Equilibrium Study in the KNO3+ NH4NO3+ H2O System at

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Cite This: J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Equilibrium Study in the KNO3 + NH4NO3 + H2O System at Temperatures from 293.15 to 323.15 K Adriana Wrob́ el-Kaszanek, Sebastian Drużynś ki,* Urszula Kiełkowska, and Krzysztof Mazurek

J. Chem. Eng. Data Downloaded from pubs.acs.org by IOWA STATE UNIV on 02/04/19. For personal use only.

Faculty of Chemistry, Nicolaus Copernicus University in Toruń, 7 Gagarin Street, 87-100 Toruń, Poland ABSTRACT: A solid−liquid equilibrium (SLE) study was conducted in the ternary system KNO3 + NH4NO3 + H2O using the isothermal solution saturation method at temperatures of 293.15, 303.15, 313.15, and 323.15 K. The experimental data obtained were used to plot a section of the solubility polytherm of the investigated system and show the dependence of the densities of equilibrium solutions on their chemical compositions. It has been found that a complex salt of the formula K0.25(NH4)0.75NO3 is formed in the system at a temperature below 303.15 K.

1. INTRODUCTION The soda-chlorine-saltpeter (SCS) method is a modified version of the Solvay method, which is the most popular method of sodium carbonate production in the world. This method produces practically no waste and its main advantage is that it gives three products at a time: soda, chlorine, and a solution consisting of approximately 80% of ammonium nitrate(V), about 18% of sodium nitrate(V), and 2% of sodium chloride.1 The authors of the SCS method assumed direct production of mixed fertilizer by crystallizing ammonium-sodium saltpeter. However, according to studies,2−4 chloride ions accelerate thermal decomposition of ammonium nitrate(V) causing a real risk of explosion during the concentration, crystallization, and storage of ammonium-sodium saltpeter. The problem can be solved by converting thermally unstable ammonium nitrate(V) to potassium nitrate(V), which is characterized by much higher thermal stability. The reaction of the conversion of ammonium nitrate(V) is carried out with the use of potassium metavanadate according to eq 1. NH4NO3 + KVO3 ↔ NH4VO3 ↓ + KNO3

NH4NO3 + NH4VO3 + H2O, and NH4VO3 + KVO3 + H2O in the temperature range tested. This knowledge will enable determination of the triple points and lines separating the areas of the cocrystallization of salts in an oblique projection on the plane according to Jänecke’s method. Sufficient data on the NH 4 VO 3 −KVO 3 −H 2 O 5 and NH4NO3−NH4VO3−H2O6 systems have been found in the available literature and equilibrium data on the KVO3−KNO3− H2O system have been published.7 The literature does not provide complete data on the ternary system KNO3− NH4NO3−H2O, and so equilibrium studies need to be conducted on this system.

2. EXPERIMENTAL SECTION 2.1. Apparatus and Reagents. During this study, a Polyscience MX water thermostat operating with an accuracy of ±0.02 K was used. The temperature was also controlled using an electronic thermometer with an accuracy of 0.01 K. Analytically pure reagents and deionized water were used in the study (Table 1). 2.2. Methods. The equilibrium study in the KNO3 + NH4NO3 + H2O system was carried out using the isothermal solution saturation method at temperatures from 293 to 323 K. Appropriate amounts of potassium nitrate(V) and ammonium nitrate(V) (of which one salt was used in excess of its solubility in water) were placed in 150 cm3 Erlenmeyer flasks. Appropriate amounts of deionized water were then added, and the flasks were provided with magnetic stirrers and rubber stoppers. The solutions were thermostated and stirred until the phase equilibrium was reached. The time required to reach equilibrium

(1)

Figure 1 shows a schematic diagram of the process. The reaction products are sparingly soluble ammonium metavanadate and a solution of potassium nitrate(V) and unreacted sodium nitrate(V). Ammonium metavanadate is processed in a known manner to potassium metavanadate,5 whereas the potassium-sodium saltpeter solution is concentrated and crystallized. Substrates and products of the conversion reaction (eq 1) form a quaternary system of exchangeable salt pairs. Determining the optimal conditions for this reaction requires thorough knowledge of mutual solubility isotherms for the quaternary system NH4NO3 + KVO3 + H2O and for the ternary systems: KNO3 + NH4NO3 + H2O, KNO3 + KVO3 + H2O, © XXXX American Chemical Society

Received: November 8, 2018 Accepted: January 23, 2019

A

DOI: 10.1021/acs.jced.8b01052 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

4NH4NO3 + 6HCHO → (CH 2)6 N4 + 4HNO3 + 6H 2O (2)

The concentrations of potassium ions in the equilibrium solutions were determined by flame atomic absorption spectroscopy. The tests were carried out using a SavantAA Sigma GBC spectrometer. The solid phase composition was analyzed for selected experimental points using a Philips X-Pert PRO powder diffractometer. To determine the composition of the solid phases, the diffractometric data obtained were compared with appropriate standards in the “Powder Diffraction File”.8−24

3. RESULTS AND DISCUSSION The obtained experimental data are compared with literature values for the binary systems KNO3 + H2O and NH4NO3 + H2O at temperatures ranging from 293.15 to 323.15 K in Tables 2 and 3. The solubility of salts is given in g·100 g−1 H2O. Table 2. Comparison of Solubility Data Presented in This Work with the Literature Data for the Binary System NH4NO3−H2O6,11,12 g·100 g−1 H2O

Figure 1. Schematic diagram of the conversion of ammonium nitrate(V) to potassium metavanadate.7

ammonium nitrate(V) potassium nitrate(V) deionized water

source Avantor Performance Materials Poland Avantor Performance Materials Poland

mol fraction purity 0.99

this work

ref 6

ref 11

ref 12

293.15 303.15 313.15 323.15

188.7 223.4 296.6 324.0

181.7 222.6 275.1 318.1

192 242 296

189.9 236.1 288.7 349.3

Table 3. Comparison of Solubility Data Presented in This Work with the Literature Data for the Binary System KNO3− H2O11,13,14

Table 1. Reagents Used in the Studies chemical name

T/K

purification method

g·100 g−1 H2O

none

0.99

none

0.06 μS cm−1a

none

T/K

this work

ref 11

ref 13

ref 14

293.15 303.15 313.15 323.15

31.64 45.82 66.88 85.58

31.6 45.3 61.3

31.93 45.56 62.87 84.16

31.15 44.60 67.70

a

Conductivity of water used in the studies.

The solubility values of ammonium nitrate(V) in water vary significantly between individual authors. These differences are particularly visible at temperatures of 313.15 K (21.44 g·100 g−1 H2O) and 323.15 K, where the difference reaches 31.21 g·100 g−1 H2O. These discrepancies result particularly from the very strong dependence of the solubility of ammonium nitrate(V) on temperature. Even a slight change in temperature during sampling causes crystallization and a change in concentration. Another factor that affects the discrepancies is the sampling procedure, which could differ depending on the author. Considering the difficulties in sampling saturated solutions of NH4NO3 at higher temperatures, the presented experimental and literature results are acceptable. The maximum differences in results do not exceed 10% of the mean value at 323.15 K. The solubility results for the KNO3−H2O system presented in Table 3 show good correlation with the literature values in the studied temperature range. Experimental data obtained for the KNO3 + NH4NO3 + H2O system in the temperature range from 293.15 to 323.15 K are presented in Table 4. The individual columns show the densities of solutions in g cm−3, the concentrations of individual salts in mol 1000 g−1 H2O, the mole fractions of individual components

was 24 h.8 After this time, the stirring was switched off and the solutions were left for 24 h for sedimentation. Samples of clear solutions were taken to calibrated Ostwald’s pycnometers in order to determine their densities. Before sampling at higher temperatures (303.15−323.15 K), the pycnometers were thermostated at a given temperature to prevent salt crystallization during the filling of the vessels. After determining the densities, the pycnometers content was quantitatively transferred to 500 cm3 volumetric flasks and deionized water was added. The composition of the equilibrium solutions was determined based on the analyses of concentrations of individual ions. X-ray diffraction (XRD) analysis of solid phases was also carried out for selected experimental points. 2.3. Analytical Methods. The concentration of ammonium ions in the equilibrium solutions was determined by the Ronchese formalin method. Formaldehyde reacts with ammonium ions forming urotropin and an equimolar amount of mineral acid (eq 2). The resulting acid is titrated with the standard solution of NaOH in the presence of phenolphthalein.9,10 B



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Table 4. Solubility Data in the NH4NO3−KNO3−H2O System as a Function of KNO3 Mole Fraction in Mixture with NH4NO3 (x) and Density (d) at Temperature T and Pressure p = 0.1 MPaa m

m

(mol·kg−1H2O)

d

x

no.

(g·cm−3)

KNO3

NH4NO3

KNO3

1 2 3 4 (E1)

1.307 1.314 1.322 1.338

0.000 0.628 1.065 1.786

T = 293.15 22.70 22.67 22.61 22.48

K 0.0000 0.0270 0.0450 0.0736

5 6 7 (E2)

1.348 1.352 1.367

3.165 3.854 4.917

21.49 20.36 19.30

0.1284 0.1592 0.2030

8 9 10 11 12 13 14 15 16 17

1.350 1.322 1.277 1.250 1.222 1.212 1.198 1.188 1.170 1.157

4.614 4.199 3.542 3.405 3.226 3.204 3.148 3.105 3.195 3.333

1 2 3 4 5 6 7 8 (E) 9 10 11 12 13 14 15

1.324 1.333 1.346 1.356 1.367 1.377 1.384 1.380 1.366 1.302 1.288 1.257 1.242 1.220 1.216

16.27 12.60 7.544 5.679 3.915 3.328 2.615 1.728 0.880 0.000 T = 303.15 0.000 27.81 0.804 27.39 1.466 27.14 2.350 26.71 2.938 26.19 3.947 25.75 4.911 24.95 6.335 24.27 5.422 15.28 4.649 7.574 4.431 4.974 4.326 3.416 4.341 2.489 4.406 0.812 4.464 0.389

0.2209 0.2500 0.3195 0.3748 0.4518 0.4905 0.5462 0.6425 0.7840 1.0000 K 0.0000 0.0285 0.0512 0.0809 0.1009 0.1329 0.1645 0.2170 0.2619 0.3803 0.4711 0.5588 0.6356 0.8444 0.9198

d solid phase composition

no.

NH4NO3 NH4NO3 NH4NO3 NH4NO3 + (NH4)xK1‑yNO3 (NH4)xK1‑yNO3 (NH4)xK1‑yNO3 KNO3 + (NH4)xK1‑yNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 + KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3

X NH4NO3 =

1.217

1 2 3 4 5 6 7 8 9 (E) 10 11 12 13 14 15

1.338 1.352 1.355 1.363 1.371 1.382 1.395 1.392 1.406 1.421 1.405 1.399 1.353 1.324 1.269

1 2 3 4 5 6 7 (E) 8 9 10 11 12 13 14

1.356 1.372 1.385 1.400 1.413 1.432 1.436 1.433 1.424 1.410 1.383 1.368 1.340 1.332

NH4NO3

T = 303.15 0.000 T = 313.15 0.000 34.37 0.676 33.96 1.478 33.25 2.132 32.78 2.886 32.13 3.730 31.39 4.736 30.48 5.999 30.51 8.000 28.67 7.351 21.16 7.026 17.90 6.512 10.68 6.252 6.095 6.191 1.154 6.615 0.000 T = 323.15 0.000 39.56 1.518 38.09 3.132 37.91 4.718 36.15 6.409 35.33 8.069 34.52 9.951 32.73 8.801 19.45 8.358 15.59 7.917 10.46 7.719 7.066 7.540 0.992 7.723 0.292 8.465 0.000 5.189

KNO3 K 1.0000 K 0.0000 0.0195 0.0426 0.0611 0.0824 0.1062 0.1345 0.1643 0.2182 0.2578 0.2819 0.3788 0.5064 0.8429 1.0000 K 0.0000 0.0383 0.0763 0.1154 0.1535 0.1895 0.2331 0.3115 0.3490 0.4308 0.5221 0.8837 0.9636 1.0000

solid phase composition KNO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3+ KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3 NH4NO3+ KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3 KNO3

respectively, I, equilibrium solutions with a solid phase in the form of ammonium nitrate(V); II, equilibrium solutions with a solid phase in the form of potassium nitrate(V) and a section between points E1 and E2, equilibrium solutions saturated with a double salt. The courses of branches marked I of the solubility isotherms (Figure 2) in the studied temperature range are linear. With the increase of KNO3 concentration in the equilibrium solutions with a solid phase in the form of NH4NO3, the concentration of NH4NO3 decreases toward the points saturated with both salts (E). The points on branches marked II of the solubility isotherms (Figure 2) correspond to solutions saturated with KNO3 with an increasing concentration of NH4NO3 toward the eutonic points. The courses of branches marked II indicate the salting-in effect of NH4NO3 on KNO3. The increase in ammonium nitrate(V) concentration increases the solubility of potassium nitrate(V) in equilibrium solutions.

[KNO3] [KNO3] + [NH4NO3] [NH4NO3] [KNO3] + [NH4NO3]

16

KNO3

x

a Standard uncertainties: u(T) = 0.02 K; u(p) = 5 kPa; u(d) = 0.002 g cm−3; ur(m) = 0.02.

(excluding the solvent), and the composition of the solid phase in equilibrium with the saturated solution. The data contained in Table 4 were used to plot a section of the solubility isotherm (Figure 2) of the tested system and create a graph showing the dependence of the densities (Figure 3) of the equilibrium solutions obtained as a function of the mole fraction of KNO3 calculated in accordance with eq 3. X KNO3 =

(g·cm−3)

(mol·kg−1H2O)

(3)

The solubility isotherms (Figure 2) for 303.15, 313.15, and 323.15 K consist of two branches marked I and II, which correspond, respectively, to equilibrium solutions saturated with ammonium nitrate (N) and potassium nitrate(V). Eutonic solutions saturated with both salts were marked as points (E). The isotherm for 293.15 K (Figure 2) consists of three branches, C



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Figure 2. Section of the solubility polytherm for the NH4NO3−KNO3−H2O system in the temperature range from 293.15 to 323.15 K.

Figure 3. Dependence of the densities of equilibrium solutions on the mole fraction of KNO3.

There are many publications describing studies on solid phases appearing in the studied system. The outlined differences are primarily due to the tendency of potassium nitrate(V) and ammonium nitrate(V) to create metastable states in the form of solid solutions.9−17 Some authors confirm the existence of double salts; however, there is no consensus on their stoichiometry and the temperatures at which they appear.16−24 Some studies indicate the formation of double salts composed of 3KNO3·NH4NO3 and two rows of solid solutions at 298 K,20−22 but there are also reports of double salts: 2KNO3·NH4NO3 and KNO3·2NH4NO3.19 Other authors describe solid phases that are in fact solid solutions: K 0.75 (NH 4 ) 0.25 NO 3 23 and K0.67(NH4)0.33NO3.24 The description of the formation of solid solutions and the detailed equilibrium study in the KNO3−NH4NO3−H2O system indicate that the possibility of the formation of mixed salts or solid solutions in this system is related to similar values of

The dependence between the system’s properties, i.e., the densities and the chemical compositions of the equilibrium solutions, was used to determine the number of solid phases in the studied system (Figure 3). Points of discontinuity of the function on the curves mark new solid phases formed in the studied system.5−8 The nature of the curves in Figure 3 indicates that each curve has one point of discontinuity at 303.15, 313.15, and 323.15 K. These points represent isothermally stable (eutonic) solutions in equilibrium with KNO3 and NH4NO3 in the solid phase. The dependence of the densities of the equilibrium solutions on the mole fraction of KNO3 at 293.15 K has two points of discontinuity, which indicates the formation of a new solid phase of a mixed salt of KNO3 and NH4NO3. The analysis of XRD, the solid phase, and the point in the curve’s turning point confirms the presence of a new salt in the system, as shown in Figure 4. D



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Figure 4. Diffractograms for the solid phase (A) for pure KNO3; (B) for pure NH4NO3; (C) sample of solid solution at 293.15 K branch E1−E2; (D) E2 sample of solid solution and KNO3 in solid phase at 293.15 K; (E) E1 sample of solid solution and NH4NO3 in solid phase at 293.15 K; (F) for eutonic point at 303.15 K; (G) for eutonic point at 313.15 K; (H) for eutonic point at 323.15 K. ●, KNO3; ■, NH4NO3; red ▲, solid solution KNO3 and NH4NO3.

ionic radii: ammonium cation NH4+ (1.43 Å) and potassium cation K+ (1.33 Å), which can be exchanged in a solid phase crystal lattice within a mole fraction of ammonium nitrate(V)

ranging from 0.667 to 0.933. This phenomenon is described by the linear dependence of the changing volume of a unit cell on the solid phase composition. This is illustrated by eq 4.16−19 E



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Figure 5. TG curve of a solid phase sample formed in the KNO3−NH4NO3−H2O system at 293.15 K.

V = 35.6806·x NH4NO3 + 287.093

nitrate(V) on ammonium nitrate(V) and the salting-in effect of ammonium nitrate(V) on potassium nitrate(V). At 293.15 K, the formation of the double salt K0.25(NH4)0.75NO3 was observed, which is consistent with the literature.16−19 However, contrary to what some authors report,20−24 no other mixed salts were observed at higher temperatures. The results obtained complement the existing state of knowledge regarding equilibrium studies in the KNO3− NH4NO3−H2O system and are necessary in the subsequent stages of research on the quaternary system NH4NO3−KVO3− H2O. Plotting solubility isotherms of the quaternary system will be the basis for determining the most favorable conditions for the conversion of ammonium nitrate to potassium nitrate.

(4)

where V is the volume of a unit cell [Å3], and xNH4NO3 is the mole fraction of NH4NO3. The experimental data obtained at 293.15 K indicate unambiguously that a new solid phase is formed in the system under investigation. A sample of the solid phase obtained was subjected to chemical, XRD, and thermogravimetric (TG) analyses. The XRD analysis indicates that the compound obtained has a different diffractogram (Figure 4C) compared to pure single salts (Figure 4A,B) and their mixtures (Figure 4F− H). The volume of the unit cell in the mole fraction range xNH4NO3 = 0.667−0.933 for the mixed salt ranges from 310.8 Å3 to 319.9 Å3.19 The result obtained for the mixed salt sample was 313.1 Å3, which indicates the formation of a solid phase K0.33(NH4)0.66NO3 in the tested system. Thermal analysis was carried out in the temperature range from 300.65 K to 782.25 K with a heating rate of 10 K/min in the air atmosphere (Figure 5). Pure ammonium nitrate(V) begins to decompose at about 443.15 K and ends at about 563.15 K depending on the size of the sample.25 Based on the recorded TG curve, the calculated stoichiometric composition of the mixed salt sample is presented by the formula K0.252(NH4)0.748NO3. The chemical analysis, including the analysis of individual ions in the solution resulting from the dissolution of the analytical sample of the tested solid phase, confirms the thermogravimetric results obtained. The stoichiometric composition obtained of the solid phase tested by chemical analysis is K0.271(NH4)0.729NO3.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Sebastian Drużyński: 0000-0002-8762-2528 Funding

This work was supported by Grant No. 2341-Ch from UMK (Nicolaus Copernicus University). Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Collaborative paper, Research on the new method of soda production; Nicolaus Copernicus University: Toruń, Poland, 1969. (2) Kołaczkowski, A. Spontaneous decomposition of ammonium nitrate; Scientific Papers of the Institute of Inorganic Technology and Mineral Fertilizers; Wrocław Technical University: Wrocław, Poland, 1980.

4. CONCLUSIONS The results of the study of the mutual solubility of salt in the studied system indicate the salting-out effect of potassium F



Article

(3) Izato, Y.; Miyake, A. Thermal decomposition mechanism of ammonium nitrate and potassium chloride mixtures. J. Therm. Anal. Calorim. 2015, 121, 287−294. (4) Li, X.-R.; Koseki, H. Study on the contamination of chlorides in ammonium nitrate. Process Saf. Environ. Prot. 2005, 83 (B1), 31−37. (5) Trypuć, M.; Stefanowicz, D. I. Solubility in the KVO3 + NH4VO3 + H2O System. J. Chem. Eng. 1997, 42, 1140−1144. (6) Trypuć, M.; Drużyński, S. Investigation of Mutual Solubility in the NH4VO3 - NH4NO3 - H2O system. Ind. Eng. Chem. Res. 2009, 48, 5058−5063. (7) Wróbel, A.; Drużyński, S.; Kiełkowska, U. Solid Liquid Equilibria Studies in the KVO3−KNO3−H2O System in the Temperature Range 293.15−323.15 K. J. Chem. Eng. Data 2017, 62, 3802−3806. ( 8 ) T r y p u ć , M . ; D r u ż y ń s k i , S . S o l u b i l i t y i n t h e NH4NO3+NaNO3+H2O System. Ind. Eng. Chem. Res. 2008, 47, 3767−3770. (9) Kocjan, R. Analytical Chemistry, Classical Quantitative Analysis, Vol. I; PZWL: Warszawa, Poland, 2015. (10) Sáez-Plaza, P.; Michałowski, T.; Navas, J. M.; Asuero, A. G.; Wybraniec, S. An Overview of the Kjeldahl Method of Nitrogen Determination. Part I. Early History, Chemistry of the Procedure, and Titrimetric Finish. Crit. Rev. Anal. Chem. 2013, 43, 178−223. (11) Speight, G. Lange’s Handbook of Chemistry; McGraw-Hill: New York, 2005. (12) Broul, M.; Nývlt, J.; Söhnel, O. Solubility in Inorganic TwoComponent Systems; Prague, Czech Republic, 1981. (13) Söhnel, O.; Novotny, P. Densities of Aqueous Solutions of Inorganic Substances; Elsevier: Amsterdam, The Netherlands, 1985. (14) Korin, E.; Soifer, L. Phase diagram for the system K2Cr2O7 + KNO3 + H2O in the temperature range 10°C to 40°C. J. Chem. Eng. Data 1997, 42, 508−510. (15) Joint Committee on Powder Diffraction Standards. Powder Diffraction File 1976, USA. (16) Dejewska, B. The Characteristics of the Mixed Crystals of the KNO3-NH4NO3-H2O System at 298 K. Cryst. Res. Technol. 2000, 35, 1059−1067. (17) Gałdecki, Z.; Czurak, W.; Fruziński, A.; Dejewska, B. The structure of mixed crystals KNO3. x NH4NO3. Cryst. Res. Technol. 1995, 30, 87−89. (18) Dejewska, B.; Sedzimir, A. X-ray powder diffraction investigations of solid solutions with limited miscibility in the KNO3 NH4NO3 - H2O system at 298 K. Cryst. Res. Technol. 1988, 23, 997− 1004. (19) Dejewska, B. Solid Solutions Crystallising Inorganic Salts in Systems: Water under Phase Equilibrium Conditions; Nicolaus Copernicus University: Toruń, Poland, 1996. (20) Kudryashova, O. S.; Kataev, A. V.; Malinina, L. N. Solubility in the NaNO3-NH4NO3-KNO3-H2O system. Russ. J. Inorg. Chem. 2015, 60 (3), 355−361. (21) Karnaukhov, A. S. The study of the ternary systems NaNO3− NH4NO3−H2O, KNO3−NH4NO3−H2O, RbNO3−NH4NO3−H2O by use of physicochemical analysis at 25 °C. Zh. Obsh. Khim. 1956, 26, 1027−1034. (22) Shenkin, Y. S.; Ruchnova, S. A. Solubility isobar in NH4NO3 KNO3 - H2O system at atmospheric pressure. Zh. Neorg. Khim. 1970, 15, 2512−2518. (23) Morrow, S. I., Campisi, J. J.; Bracuti, A. J. Investigation of propellant and explosive solid solution systems II, X- ray studies; Technical Report ADA055981, 1978. (24) Ando, J.; Smith, J. P.; Siegel, M. R.; Jordan, J. E. Quantitative analysis of mixed fertilizers by X-ray diffraction. J. Agric. Food Chem. 1965, 13, 186−195. (25) Gunawan, R.; Zhang, D. Thermal stability and kinetics of decomposition of ammonium nitrate in the presence of pyrite. J. Hazard. Mater. 2009, 165, 751−758.

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Equilibrium Study in the KNO3+ NH4NO3+ H2O System at

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