Continuous Precipitation of Nickel Hydroxide by Addition of

with a peristaltic pump (NRP-75, Nissin Rika Co.). Suspension temperature is measured by a thermometer and a suspended solution is agitated by a 3-bla...
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Continuous Precipitation of Nickel Hydroxide by Addition of Ammonium Ions Kunio Funakoshi,*,# Shogo Yoshizawa, and Masakuni Matsuoka

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Department of Chemical Engineering Tokyo University of Agriculture and Technology (TUAT) Koganei, Tokyo, 184-8588, Japan ABSTRACT: Spherical and dense nickel hydroxide, Ni(OH)2, particles are produced by addition of ammonia or ammonium ions, NH4+, using a continuous precipitation. In this study, the nickel ion, Ni2+, concentrations and the NH4+ concentrations in a suspended solution for the continuous precipitation of Ni(OH)2 are examined, and the formation behaviors of spherical and dense Ni(OH)2 particles are discussed by the relations between the Ni2+ and NH4+ concentrations. When NH4+ was supplied into a suspended solution during the continuous precipitation of Ni(OH)2, Ni2+ was present in a steady state and the Ni2+ concentrations changed with the NH4+ concentrations in a feed solution, with the solution pH and the residence times of suspended solutions in a crystallizer. The ratios of the NH4+ concentrations to the Ni2+ in a solution at steady state were different from the coordination number of NH3 in a nickel ammine complex. Formation of Ni(OH)2 precipitates when adding NH4+ is discussed by the relations between the Ni2+ and the NH4+ concentrations in suspended solution for continuous and batch operations.

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INTRODUCTION Nickel hydroxide, Ni(OH)2, is a type of cathode active material for a nickel−hydrogen secondary battery. Ni(OH)2 precipitates are deposited by mixing of nickel sulfate, NiSO4, and sodium hydroxide, NaOH, aqueous solutions NiSO4 + 2NaOH → Ni(OH)2 ↓ + Na 2SO4

suspended solution the formation behaviors of Ni(OH)2 precipitates are not revealed. The purposes of the present study are to clarify the relations between operating conditions and the properties of product particles or the Ni2+ concentrations and to discuss formation behaviors of Ni(OH)2 precipitates, concerning a reaction of nickel ammine complexes.

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(1)

The Ni(OH)2 precipitates obtained by the reaction of a NiSO4 aqueous solution with a NaOH solution are always amorphous and porous, resulting in lower apparent density.1 On the contrary, when aqueous ammonia, NH3, is added into a suspension during a continuous precipitation, Ni(OH) 2 precipitates become spherical and dense. In this case, nickel ions, Ni2+, would bond with NH3 and a nickel ammine complex would be formed. Then, nickel ammine complexes would be decomposed and Ni(OH)2 precipitates would be formed by reaction with hydroxide ions, OH−, in a solution. Effects of operating conditions on the properties of products such as mean particle size, particle size distributions, PSDs, and apparent density are experimentally examined. When the amount of ammonium ion, NH4+, added into a suspended solution increased, mean particle size became larger,2 PSDs became broader,3 tapping density became higher, and nucleation rates of Ni(OH)2 particles were reduced.4 In addition, at high pH values both particle and crystallite sizes became smaller,5,6 probably because of the enhancement of nucleation and deceleration of growth.2 Moreover, it is reported that operating temperature and agitation speeds also influenced surface shapes and tapping densities of product particles.3,4,6,7,8 Although changes in mean particle size and PSDs of Ni(OH)2 with operating conditions for continuous precipitations were indicated, the Ni2+ concentrations in suspended solutions are not pointed out. Moreover, when NH4+ exists in an Ni(OH)2 © 2016 American Chemical Society

EXPERIMENTAL SECTION

Experimental Apparatus. Schematic drawing of an experimental apparatus is illustrated in Figure 1. A 170 mL jacketed crystallizer made of glass is heated by hot water supplied from a thermostat bath. Feed solutions are provided by diaphragm pumps (DME 2, Grundfos A/S) and pH in a suspended solution is kept at a given value using a pH

Figure 1. Experimental apparatus for continuous precipitation of Ni(OH)2. Received: July 8, 2015 Revised: January 8, 2016 Published: February 5, 2016 1824

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Figure 2. Appearances of product particles at (a) steady state and (b) t = 1 min for (NH4)2SO4 concentration = 0.588 mol/L and pH = 10, and (c) at steady state for pH = 12. controller (NPH-669NDE, Nissin Rika Co.) with a peristaltic pump (NRP-75, Nissin Rika Co.). Suspension temperature is measured by a thermometer and a suspended solution is agitated by a 3-blade marine propeller type impeller with a motor. Experimental Procedures. 100 mL of 10−4 mol/L NaOH aqueous solution was poured into a crystallizer and agitated at 500 rpm and the solution temperature was kept at 50 °C. Ammonium sulfate, (NH4)2SO4, and 1.70 mol/L NiSO4 aqueous solutions were injected into a NaOH aqueous solution at the same flow rates by pumps. 5.00 mol/L NaOH aqueous solution was added into a crystallizer in order to keep the pH in the suspended solution constant. For the case in which the NH4+ concentration in a feed solution is 0 mol/L, distilled water was injected into a NaOH aqueous solution, instead of an (NH4)2SO4 aqueous solution. Suspended solution was removed by a syringe at a given time interval to keep suspension volume in a crystallizer from 120 to 132 mL during continuous precipitation of Ni(OH)2. At a given time, the removed suspension was filtered by a membrane filter. Product particles were observed by a scanning electron microscope, SEM (JSM5310, JEOL Ltd.), and were analyzed by a powder X-ray diffractometer, XRD (Ultima IV, Rigaku Co.). The Ni2+ and the NH4+ concentrations in filtered solutions were measured by the method in the following section. In this study, the (NH4)2SO4 concentrations in a feed solution changed from 0 to 0.588 mol/L, pH = 10−12, and the feeding rates of NiSO4 and (NH4)2SO4 aqueous solution also changed from 1.49 × 10−3 to 7.46 × 10−4 L/min., respectively, because the residence time in the solution changed from 30 to 60 min. Measurements of Ni2+ and NH4+ Concentrations in a Suspended Solution. Ni2+ concentrations in a suspended solution were measured by dimethylglyoxime absorptiometry. 10 mL of filtrated solution and 10 mL iodine−potassium iodide mixed solution were blended in a 100 mL volumetric flask, after which 20 mL aqueous ammonia and 3 mL dimethylglyoxime solution were added into the flask, and then this mixture was diluted with distilled water. The solution was placed into a quartz cell and its absorbance spectrum at 440 nm was measured using an absorptiometer (UV-1200, Shimazu Co.). The Ni2+ concentrations were determined by a calibration curve produced previously. NH4+ concentrations were measured by indophenol blue absorptiometry. The 10 mL filtered solution, 20 mL sodium phenate, and 10 mL sodium hypochlorite were blended in a 100 mL volumetric flask, and then the mixture was diluted with distilled water. The solution was placed in a quartz cell and its absorbance spectrum at 630 nm was measured by an absorptiometer. The NH4+ concentrations were determined by a calibration curve.

A typical powder X-ray diffraction (XRD) pattern of product particles obtained by continuous precipitation of Ni(OH)2 is shown in Figure 3. It is found that these particles are β-type

Figure 3. Powder X-ray diffraction pattern obtained (NH4)2SO4 concentration = 0.588 mol/L and pH = 10.

Ni(OH)2.8 Using Scherrer’s equation, crystallite size of product particles, D, is defined as D=

Kλ β cos θ

(2)

where K, λ, and β are shape factor, wavelength, and half width of the XRD peaks, respectively. Crystallite size of product particles obtained by a series of continuous precipitation of Ni(OH)2 is several tens of nanometers and is similar to the thickness of the “plate-like particles”. Ni2+ Concentrations in Suspended Solutions. Figure 4 shows the changes in Ni2+ concentration with experimental

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RESULTS AND DISCUSSION Crystallographic Properties of Product Particles. Appearances of Ni(OH)2 precipitates obtained for various experimental durations and solution pH are indicated in Figure 2. At the initial stages of a continuous precipitation (t = 1 min), very small spherical crystals (Figure 2a) were formed, and then they changed to “plate-like particles” of which size is less than 100 nm. Their plate-like particles changed to rosaceous agglomerates with size of a few tens of micrometers (Figure 2b). On the contrary, Ni(OH)2 precipitates at pH = 12 were agglomerates constituted by cubic particles, as shown in Figure 2c.

Figure 4. Ni2+ concentrations for various (NH4)2SO4 concentrations in feed solution for case of pH = 10 and τ = 30 min.

duration for various (NH4)2SO4 concentrations of a feed solution. For (NH4)2SO4 concentration of 0 mol/L, the Ni2+ concentrations were almost 0 mol/L during the experiment. Therefore, immediately after the NiSO4 feed solution was added into the suspended solution, Ni2+ reacted with OH− and then 1825

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Ni(OH)2 was deposited. When the (NH4)2SO4 solution was added into a solution, the Ni2+ concentrations increased with experimental duration at the initial stages of experiment, and then they became constant. The higher the (NH4 ) 2SO4 concentration was in a feed solution, the higher the Ni2+ concentration in a suspended solution at steady states. The Ni2+ concentrations at steady state were lower than the higher pH and were almost 0 mol/L for pH = 12, as shown in Figure 5. Changes in Ni2+ concentrations with the residence time

Figure 7. Experimental and calculated results of Ni2+ concentration changes when (NH4)2SO4 concentration in feed solution were 0.588 and 0.294 mol/L and pH = 10.

suspended solution is v [L/min]. The mass balance of NH4+ for a differential time, Δt [min], is indicated by the following C(t )V + C0v0Δt = C(t + Δt )V + C(t + Δt )vΔt

(3)

Equation 3 is modified to eq 3′ C v − C(t + Δt )v C(t + Δt ) − C(t ) = o0 Δt V

Figure 5. Relations between pH in solution and Ni2+ concentrations when (NH4)2SO4 concentration in feed solution was 0.588 mol/L and τ = 30 min.

(3′)

In all experiments, the volume of a suspended solution in a crystallizer was 0.12 L. For (NH4)2SO4 feed concentration of 0.588 mol/L and residence time, τ, of 30 min, C0 became 1.176 mol/L, v0 = 1.49 × 10−3 L/min, and v = 4.00 × 10−3 L/min, respectively. Using eq 3′ and these values, the changes in NH4+ concentrations with experimental duration were calculated. Changes in the calculated NH4+ concentrations with experimental duration for cases in which (NH4)2SO4 concentrations in feed solutions were 0.588 and 0.294 mol/L are also indicated in Figure 7 by solid and broken lines, respectively. The calculated NH4+ concentrations correlate well with the experimental results. Therefore, it is considered that changes in the calculated NH4+ concentrations when the (NH4)2SO4 concentration is 0.147 mol/L, as shown in Figure 7 by a dotted line, would agree with the actual experimental results. The calculated NH4+ concentrations at steady state, C*, are 0.439, 0.219, and 0.110 when the (NH4)2SO4 concentrations of a feed solution are 0.577, 0.294, and 0.147 [mol/L], respectively. Figure 8 shows the relations between NH4+ concentration and experimental duration when the residence times were 30 and 45

are indicated in Figure 6. The longer the residence time of suspended solution was, the lower the Ni2+ concentrations in suspended solutions.

Figure 6. Changes in Ni2+ concentrations with residence times when (NH4)2SO4 concentration in feed solution was 0.588 mol/L and pH = 10.

NH4+ Concentrations in Suspended Solutions. Changes in NH4+ concentration with experimental duration for the cases in which (NH4)2SO4 concentrations in feed solution were 0.588 and 0.294 mol/L and pH = 10 are indicated in Figure 7. The NH4+ concentrations increased with experimental durations at the initial stages of the experiments, and then they became constant. The NH4+ concentrations in a suspended solution at steady states were higher for the higher (NH4)2SO4 concentrations in a feed solution. Changes in the NH4+ concentrations with experimental durations during continuous precipitations of Ni(OH)2 are calculated by the use of the mass balance equation of NH4+ in a suspended solution. (NH4)2SO4 aqueous solution, of which concentration is C0 [mol/L], is fed at v0 [L/min] into the suspended solution of Ni(OH)2. The volume of a suspended solution in a crystallizer is V [L] and the NH4+ concentration in a suspended solution at time t is C(t) [mol/L]. The removal rate of

Figure 8. Relations between NH4+ concentration and experimental duration for various residence times.

min. The changes in NH4+ concentration with experimental duration hardly influences the residence time. Relations between the Ni2+ and NH4+ concentrations at steady states and pH = 10 for continuous precipitations of Ni(OH)2 are shown in Figure 9 and are indicated by the following equation. 1826

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and the pH in the suspended solution kept constant by addition of an NaOH aqueous solution. Changes in Ni2+ and NH4+ concentrations in a suspended solution with experimental duration are shown in Figure 11. Both the Ni2+ and the NH4+ concentrations in a suspended solution became 0 mol/L, but the duration until the Ni2+ concentration became 0 mol/L was shorter than for NH4+.

(4)

Figure 9. Relations between Ni2+ and NH4+ concentrations at steady state.

Figure 11. Changes in Ni2+ and NH4+ concentrations with time after feeding of NiSO4 and (NH4)2SO4 solution were stopped.

Using eq 4 and the changes in the NH4+ concentration in a suspended solution with experimental duration during continuous precipitation, the Ni2+ concentrations are calculated. Changes in the “calculated” Ni2+ concentration with experimental duration agree well with the experimental results, as shown in Figure 10.

The Ni2+ concentrations in suspended solutions at steady state for a continuous precipitation of Ni(OH)2 is higher than the solubility of Ni2+ at a given NH4+ concentration and pH. When NH4+ and/or their complexes existed in suspended solution, Ni2+ is able to exist at comparatively high concentration levels. Ni2+ also forms nickel hydroxide complexes with hydroxide ions, OH−. Moreover, nickel hydroxide complexes would exist in nickel hydroxide precipitates and/or in a solution. Therefore, NH4+, nickel ammine complexes, and/or nickel hydroxide ones in suspended solutions suppressed the crystallization of Ni(OH)2.

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CONCLUSIONS Continuous precipitation of Ni(OH)2 was conducted by mixing aqueous solutions of NiSO4 and NaOH in the presence and the absence of NH4+, and the formation behavior of Ni(OH)2 precipitates was discussed by the relations between the Ni2+ and NH4+ concentrations in suspended solutions at various operating conditions. The appearance and crystallite size of product particles changed with pH or NH4+ concentrations in a suspended solution. Ni2+ concentrations in a suspended solution increased to constant values which changed with the operating conditions. However, in the absence of NH4 +, the Ni2+ concentration in a suspended solution was almost 0 mol/L. NH4+ concentration in a suspended solution for a continuous precipitation of Ni(OH)2 changed with experimental duration but was independent of the pH in a suspended solution. Measured NH4+ concentrations for continuous precipitation of Ni(OH)2 agreed with the calculated values estimated by a mass balance. The ratio of the NH4+ concentrations to the Ni2+ ones at steady state for continuous precipitation of Ni(OH)2 is different from the coordination number of NH3 in nickel ammine complexes. When feeding of NiSO4 and (NH4)2SO4 aqueous solutions was stopped after the attainment of steady state for continuous precipitation of Ni(OH)2, both Ni2+ and NH4+ concentrations became 0 mol/L. The durations in which the Ni2+ concentrations became 0 mol/L, however, were shorter than those for the NH4+ concentrations. The Ni2+ concentrations in suspended solutions at steady state for a continuous

Figure 10. Comparison of dissolved Ni2+ concentration for continuous precipitation of Ni(OH)2 between calculated result and experimental result.

It is said that Ni2+ forms a nickel ammine complex with NH3 and the coordination number of NH3 in a nickel ammine complex is 6 or 4.9 In this study, it is difficult to approximate the relations between Ni2+ and NH4+ concentrations when pH = 10 and τ = 30 min with a linear slope of 6 or 4, as shown in Figure 9. As also shown in Figure 9, Ni2+ which does not form nickel ammine complexes would exist in a suspended solution. On the contrary, it is said that the color of the nickel ammine complex solution is blue or pansy when the number of NH3 ligands forming nickel ammine complexes is 6 or 4, respectively. In this study, the color of the filtered solution was blue except for the case in which the (NH4)2SO4 concentration in a feed solution was 0 mol/L or the pH of the suspended solution was 12. Some NH4+ existed in a suspended solution bound with Ni2+ and would form the nickel ammine complex. The others existed as NH4+ in the suspended solutions. After continuous precipitation of Ni(OH)2 attained a steady state, feeding of the NiSO4 and (NH4)2SO4 solutions stopped 1827

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precipitation of Ni(OH)2 are higher than the solubility of Ni2+ in a solution at a given NH4+ concentration and pH.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Present Address #

National Institute of Technology, Suzuka Collage (NIT, Suzuka Collage) Suzuka, Mie, 510−0284, Japan Notes

The authors declare no competing financial interest.

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GLOSSARY CNH4+0, NH4+ concentration in (NH4)2SO4 feed solution [mol/L] CNH4+(t), NH4+ concentration at time t [mol/L] CNH4+, NH4+ concentration in suspended solution at steady state [mol/L] CNi2+, Ni2+ concentration in suspended solution at steady state[mol/L] D, crystallite size [m] K, shape factor t, time [s] V, volume of suspended solution [L] v0, feeding rate of (NH4)2SO4 aqueous solution [L/s] v, removal rate of suspended solution [L/s] β, half width of XRD peak [°] λ, wavelength of XRD [Å] θ, angle [°]

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REFERENCES

(1) Ikeda, K.; Iwakura, T.; Matsuda, Y. All about nickel-hydrogen secondary batteries; NTS Inc.: Tokyo, 2001; pp 91−93. (2) Kaitani, H.; Imari, A.; Yamashita, K. Electrochemistry 1995, 63, 752−758. (3) Tsuge, H.; Sakamoto, S. 16th Int. Symp. on Ind. Cryst. 2005, 307− 312. (4) Ito, T.; Tsuge, H.; et al. 14th Int. Symp. on Ind. Cryst. 1999, 143− 150. (5) Song, Q.; Tang, Z.; et al. J. Power Sources 2002, 112, 428−434. (6) Coudun, C.; Grillon, F.; Hochepied, J. F. Colloids Surf., A 2006, 280, 23−31. (7) Tsuge, H.; Watanabe, A. Sixth Japan-Korea Symposium on Separation Technology 2002, 273−276. (8) He, X.; Li, J.; et al. J. Power Sources 2005, 152, 285−290. (9) Tanaka, M. et al., Drill for Analytical Chemistry; Sankyo Shuppan: Tokyo, 1993; p192.

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