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Investigation of phase evolution during the formation of calcium potassium sodium orthophosphate
Materials Chemistry and Physics 78 (2002) 308–312
Material science communication
Investigation of phase evolution during the formation of calcium potassium sodium orthophosphate Niu Jinlong a,b,∗ , Zhang Zhenxi a , Jiang Dazong a a
Institute of Biomedical Engineering, School of Life Science and Technology, Xi’an Jiaotong University, Xi’an 710049, China b Northwest Institute for Nonferrous Metal Research, Xi’an 710016, China Received 14 September 2001; received in revised form 7 March 2002; accepted 6 April 2002
Abstract The formation of calcium potassium sodium orthophosphate (Ca2 KNa(PO4 )2 ) from sodium dihydrogen phosphate dihydrate (NaH2 PO4 · 2H2 O), potassium dihydrogen phosphate (KH2 PO4 ) and calcium carbonate (CaCO3 ) was investigated using differential thermal analysis (DTA), thermal gravimetric analysis (TGA) and powder X-ray diffraction (XRD). DTA showed five distinct thermal events attributed to dehydration of NaH2 PO4 ·2H2 O, dehydration of NaH2 PO4 and KH2 PO4 , the transition of amorphous calcium sodium phosphate to -CaNaPO4 , the decomposition of CaCO3 to CaO, and the crystallization of Ca2 KNa(PO4 )2 . TGA showed 5.6% weight loss between 70 and 170 ◦ C due to the dehydration of NaH2 PO4 ·2H2 O to NaH2 PO4 ; the 7.3% weight loss between 170 and 280 ◦ C is due to the dehydration of NaH2 PO4 and KH2 PO4 ; a continuous weight loss over 280–760 ◦ C due to the reaction of melted alkali metal phosphates with CaCO3 and an approximate 8% weight loss between 670 and 840 ◦ C due to the calcination of residual CaCO3 . XRD analysis, as a function of temperature, supported the evolution of these events and phases, showed the formation of other intermediates: -CaKPO4 and calcium potassium pyrophosphate (CaK2 P2 O7 ). The proposed mechanism of the reaction to Ca2 KNa(PO4 )2 involves the formation and consumption of these intermediates: acid alkali metal pyrophosphate, metaphosphate, CaK2 P2 O7 and -CaNaPO4 . © 2002 Elsevier Science B.V. All rights reserved. Keywords: Resorbable bioceramic; Calcium alkali phosphate; Thermochemical synthesis
1. Introduction Being the main inorganic constituent of hard tissue, calcium phosphates have long been attractive in hard tissue repair [1]. Calcium hydroxyapatite (Ca10 (PO4 )6 (OH)2 , HA) and the more resorbable -tricalcium phosphate (-Ca3 (PO4 )2 , -TCP), are currently recognized as ceramic materials that significantly simulate the mineralogical structure of bone [2,3]. When calcium phosphates ceramics implanted in vivo, these materials are non-toxic, antigenically inactive, noncarcinogenic and bond directly to bone without any intervening connective tissue layer. They show good biological compatibility, safety and osteoconductivity in living tissues; therefore, they are clinically used for hard tissue replacement [4,5]. It is of great interest for both phosphorous chemists and biomedical engineer, as well as material scientists to understand the mechanism of ∗ Corresponding author. Present address: Institute of Biomedical Engineering, School of Life Science and Technology, Xi’an Jiaotong University, Xi’an 710049, China. Tel.: +86-29-2668664. E-mail address: [email protected] (N. Jinlong).
the formation of calcium phosphate ceramics at high temperature. Because of different process, calcium phosphate ceramic would have different physicochemical, mechanical and biological properties. HA and -TCP both are bioactive materials, but their biodegrabilities are different. HA is much less degraded and resorbed by living tissue than -TCP [6]. Though -TCP is biodegradable and resorbable bioceramics, its biodegrability and/or bioresorbability are worth improving for fast repair of bone defect. To improve the degradability of calcium phosphate ceramic with no toxic additive, alkali metal’s phosphate was added to tricalcium phosphate, and a new kind of crystal ceramic was formed with the chemical formula Ca2 KNa(PO4 )2 [7,8]. Because K+ and Na+ are the basic constituent of body fluid and whole blood, the relatively high solubility of Ca2 KNa(PO4 )2 ceramic would make it a promising material for fast healing of bone defect in living body. A solid state reaction was selected to synthesize crystal Ca2 KNa(PO4 )2 , the phase evolution during the formation of Ca2 KNa(PO4 )2 was investigated in this project to understand the mechanism of the reaction path.
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2. Experiments Because of the difference between solubility of calcium and alkali metal phosphates, crystal ceramic Ca2 KNa(PO4 )2 cannot coprecipitate from aqueous system, the only synthesis method may be solid state reaction [7,8]. Analytical purity chemical reagents of calcium carbonate (CaCO3 —CC), potassium dihydrogen phosphate (KH2 PO4 —KHP), sodium dihydrogen phosphate dihydrate (NaH2 PO4 ·2H2 O—NHP) were mixed stoichiometrically with formula Ca2 KNa(PO4 )2 in 2:1:1 molar ratio and milled for 2 h in alumina roller. Each 10 g of the mixture was uniaxially compressed under 50 MPa and placed in alumina crucible, heated in SiC furnace. Samples were removed from the furnace at 400, 600, 800, 900, 1000, 1100, 1200 and 1300 ◦ C after held at the given temperature for half-hour. These samples were removed out from crucible and crushed, grounded immediately in an alumina mortar to powder and kept in a dry container. The sample powder was packed into a special aluminum container used for powder X-ray diffraction (XRD) measurement. Powder XRD was performed at room temperature using a diffractometer (PW1710, Philips) with Cu K␣ radiation λ = 1.5406 Å, generator tension 40 kV and current 30 mA. Monochrometer and Scintillation detector were used. Sample was scanned step by step over the 2θ range of 20–55◦ , using a step size of 0.02◦ and a scan rate of 0.02◦ s−1 . Thermal gravimetric analysis (TGA) (TGA 51 Thermogravimetric Analyzer, TA Instruments) and differential
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thermal analysis (DTA) (DSC 910S Differential Scanning Calorimeter, TA Instruments) were performed on precursor powder over a temperature range of 25–1000 and 25–1300 ◦ C, respectively with a heating rate of 20 ◦ C min−1 .
3. Results and discussion DTA and TGA data during the reaction of CC (CaCO3 ) and KHP (KH2 PO4 ) with NHP (NaH2 PO4 ·2H2 O) to form Ca2 KNa(PO4 )2 are shown as curves DTA and TGA in Fig. 1. The bulk reaction may be shown as 2CaCO3 + KH2 PO4 + NaH2 PO4 · 2H2 O → Ca2 KNa(PO4 )2 + 4H2 O ↑ +2CO2 ↑
(1)
The theoretical total weight loss, due to the escape of all gaseous species, calculated from the reaction (Eq. (1)) is 32.5%. The total weight loss, observed upon heating to 1000 ◦ C, was 30%. Powder XRD was used as a method to detect phase evolution during the reaction. All the samples removed at different temperatures were measured and their XRD patterns were shown in Fig. 2. All detected crystalline phases are labeled (Fig. 2). Powder XRD shows, after rolled in alumina roller for 2 h, the main crystalline phases of the precursor powder were CC (CaCO3 , JCPDS #5–586) and KHP (KH2 PO4 , JCPDS
Fig. 1. Thermal analysis data (DTA and TGA) for the reaction of CaCO3 with KH2 PO4 and NaH2 PO4 ·2H2 O to form Ca2 KNa(PO4 )2 at a heating rate of 20 ◦ C min−1 .
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Fig. 2. XRD patterns of samples showing the phase evolution during the formation of Ca2 KNa(PO4 )2 . Temperatures at which samples were removed are shown. Labels are as follows: CC (CaCO3 ); KHP (KH2 PO4 ); NHP (NaH2 PO4 ·2H2 O); CKP (CaK2 P2 O7 ); CNP (-CaNaPO4 ); CKNP (Ca2 KNa(PO4 )2 ).
#35–807). The diffraction peaks of NHP (NaH2 PO4 ·2H2 O, JCPDS #10–198) were not clearly shown in the XRD pattern of precursor, but there were two weak peaks of sodium dihydrogen phosphate monohydrate (NaH2 PO4 ·H2 O, JCPDS #11–651) of d = 3.37 Å (2θ = 26.42◦ ) and d = 3.16 Å (2θ = 28.18◦ ). There is a possibility that the first strong peak (d = 3.70 Å) of NaH2 PO4 ·2H2 O may be overlapped by the strongest peak of KH2 PO4 (d = 3.72 Å, JCPDS #35–807), as shown in Fig. 2, the observed peak is d = 3.71 Å (2θ = 23.93◦ ). This suggested that NaH2 PO4 ·2H2 O would lose a little crystal water and NaH2 PO4 ·H2 O would form after long time rolling, because of the frictional heat and long time contact of NaH2 PO4 ·2H2 O with CaCO3 . When the mixture was heated to 400 ◦ C, the XRD pattern of the resulting material showed that the peaks of KH2 PO4 and NaH2 PO4 ·2H2 O disappeared, yet crystalline phase of CaCO3 still remained, and a new phase was formed which was determined as calcium potassium pyrophosphate (CaK2 P2 O7 , CKP, JCPDS #22–805). DTA and TGA showed, as comparison with XRD, that a substantial endothermic event and weight loss occurred between 70 and 170 ◦ C, peak value was about 83 ◦ C. This is suggested as the dehydration process of NaH2 PO4 ·2H2 O to NaH2 PO4 . The weight loss is about 5.6%, slightly lower than theoretical value (7.3%). The slightly small weight loss may be due to the little water loss at the beginning of the mix stage. NaH2 PO4 ·2H2 O would lose all its crystal water before 100 ◦ C [9]. The experiment showed that the fast dehydration process began at about 70 ◦ C, the peak was about 83 ◦ C. It
is proposed consequently that NaH2 PO4 ·2H2 O has lost its crystal water completely by 170 ◦ C. The reaction, occurred within this stage, may be described as follows: NaH2 PO4 · 2H2 O → NaH2 PO4 + 2H2 O ↑
(2)
At about 241 ◦ C (between 170 and 280 ◦ C), there was another significant endothermic event and a substantial weight loss (6.7%), this is due to the melting process and the dehydration of KH2 PO4 and NaH2 PO4 . Since the melting points are about 253 and 190 ◦ C, respectively, the mixture of KH2 PO4 and NaH2 PO4 may have a melting point between theirs. It is known that KH2 PO4 and NaH2 PO4 will decompose and dehydrate to acid pyrophosphate when the temperature is higher than their melting points [9]. When the temperature was increased to a much higher degree, the acid pyrophosphate may become to metaphosphate. The melted salt would solidify to glassy state when they were cooled. The glassy phase gives very weak XRD intensity, so that it cannot be detected by XRD. Meanwhile, the melted acid pyrophosphate would react with CaCO3 inevitably, which resulted in the release of carbon dioxide. The changes occurred at this stage may include the following reactions: 2NaH2 PO4 → Na2 H2 P2 O7 + H2 O ↑
(3)
2KH2 PO4 → K2 H2 P2 O7 + H2 O ↑
(4)
K2 H2 P2 O7 +CaCO3 → CaK2 P2 O7 +H2 O ↑ +CO2 ↑
(5)
TGA showed a continuous weight losing (about 8.5%) between 280 and 670 ◦ C. This means that the melted acid
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pyrophosphate reacted with CC continuously, and carbon dioxide and water vapor were released simultaneously. XRD pattern of the sample produced at 400 ◦ C has shown the disappearance of dihydrogen phosphates and the formation of CKP (CaK2 P2 O7 , JCPDS #22–805). XRD pattern of sample produced at 600 ◦ C showed clearly the existence of CaK2 P2 O7 and CaCO3 phases. Besides the two main phases, a new phase was detected as -calcium sodium phosphate (-CaNaPO4 —-CNP, JCPDS #29–1193), because its three strong lines, d = 2.75, 2.70, 2.66 Å (2θ = 32.58, 33.10, 33.71◦ , respectively) were clearly found. The existence of -calcium potassium phosphate (-CaKPO4 , JCPDS #33–1002) may be possible, because its three strong lines d = 2.98, 2.79, 2.06 Å may be overlapped by the three lines of CaK2 P2 O7 (JCPDS #22–805) d = 3.00, 2.83, 2.06 Å, respectively. It can be seen from the XRD pattern that the two peaks, d = 3.04 and 2.83 Å (2θ = 29.40 and 31.60◦ ) are not symmetrical. The right extension of the two peaks may conceal the two lines of -CaKPO4 (d = 2.98 and 2.78 Å). DTA showed an endothermal peak at about 500 ◦ C without substantial weight loss. This may be the phase transition of amorphous calcium sodium and potassium phosphate to -CaNaPO4 and -CaKPO4 . NHP and KHP will melt and dehydrate above 200 ◦ C; a glassy phase of pyrophosphate and/or metaphosphate would be formed. The liquid glassy phosphates covered the reaction products of melt acid pyrophosphate with CaCO3 and an intermediate with molar ratio Ca/P < 1 was formed. With the increase of reaction temperature, the Ca/P of intermediate would increase gradually. CKP (CaK2 P2 O7 ) with Ca/P = 0.5 was first formed, then at about 500 ◦ C, amorphous calcium potassium or -CNP with Ca/P = 1 would produced in a little quantity. The intermediate phases were amorphous, diffraction peaks were very weak, until the amorphous intermediate with Ca/P of 0.5 or 1 changed to crystalline phase. Besides the main crystalline phases of CaCO3 and CaK2 P2 O7 , XRD pattern of sample at 600 ◦ C showed the existence of -CaNaPO4 , supported the discussion above. So the reaction occurred between 280 and 670 ◦ C may include Eq. (5) as well as the following equations:
showed significant amount of CaCO3 , which means CaCO3 would not decomposed completely at 800 ◦ C. -CaNaPO4 and CaK2 P2 O7 were also found as crystalline components of the product. Besides CaCO3 , -CaNaPO4 and CaK2 P2 O7 crystals, there appeared some new peaks that were determined as diffraction lines of Ca2 KNa(PO4 )2 [8], because these peaks appeared at the same position in the XRD pattern of sample produced at higher temperature. The reaction may be as follows:
Na2 H2 P2 O7 + 2CaCO3
The formation of Ca2 KNa(PO4 )2 from the reaction of CC (CaCO3 ), KHP (KH2 PO4 ) and NHP (NaH2 PO4 ·2H2 O) with a molar ratio of 2:1:1 followed a complex chemical reaction path. At the beginning stage of the reaction, NaH2 PO4 ·2H2 O lost its total crystal water and formed NaH2 PO4 . Then, between 170 and 400 ◦ C the acid phosphate of sodium and potassium melted and dehydrated to acid pyrophosphate, and the melted salt reacted with CC to form CKP. The amorphous intermediate of calcium potassium and/or sodium phosphate transformed to -CNP and -calcium potassium phosphate by 600 ◦ C. The residual CC decomposed at about 800 ◦ C, and Ca2 KNa(PO4 )2 began to form. The perfect crystalline phase of Ca2 KNa(PO4 )2 was formed completely by 1300 ◦ C.
→ 2CaNaPO4 (amorphous) + H2 O ↑ +2CO2 ↑ CaNaPO4 (amorphous) → -CaNaPO4
(6) (7)
DTA showed a significant endothermic event between 670 and 840 ◦ C that owed to the decomposition of CC. TGA showed 8% weight loss over the same temperature range. The reaction is shown below: CaCO3 → CaO + CO2 ↑
(8)
Although DTA showed that CaCO3 decomposed at about 780 ◦ C, TGA showed the calcination was not completed until 840 ◦ C. XRD pattern of sample produced at 800 ◦ C
CaK2 P2 O7 + 2CaNaPO4 + CaO → 2Ca2 KNa(PO4 )2
(9)
A slowly endothermic event was shown in DTA between 840 and 1300 ◦ C; it can be considered as the melting process of Ca2 KNa(PO4 )2 at high temperature, because the only crystalline phase was Ca2 KNa(PO4 )2 in the samples produced at 900 ◦ C or higher temperature. The formation of Ca2 KNa(PO4 )2 began at about 800 ◦ C with the decomposition of CaCO3 , but it was a solid state reaction of CaO with CaK2 P2 O7 and -CaNaPO4 ; the lattice structure would not be perfect at the initial stage of the reaction, it is necessary to increase the reaction temperature. The appearance of the sample tablet, produced at different temperatures, was of benefit to the explanation of the reaction process. The white block was formed at 900, 1000, 1100 and 1200 ◦ C, but the tablet produced at 1300 ◦ C was semitransparent crystalline. Hardness and strength of the sample increased with the reaction temperature. It is suggested that the melting degree increased with the reaction temperature. Between 840 and 1000 ◦ C, there was a weight loss of 1.2%. The little weight loss might attribute to the decomposition of CaCO3 that did not calcine completely below 840 ◦ C. The measured total weight loss of the precursor was about 30% between room temperature to 1000 ◦ C; it corresponded to theoretical weight loss of 32.5%. The little difference between the observed weight loss and calculated value may be resulted from the dehydration of NaH2 PO4 ·2H2 O at the initial mixing stage or the dehydration of the precursor stored in the desiccator.
4. Conclusion
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Acknowledgements The project was supported by the Doctorate Foundation of Xi’an Jiaotong University (DFXJU2000-9) and Shaanxi Provincial Science Foundation of China (2000C17). References [1] K. de Groot, Bioceramics consisting of calcium phosphate salts, Biomaterials 1 (1980) 47. [2] E.C. Shors, R.E. Holmes, Porous hydroxyapatite, in: L.L. Hench, J. Wilson (Eds.), An Introduction to Bioceramics, World Scientific, Singapore, 1993, p. 181.
[3] J.B. Park, R.S. Lakes, Biomaterials: An Introduction, Plenum Press, New York, 1992 (Chapter 6). [4] L.L. Hench, Bioceramics, J. Am. Ceram. Soc. 81 (1998) 1705. [5] D.C. Greenspan, Bioactive ceramic implant materials, Curr. Opin. Solid State Mater. Sci. 4 (1999) 389. [6] W. Suchanek, M. Yoshimura, Processing and properties of hydroxyapatite-based biomaterials for use as hard tissue replacements, J. Mater. Res. 13 (1998) 94. [7] G. Berger, R. Gildenhaar, U. Ploska, Rapid resorbable, glassy crystalline materials on the basis of calcium alkali orthophosphates, Biomaterials 16 (1995) 1241. [8] J.L. Niu, Z.X. Zhang, D.Z. Jiang, Preparation structure and solubility of Ca2 KNa(PO4 )2 , J. Mater. Sci. 36 (2001) 3805. [9] D.L. Perry, S.L. Phillips (Eds.), Handbook of Inorganic Compounds, World Publ. Co., Beijing, 1998.
Material science communication
Investigation of phase evolution during the formation of calcium potassium sodium orthophosphate Niu Jinlong a,b,∗ , Zhang Zhenxi a , Jiang Dazong a a
Institute of Biomedical Engineering, School of Life Science and Technology, Xi’an Jiaotong University, Xi’an 710049, China b Northwest Institute for Nonferrous Metal Research, Xi’an 710016, China Received 14 September 2001; received in revised form 7 March 2002; accepted 6 April 2002
Abstract The formation of calcium potassium sodium orthophosphate (Ca2 KNa(PO4 )2 ) from sodium dihydrogen phosphate dihydrate (NaH2 PO4 · 2H2 O), potassium dihydrogen phosphate (KH2 PO4 ) and calcium carbonate (CaCO3 ) was investigated using differential thermal analysis (DTA), thermal gravimetric analysis (TGA) and powder X-ray diffraction (XRD). DTA showed five distinct thermal events attributed to dehydration of NaH2 PO4 ·2H2 O, dehydration of NaH2 PO4 and KH2 PO4 , the transition of amorphous calcium sodium phosphate to -CaNaPO4 , the decomposition of CaCO3 to CaO, and the crystallization of Ca2 KNa(PO4 )2 . TGA showed 5.6% weight loss between 70 and 170 ◦ C due to the dehydration of NaH2 PO4 ·2H2 O to NaH2 PO4 ; the 7.3% weight loss between 170 and 280 ◦ C is due to the dehydration of NaH2 PO4 and KH2 PO4 ; a continuous weight loss over 280–760 ◦ C due to the reaction of melted alkali metal phosphates with CaCO3 and an approximate 8% weight loss between 670 and 840 ◦ C due to the calcination of residual CaCO3 . XRD analysis, as a function of temperature, supported the evolution of these events and phases, showed the formation of other intermediates: -CaKPO4 and calcium potassium pyrophosphate (CaK2 P2 O7 ). The proposed mechanism of the reaction to Ca2 KNa(PO4 )2 involves the formation and consumption of these intermediates: acid alkali metal pyrophosphate, metaphosphate, CaK2 P2 O7 and -CaNaPO4 . © 2002 Elsevier Science B.V. All rights reserved. Keywords: Resorbable bioceramic; Calcium alkali phosphate; Thermochemical synthesis
1. Introduction Being the main inorganic constituent of hard tissue, calcium phosphates have long been attractive in hard tissue repair [1]. Calcium hydroxyapatite (Ca10 (PO4 )6 (OH)2 , HA) and the more resorbable -tricalcium phosphate (-Ca3 (PO4 )2 , -TCP), are currently recognized as ceramic materials that significantly simulate the mineralogical structure of bone [2,3]. When calcium phosphates ceramics implanted in vivo, these materials are non-toxic, antigenically inactive, noncarcinogenic and bond directly to bone without any intervening connective tissue layer. They show good biological compatibility, safety and osteoconductivity in living tissues; therefore, they are clinically used for hard tissue replacement [4,5]. It is of great interest for both phosphorous chemists and biomedical engineer, as well as material scientists to understand the mechanism of ∗ Corresponding author. Present address: Institute of Biomedical Engineering, School of Life Science and Technology, Xi’an Jiaotong University, Xi’an 710049, China. Tel.: +86-29-2668664. E-mail address: [email protected] (N. Jinlong).
the formation of calcium phosphate ceramics at high temperature. Because of different process, calcium phosphate ceramic would have different physicochemical, mechanical and biological properties. HA and -TCP both are bioactive materials, but their biodegrabilities are different. HA is much less degraded and resorbed by living tissue than -TCP [6]. Though -TCP is biodegradable and resorbable bioceramics, its biodegrability and/or bioresorbability are worth improving for fast repair of bone defect. To improve the degradability of calcium phosphate ceramic with no toxic additive, alkali metal’s phosphate was added to tricalcium phosphate, and a new kind of crystal ceramic was formed with the chemical formula Ca2 KNa(PO4 )2 [7,8]. Because K+ and Na+ are the basic constituent of body fluid and whole blood, the relatively high solubility of Ca2 KNa(PO4 )2 ceramic would make it a promising material for fast healing of bone defect in living body. A solid state reaction was selected to synthesize crystal Ca2 KNa(PO4 )2 , the phase evolution during the formation of Ca2 KNa(PO4 )2 was investigated in this project to understand the mechanism of the reaction path.
0254-0584/02/$ – see front matter © 2002 Elsevier Science B.V. All rights reserved. PII: S 0 2 5 4 - 0 5 8 4 ( 0 2 ) 0 0 2 0 0 - 6
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2. Experiments Because of the difference between solubility of calcium and alkali metal phosphates, crystal ceramic Ca2 KNa(PO4 )2 cannot coprecipitate from aqueous system, the only synthesis method may be solid state reaction [7,8]. Analytical purity chemical reagents of calcium carbonate (CaCO3 —CC), potassium dihydrogen phosphate (KH2 PO4 —KHP), sodium dihydrogen phosphate dihydrate (NaH2 PO4 ·2H2 O—NHP) were mixed stoichiometrically with formula Ca2 KNa(PO4 )2 in 2:1:1 molar ratio and milled for 2 h in alumina roller. Each 10 g of the mixture was uniaxially compressed under 50 MPa and placed in alumina crucible, heated in SiC furnace. Samples were removed from the furnace at 400, 600, 800, 900, 1000, 1100, 1200 and 1300 ◦ C after held at the given temperature for half-hour. These samples were removed out from crucible and crushed, grounded immediately in an alumina mortar to powder and kept in a dry container. The sample powder was packed into a special aluminum container used for powder X-ray diffraction (XRD) measurement. Powder XRD was performed at room temperature using a diffractometer (PW1710, Philips) with Cu K␣ radiation λ = 1.5406 Å, generator tension 40 kV and current 30 mA. Monochrometer and Scintillation detector were used. Sample was scanned step by step over the 2θ range of 20–55◦ , using a step size of 0.02◦ and a scan rate of 0.02◦ s−1 . Thermal gravimetric analysis (TGA) (TGA 51 Thermogravimetric Analyzer, TA Instruments) and differential
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thermal analysis (DTA) (DSC 910S Differential Scanning Calorimeter, TA Instruments) were performed on precursor powder over a temperature range of 25–1000 and 25–1300 ◦ C, respectively with a heating rate of 20 ◦ C min−1 .
3. Results and discussion DTA and TGA data during the reaction of CC (CaCO3 ) and KHP (KH2 PO4 ) with NHP (NaH2 PO4 ·2H2 O) to form Ca2 KNa(PO4 )2 are shown as curves DTA and TGA in Fig. 1. The bulk reaction may be shown as 2CaCO3 + KH2 PO4 + NaH2 PO4 · 2H2 O → Ca2 KNa(PO4 )2 + 4H2 O ↑ +2CO2 ↑
(1)
The theoretical total weight loss, due to the escape of all gaseous species, calculated from the reaction (Eq. (1)) is 32.5%. The total weight loss, observed upon heating to 1000 ◦ C, was 30%. Powder XRD was used as a method to detect phase evolution during the reaction. All the samples removed at different temperatures were measured and their XRD patterns were shown in Fig. 2. All detected crystalline phases are labeled (Fig. 2). Powder XRD shows, after rolled in alumina roller for 2 h, the main crystalline phases of the precursor powder were CC (CaCO3 , JCPDS #5–586) and KHP (KH2 PO4 , JCPDS
Fig. 1. Thermal analysis data (DTA and TGA) for the reaction of CaCO3 with KH2 PO4 and NaH2 PO4 ·2H2 O to form Ca2 KNa(PO4 )2 at a heating rate of 20 ◦ C min−1 .
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Fig. 2. XRD patterns of samples showing the phase evolution during the formation of Ca2 KNa(PO4 )2 . Temperatures at which samples were removed are shown. Labels are as follows: CC (CaCO3 ); KHP (KH2 PO4 ); NHP (NaH2 PO4 ·2H2 O); CKP (CaK2 P2 O7 ); CNP (-CaNaPO4 ); CKNP (Ca2 KNa(PO4 )2 ).
#35–807). The diffraction peaks of NHP (NaH2 PO4 ·2H2 O, JCPDS #10–198) were not clearly shown in the XRD pattern of precursor, but there were two weak peaks of sodium dihydrogen phosphate monohydrate (NaH2 PO4 ·H2 O, JCPDS #11–651) of d = 3.37 Å (2θ = 26.42◦ ) and d = 3.16 Å (2θ = 28.18◦ ). There is a possibility that the first strong peak (d = 3.70 Å) of NaH2 PO4 ·2H2 O may be overlapped by the strongest peak of KH2 PO4 (d = 3.72 Å, JCPDS #35–807), as shown in Fig. 2, the observed peak is d = 3.71 Å (2θ = 23.93◦ ). This suggested that NaH2 PO4 ·2H2 O would lose a little crystal water and NaH2 PO4 ·H2 O would form after long time rolling, because of the frictional heat and long time contact of NaH2 PO4 ·2H2 O with CaCO3 . When the mixture was heated to 400 ◦ C, the XRD pattern of the resulting material showed that the peaks of KH2 PO4 and NaH2 PO4 ·2H2 O disappeared, yet crystalline phase of CaCO3 still remained, and a new phase was formed which was determined as calcium potassium pyrophosphate (CaK2 P2 O7 , CKP, JCPDS #22–805). DTA and TGA showed, as comparison with XRD, that a substantial endothermic event and weight loss occurred between 70 and 170 ◦ C, peak value was about 83 ◦ C. This is suggested as the dehydration process of NaH2 PO4 ·2H2 O to NaH2 PO4 . The weight loss is about 5.6%, slightly lower than theoretical value (7.3%). The slightly small weight loss may be due to the little water loss at the beginning of the mix stage. NaH2 PO4 ·2H2 O would lose all its crystal water before 100 ◦ C [9]. The experiment showed that the fast dehydration process began at about 70 ◦ C, the peak was about 83 ◦ C. It
is proposed consequently that NaH2 PO4 ·2H2 O has lost its crystal water completely by 170 ◦ C. The reaction, occurred within this stage, may be described as follows: NaH2 PO4 · 2H2 O → NaH2 PO4 + 2H2 O ↑
(2)
At about 241 ◦ C (between 170 and 280 ◦ C), there was another significant endothermic event and a substantial weight loss (6.7%), this is due to the melting process and the dehydration of KH2 PO4 and NaH2 PO4 . Since the melting points are about 253 and 190 ◦ C, respectively, the mixture of KH2 PO4 and NaH2 PO4 may have a melting point between theirs. It is known that KH2 PO4 and NaH2 PO4 will decompose and dehydrate to acid pyrophosphate when the temperature is higher than their melting points [9]. When the temperature was increased to a much higher degree, the acid pyrophosphate may become to metaphosphate. The melted salt would solidify to glassy state when they were cooled. The glassy phase gives very weak XRD intensity, so that it cannot be detected by XRD. Meanwhile, the melted acid pyrophosphate would react with CaCO3 inevitably, which resulted in the release of carbon dioxide. The changes occurred at this stage may include the following reactions: 2NaH2 PO4 → Na2 H2 P2 O7 + H2 O ↑
(3)
2KH2 PO4 → K2 H2 P2 O7 + H2 O ↑
(4)
K2 H2 P2 O7 +CaCO3 → CaK2 P2 O7 +H2 O ↑ +CO2 ↑
(5)
TGA showed a continuous weight losing (about 8.5%) between 280 and 670 ◦ C. This means that the melted acid
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pyrophosphate reacted with CC continuously, and carbon dioxide and water vapor were released simultaneously. XRD pattern of the sample produced at 400 ◦ C has shown the disappearance of dihydrogen phosphates and the formation of CKP (CaK2 P2 O7 , JCPDS #22–805). XRD pattern of sample produced at 600 ◦ C showed clearly the existence of CaK2 P2 O7 and CaCO3 phases. Besides the two main phases, a new phase was detected as -calcium sodium phosphate (-CaNaPO4 —-CNP, JCPDS #29–1193), because its three strong lines, d = 2.75, 2.70, 2.66 Å (2θ = 32.58, 33.10, 33.71◦ , respectively) were clearly found. The existence of -calcium potassium phosphate (-CaKPO4 , JCPDS #33–1002) may be possible, because its three strong lines d = 2.98, 2.79, 2.06 Å may be overlapped by the three lines of CaK2 P2 O7 (JCPDS #22–805) d = 3.00, 2.83, 2.06 Å, respectively. It can be seen from the XRD pattern that the two peaks, d = 3.04 and 2.83 Å (2θ = 29.40 and 31.60◦ ) are not symmetrical. The right extension of the two peaks may conceal the two lines of -CaKPO4 (d = 2.98 and 2.78 Å). DTA showed an endothermal peak at about 500 ◦ C without substantial weight loss. This may be the phase transition of amorphous calcium sodium and potassium phosphate to -CaNaPO4 and -CaKPO4 . NHP and KHP will melt and dehydrate above 200 ◦ C; a glassy phase of pyrophosphate and/or metaphosphate would be formed. The liquid glassy phosphates covered the reaction products of melt acid pyrophosphate with CaCO3 and an intermediate with molar ratio Ca/P < 1 was formed. With the increase of reaction temperature, the Ca/P of intermediate would increase gradually. CKP (CaK2 P2 O7 ) with Ca/P = 0.5 was first formed, then at about 500 ◦ C, amorphous calcium potassium or -CNP with Ca/P = 1 would produced in a little quantity. The intermediate phases were amorphous, diffraction peaks were very weak, until the amorphous intermediate with Ca/P of 0.5 or 1 changed to crystalline phase. Besides the main crystalline phases of CaCO3 and CaK2 P2 O7 , XRD pattern of sample at 600 ◦ C showed the existence of -CaNaPO4 , supported the discussion above. So the reaction occurred between 280 and 670 ◦ C may include Eq. (5) as well as the following equations:
showed significant amount of CaCO3 , which means CaCO3 would not decomposed completely at 800 ◦ C. -CaNaPO4 and CaK2 P2 O7 were also found as crystalline components of the product. Besides CaCO3 , -CaNaPO4 and CaK2 P2 O7 crystals, there appeared some new peaks that were determined as diffraction lines of Ca2 KNa(PO4 )2 [8], because these peaks appeared at the same position in the XRD pattern of sample produced at higher temperature. The reaction may be as follows:
Na2 H2 P2 O7 + 2CaCO3
The formation of Ca2 KNa(PO4 )2 from the reaction of CC (CaCO3 ), KHP (KH2 PO4 ) and NHP (NaH2 PO4 ·2H2 O) with a molar ratio of 2:1:1 followed a complex chemical reaction path. At the beginning stage of the reaction, NaH2 PO4 ·2H2 O lost its total crystal water and formed NaH2 PO4 . Then, between 170 and 400 ◦ C the acid phosphate of sodium and potassium melted and dehydrated to acid pyrophosphate, and the melted salt reacted with CC to form CKP. The amorphous intermediate of calcium potassium and/or sodium phosphate transformed to -CNP and -calcium potassium phosphate by 600 ◦ C. The residual CC decomposed at about 800 ◦ C, and Ca2 KNa(PO4 )2 began to form. The perfect crystalline phase of Ca2 KNa(PO4 )2 was formed completely by 1300 ◦ C.
→ 2CaNaPO4 (amorphous) + H2 O ↑ +2CO2 ↑ CaNaPO4 (amorphous) → -CaNaPO4
(6) (7)
DTA showed a significant endothermic event between 670 and 840 ◦ C that owed to the decomposition of CC. TGA showed 8% weight loss over the same temperature range. The reaction is shown below: CaCO3 → CaO + CO2 ↑
(8)
Although DTA showed that CaCO3 decomposed at about 780 ◦ C, TGA showed the calcination was not completed until 840 ◦ C. XRD pattern of sample produced at 800 ◦ C
CaK2 P2 O7 + 2CaNaPO4 + CaO → 2Ca2 KNa(PO4 )2
(9)
A slowly endothermic event was shown in DTA between 840 and 1300 ◦ C; it can be considered as the melting process of Ca2 KNa(PO4 )2 at high temperature, because the only crystalline phase was Ca2 KNa(PO4 )2 in the samples produced at 900 ◦ C or higher temperature. The formation of Ca2 KNa(PO4 )2 began at about 800 ◦ C with the decomposition of CaCO3 , but it was a solid state reaction of CaO with CaK2 P2 O7 and -CaNaPO4 ; the lattice structure would not be perfect at the initial stage of the reaction, it is necessary to increase the reaction temperature. The appearance of the sample tablet, produced at different temperatures, was of benefit to the explanation of the reaction process. The white block was formed at 900, 1000, 1100 and 1200 ◦ C, but the tablet produced at 1300 ◦ C was semitransparent crystalline. Hardness and strength of the sample increased with the reaction temperature. It is suggested that the melting degree increased with the reaction temperature. Between 840 and 1000 ◦ C, there was a weight loss of 1.2%. The little weight loss might attribute to the decomposition of CaCO3 that did not calcine completely below 840 ◦ C. The measured total weight loss of the precursor was about 30% between room temperature to 1000 ◦ C; it corresponded to theoretical weight loss of 32.5%. The little difference between the observed weight loss and calculated value may be resulted from the dehydration of NaH2 PO4 ·2H2 O at the initial mixing stage or the dehydration of the precursor stored in the desiccator.
4. Conclusion
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Acknowledgements The project was supported by the Doctorate Foundation of Xi’an Jiaotong University (DFXJU2000-9) and Shaanxi Provincial Science Foundation of China (2000C17). References [1] K. de Groot, Bioceramics consisting of calcium phosphate salts, Biomaterials 1 (1980) 47. [2] E.C. Shors, R.E. Holmes, Porous hydroxyapatite, in: L.L. Hench, J. Wilson (Eds.), An Introduction to Bioceramics, World Scientific, Singapore, 1993, p. 181.
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