ALLRED and E. M. WOOLLEY. J. Chem. Thermodyn. 13, 147 (1981). 11. A. H. ROUX, G. PERRON, and J. E. DESNOYERS. Can. J. Chem. 62, 878 (1984). 12. G. CONTI, P. GIANNI, A. PAPINI, and E. MAT~EOLI. J. Solution Chem. 17,481 (1988). 13. J. M. SIMONSON, R. E. MESMER, and P. S. Z. ROGERS. J. Chem. Thermodyn. 21, 561 (1989). 14. H. F. HOLMES and R. E. MESMER. J. Phys. Chem. 87, 1242 (1983). 15. H. C. HELGESON, D. H. KIRKHAM, and G. C. FLOWERS. Am. J. Sci. 281, 1249 (1981). 16. P. R. TREMAINE, K. SWAY, and J. A. BARBERO. J. Solution Chem. 15, 1 (1986). 17. K. S. PITZER and J. C. PEIPER. J. Phys. Chem. Ref. Data, 13, 1 (1984). 18. B. S. HEMINGWAY, R. A. ROBIE, and J. A. KITTRICK. Geochim. Cosmochim. Acta, 42, 1533 (1978). 19. B.-T. CHANG. Bull. Chem. Soc. Jpn. 54, 1960 (1981). 20. B.-T. CHANG. Bull. Chem. Soc. Jpn. 55, 1949 (1982). 21. N. S. KUYUNKO, S. D. MALININ, and I. L. KHODAKOVSKY. Geochim. Int. 20, 76 (1983). 22. J. A. APPS and J. M. NEIL. Am. Chem. Soc. Symp. Ser. 416, 414 (1990). 23. J. A. APPS, J. M. NEIL, and C.-H. JUN. Lawrence Berkeley Laboratory, LBL Q. CHEN and L. G. HEPLER. To be published.