ARTICLE pubs.acs.org/IECR
Decomposition Thermodynamics of Magnesium Sulfate Madeleine N. Scheidema* and Pekka Taskinen Department of Materials Science and Engineering, Aalto University, Vuorimiehentie 2, FI-02150 Espoo, Finland ABSTRACT: The stability of MgSO4 was evaluated with varying conditions using HSC Chemistry software and thermo analysis (TG-DSC). The available thermodynamic data were reviewed and the equilibrium calculations using these data were compared to the results of the thermo analysis, in order to select the most reliable data. An enthalpy of formation for MgSO4 of 1284.9 kJ/mol was used in the calculations. Equilibrium calculations of magnesium sulfate in the absence of any reducing agent resulted in a decomposition temperature of 1080 °C. This temperature can be lowered in the presence of a reducing agent, which reduces the energy requirement. The calculations resulted in a decomposition temperature of 625 °C, 600 and 610 °C in the presence of sulfur, carbon monoxide and hydrogen gas as a reducing agent, respectively. The thermo analysis showed that magnesium sulfate decomposed at 1100 °C under N2 atmosphere; at 1070 °C in an atmosphere of 10% CO and 90% N2; and at 950 °C in an atmosphere with 5% H2 and 95% N2.
1. INTRODUCTION In many hydrometallurgical processes, sulfuric acid is used for leaching. The sulfuric acid residue may contain metal sulfates in solution when the leaching process is completed and most of the metals of interest have been extracted. This acid has no direct use and can be hazardous from an environmental point of view. Before the European Community prohibited the dumping of sulfuric acid in 1993, dumping it in the sea was a common practice. In some cases, the acid was neutralized with lime or limestone prior to discarding it.1 Magnesium sulfate is present in the leaching residue of hydrometallurgical processes where magnesium rich ores or concentrates, such as laterites, are treated. Thus, magnesium sulfate is the main component of the waste stream. In plant-scale processes, the magnesium sulfate leaching residue is generally not pure; it contains minor amounts of other sulfates, for example iron-, aluminum-, and manganese sulfate. Regeneration of magnesium sulfate as a neutralizing agent will reduce the requirement of other neutralizing agents, such as lime, to the leaching process. An environmentally friendly way to dispose of sulfuric acid containing metal sulfates is thermal decomposition. The main decomposition products are metal oxide and sulfur dioxide. Thermal decomposition of magnesium sulfate has already been applied with the focus on producing sulfur dioxide for the production of sulfuric acid while the quality of the oxide product was not regarded.24 As a result, a waste stream with considerable amounts of metals is being left unused. The recycling of these metals would both reduce the waste stream and limit the need of additional substances into the leaching process.1,5 The magnesium oxide product can find various applications. In case the magnesium sulfate is obtained from a leaching process, the magnesium oxide can be recycled back into the process as a neutralizing agent, for pH control during the precipitation of other sulfates in that process.6 Other applications for MgO are as a raw material for the production of refractory materials and for the formation of MgCl2, which even can be used to produce electrolytic Mg.7 Reactive MgO is required when it is used as a neutralizing agent. Decomposition of magnesium r 2011 American Chemical Society
sulfate at 9001200 °C in air according to patent WO81021537 and decomposition in excess of 1000 °C according to patent US42255734 yields an MgO product of low reactivity. In practice, not the completely dehydrated form of magnesium sulfate is used as feed material; often it contains some crystal water, typically as monohydrate, MgSO4 3 H2O. This leads to the contribution of water vapor to the off gas. Completely dehydrated magnesium sulfate forms a fine dust as MgO product, which is difficult to recover in the gas train and process further. The formation of this fine dust can be prevented by leaving a small amount of crystal water in the feed material.3 Magnesium sulfate does not occur in the nature in anhydrous form, but only as hydrates. Magnesium sulfate hydrates that occur as natural deposits are epsomite and kieserite; magnesium sulfate with seven and one moles of crystal water, respectively. Stable and metastable forms of magnesium sulfate hydrate crystallized from solutions reported in the literature are hydrates with 111 mols of crystal water.8 Stepwise dehydration of a higher hydrate, for example epsomite, can produce hydrates with 6, 4, 3, 2, 1.25, and 1 mols of crystal water.9 The product is highly hygroscopic: at room temperature, defined as 25 °C, it will attract water. The amount of water that is attracted depends on the relative humidity: at a relative humidity of below 39%, starkeyite is formed, between 39 and 52% hexahydrate is the stable hydrate and at a relative humidity above 52%, epsomite is the stable hydrate.10Metastable phases are an unknown 2.4 hydrate in the range of 722% and pentahydrate between 22 and 39% relative humidity.10 Another recent publication on the stability of magnesium sulfate hydrates, at relatively low temperatures and varying relative humidity was made by Grevel and Majzlan.8 The aim of this study is to thermodynamically evaluate the stability of magnesium sulfate at a temperature in excess of 400 °C, with the most reliable thermodynamic data available. Received: December 22, 2010 Accepted: July 11, 2011 Revised: July 11, 2011 Published: July 11, 2011 9550
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The calculations are carried out with the HSC Chemistry 7.0 software, using the Equilibrium Module.11 The thermodynamic data available on magnesium sulfate are evaluated and compared with experimental results from thermo analysis.
2. THEORETICAL BACKGROUND Magnesium sulfate hydrate will dehydrate below 500 °C to form anhydrous magnesium sulfate. Different temperatures for dehydration are reported in the literature: 330365 °C,12 400500 °C9 and at least at 450 °C.12 Magnesium sulfate will decompose into magnesium oxide and sulfur trioxide when the temperature is 700 °C or higher, depending on the atmosphere under which the decomposition is carried out. Sulfur trioxide is not stable at temperatures in excess of 700 °C and spontaneously forms sulfur dioxide and oxygen. In the presence of a reducing agent, such as carbon, the decomposition temperature lies around 750 °C. In the absence of any reducing agent, magnesium sulfate decomposes in the temperature range of 9001100 °C.14 2.1. MgSO4: Different Prefixes Used in Literature. Magnesium sulfate is reported in the literature to exist in different forms: magnesium sulfate, R-magnesium sulfate, β-magnesium sulfate, and γ-magnesium sulfate. It appears that the different prefixes are used to indicate both the crystal structure and to make a distinction between the phase that is stable at high and low temperature. However, not in all literature sources a distinction has been made between different structural forms of MgSO4. In crystallographic literature sources, the prefix R or β is used to make a distinction between variations in the crystal structure. This difference is often a result of the production method that is used to obtain the crystals. The R-phase is grown from an aqueous solution of MgO and H2SO4; the β-phase can be formed either by dehydration of MgSO4 hydrates or by heating R-MgSO4 to 595 °C. It can be quenched to room temperature, where it persists as β-MgSO4 (possibly metastable).15 In the crystal data book,16 two polymorphs of MgSO4 are given, both orthorhombic in crystal structure, but with different space groups. As the β-phase is claimed to be the form that is commercially available,15 data of the R-phase from crystallographic sources will not be regarded in this paper. Fortes et al.15 mention a third form; γ-MgSO4, which is referred to be the same material as the high-temperature form in the article by Rowe et al.,17 which is obtained when heating β-MgSO4. The R-MgSO4 and β-MgSO4 can coexist below 527 °C according to Fortes et al.15 These two phases can also coexist according to Du,18 having a difference in enthalpy of formation of 3.01 kJ/mol. Different values for the transition temperature from the low-high temperature forms are given by different literature sources: the transition seems to take place between 997 and 1095 °C.1517,19 2.2. Theory of MgSO4 Decomposition. The reaction for the decomposition of anhydrous magnesium sulfate by heat alone is given in eq 1. The products are magnesium oxide and sulfur trioxide. Sulfur trioxide is not stable at elevated temperatures, >700 °C, and spontaneously forms sulfur dioxide and oxygen, as shown in eq 2. MgSO4 f MgO þ SO3 ðgÞ SO3 ðgÞ f SO2 ðgÞ þ
1 O2 ðgÞ 2
ð1Þ ð2Þ
As oxygen is generated in reaction 2, reducing conditions should lower the decomposition temperature of the sequence 1
and 2. The reducing agents used in current research are solid carbon or an atmosphere containing carbon monoxide gas, obtained from incomplete combustion of a carbonaceous fuel.7,20 The theoretical decomposition products are magnesium oxide, sulfur dioxide, and carbon dioxide, assumed that stoichiometric amounts of the feed materials are used. The overall reactions are given in eq 35 and eq 47,20 for solid carbon and carbon monoxide, respectively. The use of a deficiency of carbon leads to partial decomposition of MgSO4. On the other hand, with the use of an excess of carbon, species such as COS, CS2, and Sx may form. When a hydrated form of magnesium sulfate is used, H2S might be formed when the decomposition gases react with the evaporated crystal water.21 MgSO4 þ
1 1 C f MgO þ SO2 þ CO2 2 2
MgSO4 þ CO T MgO þ SO2 þ CO2
ð3Þ ð4Þ
Alternative reducing agents to carbon that can be possibly used are sulfur (gas)6,22 or hydrogen gas.7,23 The decomposition with sulfur takes place according to eq 56 and with hydrogen according to eq 6.7 The use of these reducing agents instead of a carboncontaining reducing agent results in an off gas free of CO2 and CO. 1 3 MgSO4 ðsÞ þ S2 ðgÞ f MgO þ SO2 ðgÞ 4 2
ð5Þ
MgSO4 þ H2 T MgO þ SO2 þ H2 OðvÞ
ð6Þ
3. EXPERIMENTAL SET UP AND RESULTS 3.1. Thermodynamic Data Selection. In the HSC Main Database,11 thermodynamic data of MgSO4 from different sources are available. For the equilibrium calculations, the most reliable data should be selected from this database. In general, all sources give more or less the same thermodynamic properties for MgSO4, except, however, the standard enthalpy of formation. Table 1 shows an overview of different values of enthalpy of formation of MgSO4 available from the literature. The value of the standard enthalpy of formation of MgSO4 generally lies between 1255 and 1289 kJ/mol. The two most frequently used values are 1261.8 kJ/mol24,25 and 1284.9 kJ/ mol.13,18,26,27 A recent value proposed by Landolt-B€ornstein (2000)19 is 1288.8 kJ/mol, which is accepted in the SGTE database [MTDATA28], too. The choice of standard enthalpy of formation has a large impact on the calculated stability of MgSO4 calculated as a function of temperature as can be seen from Figure 1 and Figure 2. With an enthalpy of formation of MgSO4 of 1261.8 kJ/mol, all MgSO4 is decomposed at 998 °C and with 1284.9 kJ/mol, MgSO4 is fully decomposed at 1085 °C. The choice of enthalpy of formation of MgSO4 thus results in a variation of the calculated decomposition temperature of 87 °C. In further equilibrium calculations, the value of 1284.9 kJ/mol will be used. The choice of this value is motivated in the discussion chapter. The conversion of MgSO4 to MgO does not occur sharply, as may be expected. Nitrogen and decomposition gases in the system dilute the atmosphere and the transition from sulfate to 9551
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Table 1. Overview of Enthalpy of Formation of MgSO4 Given by Various Literature Sources, With References enthalpy of source
species
formation (kJ/mol) year
used reference
Binnewies 0224
MgSO4
1261.80
2002
Landolt 0119
MgSO4
1288.76
2001 V.P. Glushko et al., unpublsihed results.
Du18
R-MgSO4
1288.80
1999 SGTE data for pure elements, Calphad, 1991, 15, 317425.
β-MgSO4
1285.79
1999
JANAF 9825
MgSO4
1261.80
1998 Knopf, H. J., Staude, H., Z. Physik. Chem., Leipzig, 1955. 204, 265275.
Pankratz 9529
R-MgSO4
1288.80
1995 DeKock, C. W., Thermodynamic Properties of Selected Metal Sulfates and their
β-MgSO4
1284.95
1995
MgSO4 (cr)
1284.91
1993 Wagman, D. D., Evans, W. H., Parker, V. B., Schumm, R. H., Halow, I., Bailey, S. M.,
Barin 9326
Hydrates. USBM IC 9081, 1986, 59pp. Based on Ko and Daut, correction for sulfate ion. Churney, K. L., Nuttall, R. L. The NBS Tables of Chemical Thermodynamic Properties, Washington DC. 1982.
Ko and Daut13
R-MgSO4
1288.46
1980
β-MgSO4
1284.61
1980
MgSO4
1309.84
1937 Kelley, K. K, Contributions to the Data of Theoretical Metallurgy. The Thermodynamic Properties of Sulfur and its Inorganic Compounds, BuMines Bull. 406, 1937, 154 pp.
MgSO4
1258.97
1897 Berthelot, M.P., Thermochimie. Gauthier-Villars Fils, Paris, France, 1897. p. 261.
MgSO4
1264.87
1883 Thomsen, J., Termochemische Untersuchungen. Johann Ambrosium Barth,
MgSO4
1284.91
1971
MgSO4
1278.21
1952 Averaged values: Thomsen, Berthelot and Ilosvay, Pickering and Kelley
Leipzig, 1883, 567 pp. Parker, V.B, Wagman, D.D. Evans, W.H.27 Rossini F.D., Wagman, D.D., Evans, W.H., Levine, S., Jaffe, I30
Figure 1. Decomposition equilibrium of 1 kmol MgSO4 in 0.1 kmol N2(g), with a standard enthalpy of formation for MgSO4 of 1261.8 kJ/mol.
oxide therefore occurs over a temperature range rather than at a certain temperature. 3.2. HSC Chemistry Equilibrium Diagrams. The thermodynamics of decomposition reaction 1 was evaluated with HSC Chemistry 7.0 software,11 using the Equilibrium Module. With this calculation approach, the stability of magnesium sulfate was evaluated as a function of temperature at ambient pressure. The thermodynamic properties of JANAF25 were used in the calculations, but with an enthalpy of formation of 1284.9 kJ/mol instead of 1261.8 kJ/mol.
Figure 2. Decomposition equilibrium of 1 kmol MgSO4 in 0.1 kmol N2(g), with a standard enthalpy of formation for MgSO4 of 1284.9 kJ/mol.
Equilibrium calculations were made with gaseous carbon monoxide, sulfur and hydrogen as reducing agents. As feed materials, 1 kmol of MgSO4 and 0.1 kmol of N2 were used in each calculation. Anhydrous magnesium sulfate was used in the calculations and not the hydrate, in order to avoid the presence of water vapor in the closed system. In reality, it is likely that a hydrate is used as the feed material, from which the crystal water will evaporate and leave the system prior to decomposition. The stoichiometric amount of each reducing agent was used, which is respectively, 1 kmol CO, 0.25 kmol S2 and 1 kmol of H2. 9552
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Figure 3. Decomposition equilibrium of magnesium sulfate with carbon monoxide: 1 kmol MgSO4 with 0.1 kmol N2 and 1 kmol CO(g).
Figure 4. Decomposition equilibrium of magnesium sulfate in the presence of sulfur: 1 kmol MgSO4 with 0.1 kmol N2 and 0.25 kmol S2(g).
The results of these calculations are shown in Figure 3, Figure 4 and Figure 5, respectively. The stoichiometric amount of reducing agent was used in the calculations, because a deficiency leads definitely to a partial decomposition. These thermodynamic calculations show that with a stoichiometric amount of the reducing agent in the system, MgO is the stable magnesium species above 590 °C, which means a decrease of the decomposition temperature of MgSO4 with approximately 495 °C compared to the calculation without reducing agents, as shown in Figure 2. In the presence of carbon monoxide as a reducing agent, magnesium sulfate is fully decomposed to magnesium oxide at a temperature of 592 °C, as can be seen from Figure 3. Upon decomposition, 1 kmol of CO2 and 1 kmol of SO2 gas is formed provided no side-reactions occur. With sulfur gas as a reducing agent, magnesium sulfate is fully decomposed to magnesium oxide at 623 °C. Per kmol of MgSO4, 1.5 kmol of sulfur dioxide is formed during decomposition and there is no carbon-containing species present in the off gas, as can be seen from Figure 4. Using hydrogen as a reducing agent, as shown in Figure 5, all
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Figure 5. Decomposition equilibrium of magnesium sulfate in the presence of hydrogen: 1 kmol MgSO4 with 0.1 kmol N2 and 1 kmol H2(g).
magnesium sulfate is converted to magnesium oxide at a temperature of 606 °C. One kmol of SO2 and one kmol of water vapor are formed per kmol of MgSO4; the hydrogen gas reacts with the oxygen released according to eq 2 to form water vapor. At temperatures below 606 °C, H2S gas is stable, which may be unwanted. However, since at temperatures below 606 °C not all magnesium sulfate is decomposed to magnesium oxide, it is not likely that such a low temperature will be used. However, the presence of H2S gas should be kept in mind, for example when the off gas is further processed. 3.3. Thermo Analysis. Thermo analysis was carried out with a NETZSCH STA 409 C/3/F. In this equipment, thermogravimetry (TG) and differential scanning calorimetry (DSC) are carried out simultaneously.31 The aim of these experiments is to use them as a comparison with the thermodynamic calculations and also to help selecting the correct enthalpy of formation of MgSO4 for the thermodynamic calculations, besides the literature review that has been carried out. Magnesium sulfate heptahydrate was subjected to different atmospheres at increasing temperature; with a heating rate of 5 °C/min. A sample mass of 29.000 mg was used in each experiment. Figure 6 shows the results of the experiment in 100% N2; Figure 7 shows the results in 90% N2 and 10% CO and Figure 8 shows the results of the experiment with 95% N2 and 5% H2. The CO gas flow in Figure 7 was started when a temperature of 700 °C was reached; in order to prevent carbon formation on the surface of the magnesium sulfate powder. In each case, there is an initial mass decrease at low temperatures (
