In the Laboratory
Hydrolysis Studies and Quantitative Determination of Aluminum Ions Using 27Al NMR An Undergraduate Analytical Chemistry Experiment Maria A. Curtin,* Laura R. Ingalls, Andrew Campbell, and Magdalena James-Pederson Department of Chemistry, Stonehill College, Easton, MA 02357; *
[email protected] The hydrolysis of cations is a topic that we often leave to a short discussion during the analytical chemistry sequence of undergraduate chemistry courses, although the subject is essential to understanding equilibrium and speciation in aqueous solutions. Metal ions in aqueous solution have always been complicated systems. Metal ion hydrolysis is complex (1) and difficult to understand owing to the lack of instrumentation capable of distinguishing directly between the different hydrolysis species of metal ions. Traditionally, studies rely on accurate potentiometric methods of analysis with careful control of ionic strength but without direct observation of the species involved (2). So the knowledge we transfer to undergraduates must be accepted as true without a real sense of what it means in an experimental lab. This is a big problem since it is essential knowledge in the study and understanding of aqueous systems especially as it applies to environmental issues (3–5). A thorough search of the chemical education literature shows a lack of experiments on either metal ion hydrolysis or 27Al NMR for undergraduates. Aluminum is an excellent element to explore metal ion hydrolysis since it exists in a single oxidation state in solution and hydrolyzes at moderately acidic pHs (6). Aluminum is found in nature as a single isotope with a nuclear spin of 5/2 making it NMR active. In this experiment we look at the NMR peak for the monomeric aluminum hexaaquocomplex, Al(H2O)63+, as a function of pH and we observe the effect of hydrogen ion concentration on this peak. Similar to 13C or 1H NMR of organic molecules, 27Al NMR spectroscopy will show different resonances for aluminum ions with different chemical environments. Because of the nuclear spin of 5/2, aluminum nuclei have a quadrupole moment that results in broad resonance peaks (7). The chemical shifts depend largely on the type of ligand and coordination number, while line widths are strongly influenced by symmetry and the rate of chemical exchange among aluminum species in equilibrium (8, 9). Background Many articles have been written on the subject of aluminum speciation analysis using 1H NMR (10) and 27Al NMR (11–13). Using 27Al NMR, a single peak is observed at ~0 ppm attributed to Al(H2O)63+ (14) in acidic solutions, pH ≤ 1. As the aluminum ion hydrolyzes with increasing pH through addition of base, the peak at ~ 0 ppm decreases and broader peaks are observed downfield owing to the formation of different hydrolysis products. A simple stepwise hydrolysis model for aluminum ions in solution is shown below, however oligomeric and polymeric aluminum species can also be formed in aqueous solutions. Al(H2O)63 (aq) (1) Al(OH)(H2O)52 (aq) H (aq)
Al(OH)(H2O)52 (aq) (2) Al(OH)2(H2O)4 (aq) H (aq)
Al(OH)2(H2O)4 (aq) (3) Al(OH)3(s) H (aq) 3H2O
The small peak at 0.34 ppm has been attributed to the mono hydrolyzed species Al(OH)(H2O)52+ formed in the first step of hydrolysis, peaks at 0.77 ppm and 4.3 ppm have been attributed to small oligomeric polynuclear species, Al2(OH)2(H2O)84+ and Al3(OH)4(H2O)95+(15). These assignments are not definite, molecular modeling calculations indicate the Al(OH)(H2O)52+ peak might be further downfield at 9 ppm (16) or around 11.6 ppm (17). More complex aluminum hydrolysis products tend to produce even broader peaks, making their analysis much more difficult by NMR spectroscopy. A tridecamer or a polymeric species containing thirteen aluminums AlO4Al12(OH)24(H2O)127+ has been identified (18) and can be formed easily through controlled addition of base to aqueous aluminum chloride or nitrate solutions (19). The tridecamer is unusual because it contains a tetrahedral aluminum surrounded by twelve octahedral aluminums all bound to each other through oxygen. The tetrahedral aluminum has a distinct sharp peak at 62.5 ppm relative to the monomeric Al(H2O)63+. The other twelve aluminums in the tridecamer as well as other large amorphous polymeric aluminum species produce a very broad peak around 40 ppm that is not easily integrated because it is not clearly defined. One other distinct sharp resonance at around 80 ppm is observed in basic solutions containing aluminum ions and this is due to the tetrahedrally coordinated Al(OH)4− ions present. We can clearly differentiate by 27Al NMR the pure Al3+ aqueous ion, Al(H2O)63+, in dilute solutions from other forms of aluminum ions such as Al(OH)(H2O)52+, Al(OH)2(H2O)4+, and other hydrolysis species. The peak intensity of the monomeric unhydrolyzed aluminum decreases as the pH of the solution is increased and is proportional to the concentration present in the solution. In addition, Bertsch et. al. (20) have shown a large linear dynamic range of peak area versus concentrations for aqueous monomeric unhydrolyzed aluminum Al(H2O)63+ solutions. This makes it possible to relate the peak area attributed to Al(H2O)63+ to actual concentration of this species in solution. Experimental The purpose of this experiment is to use NMR spectroscopy as an analytical tool to determine total aluminum content in a sample and to get a better understanding of metal ion hydrolysis and metal ion speciation in aqueous solutions.
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 2 February 2008 • Journal of Chemical Education
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y = 379365.4057x R2 = 0.9994
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In the Laboratory
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Figure 1. Calibration curve for the quantitative determination of aluminum in aqueous solutions using 27Al NMR spectroscopy.
Figure 2. Al(H2O)63+ content for solutions containing 225 mg/L of total aluminum neutralized with Na2CO3 to various hydrogen concentrations.
A 0.1 M Al(NO3)3 at pH 1 in nitric acid is used as the standard reference (prepared as described in the online material) with a single sharp peak at 0 ppm. The samples were analyzed using a 300 MHz Bruker Avance spectrometer (27Al at 78.2 MHz), but an Anasazi 90 MHz multiprobe NMR can also be used (27Al at 23.4 MHz). The number of acquisition scans was limited to 250 to keep the analysis time to an average of 10 minutes per sample and Xwin-nmr (version 3.1) was used to analyze the spectra.
The solutions were then transferred to 50 mL volumetric flasks and diluted to the mark with de-ionized water. The pH of each solution was measured again after one hour of equilibration prior to NMR analysis. The Al(H2O)63+ peak shows an increase in width as pH increases, probably owing to either the presence of Al(OH)(H2O)52+ or the slow equilibrium exchange between species. In this case, it is assumed that the increase in width does not affect the peak height of the Al(H2O)63+ peak. The Al(H2O)63+ concentration was again determined by measuring the signal intensity of the peak at its highest point and finding the equivalent concentration from the calibration curve. The data shown in Figure 2 clearly show a decrease in the unhydrolyzed aluminum concentration of the samples as the pH increases. The results show a very clear curve for the hydrolysis of Al(H2O)63+ as the [H+] is decreased in the solution. Sodium hydroxide can be used instead of sodium carbonate but it is not recommended because it is difficult to control pockets of high [OH−] created during addition, which leads to some error owing to formation of Al(OH)3 solid aggregates. Since there is only one NMR instrument available, teams of students were asked to perform different tasks. The NMR time was limited to the four hours reserved for this experiment; therefore, one group of students was asked to prepare the standards and another to work on the hydrolysis experiment. All students were allowed to prepare at least one NMR tube and analyze the sample by NMR spectroscopy. The data were then pooled and given to each student to work with. The results of the pooled data showed a good correlation; however, the calibration curve had some obvious errors owing to differences in sample preparation and techniques between individual students. The students themselves reported one of the reasons for the errors was having more than one person prepare the samples. The hydrolysis experiment was a success. The student data were more scattered than shown in Figure 2; however, the trend was obvious and they all understood the general process of metal ion hydrolysis. A simplified calculation of the equilibrium constant for the first hydrolysis step, disregarding activities and activity coefficients, can be carried out. For this calculation, [Al(H2O)63+] is the concentration obtained from the NMR analysis using the standard curve. The concentration of the hydrolyzed species, [Al(OH)(H2O)52+], is then obtained by subtracting the concentration of unhydrolyzed
Hazards Concentrated nitric acid is highly corrosive and must be kept in a fume hood. Care should be taken not to inhale the vapors. The 2% nitric acid can be dispensed at the bench. Skin contact with the 2% solution will cause skin discoloration and even burns if not washed right away, so care should be taken to use gloves when making solutions and to keep benches clean. Aluminum nitrate is a strong oxidizer and contact with other material may cause fire. It is harmful if swallowed or inhaled and causes irritation to skin, eyes, and respiratory tract. Sodium carbonate may cause eye burns, is harmful if swallowed or inhaled, and causes irritation to skin and respiratory tract. Results and Conclusion The sample preparation was carried out as outlined in the laboratory instructions for this experiment (provided in the online material). To save time students were asked to only record the signal intensity of the unhydrolyzed aluminum peak at the highest point and plot this number versus concentration. The Al(H2O)63+ peak is very narrow and therefore the signal intensity can be easily substituted for the peak area or integral. The linear correlation is excellent as seen in one student’s curve as shown in Figure 1. The second part of the experiment required the preparation of several aluminum solutions all containing the same amount of aluminum but adjusted with Na2CO3 solution to different pHs. The solutions were prepared by taking equal aliquots of the 1000 mg/L standard and adding 20 mL of de-ionized water. The pH of each solution was monitored using a calibrated pH electrode and adjusted to the desired pH with 0.1 M Na2CO3. 292
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In the Laboratory
species, [Al(H2O)63+], from the total aluminum concentration. As expected the equilibrium constant calculated from the experimental data (i.e., log K = ‒2.7 from data in Figure 2) is higher than the constant reported in the literature (log K = ‒4.97) (21). The assumption that the only hydrolyzed species in solution is the Al(OH)(H2O)52+ is not correct. This species also undergoes hydrolysis and this is corroborated by the fact that the experimentally calculated value is not constant. The calculated equilibrium constant increases as the hydrogen ion concentration decreases. The higher the pH, the larger the discrepancy between the calculated Al(OH)(H2O)52+ concentration and the actual concentration. To determine the equilibrium constant more accurately, the aluminum concentration should be at least an order of magnitude more dilute as suggested by Baes and Mesmer to reduce the further hydrolysis of Al(OH)(H2O)52+. Also, as the amount of Al(H2O)63+ approaches zero, a discussion on signal to noise ratio and sensitivity becomes very important and can be expanded on in this experiment. Literature Cited 1. Henry, M.; Jolivet, L. Struct. Bonding (Berlin) 1992, 77, 155–206. 2. Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cations; John Wiley and Sons: New York, 1976; pp 9–52. 3. Martin, T. A.; Kempton, J. H. Environ. Sci. Technol. 2000, 34, 3229–3234. 4. Dabbs, D. M.; Ramachandran, U.; Lu, S.; Liu, J.; Wang, L.; Aksay, I. A. Langmuir 2005, 21, 11690–11695. 5. Saleh, F. Y.; Mbamalu, G. E.; Jarada, Q. H.; Brungardt, C. E. Anal. Chem. 1996, 68, 740–745. 6. Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cations; John Wiley and Sons: New York, 1976; pp 112–123. 7. Iggo, J. A. NMR Spectroscopy in Inorganic Chemistry. Oxford Chemistry Primers; Oxford University Press: New York, 1999; pp 23–25.
8. Garrison, J. M.; Crumbliss, A. L. Inorg. Chem. 1988, 27, 3058– 3060. 9. Karlik, S. J.; Tarien, E.; Elgavish, G. A.; Eichhorn, G. L. Inorg. Chem. 1983, 22, 525–529. 10. Fong, D.; Grunwald, E. J. Am. Chem. Soc. 1969, 91, 2413– 2422. 11. Bottero, J. Y.; Cases, J. M.; Flessinger, F.; Poirier, J. E. J. Phys. Chem. 1980, 84, 2933–2939. 12. Akitt, J. W.; Farthing, A. J. Chem. Soc., Dalton Trans. 1981, 1606–1608. 13. Akitt, J. W.; Farthing, A.; Howarth, O. W. J. Chem. Soc., Dalton Trans. 1981, 1609–1614. 14. Akitt, J. W.; Greenwood, N. N.; Khandelwal, B. L.; Lester, G. D. J. Chem. Soc., Dalton Trans. 1972, 604. 15. Perry, C. C.; Shafran, K. L. J. Inorg. Biochem. 2001, 87, 115– 124. 16. Tossell, J. A. J. Magn. Reson. 1998, 135, 203–207. 17. Kubicki, J. D.; Sykes, D.; Apitz, S. E. J. Phys. Chem. A 1999, 103, 903–915. 18. Johansson, G.; Lundgren, G.; Sillen, L. G.; Sodequist, R. Acta Chem. Scand. 1960, 14, 769. 19. Kloprogge, J. T.; Seykens, D.; Jansen, J. B. H.; Geus, J. W. J. NonCryst. Solids 1992, 142, 87–93. 20. Bertsch, P. M.; Barnhisel, R. I.; Thomas, G. W.; Layton, W. J.; Smith, S. L. Anal. Chem. 1986, 58, 2563–2565. 21. Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cations; John Wiley and Sons: New York, 1976; p 113.
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