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Effects of temperature on the scaling of calcium sulphate in pipes Tung A. Hoang a,b, , H. Ming Ang a , Andrew L. Rohl b,c a Department of Chemical Engineering, Curtin University of Technology, PO Box U 1987, Perth 6845, WA, Australia b Nanochemistry Research Institute and A.J. Parker CRC for Hydrometallurgy, Curtin University of Technology, PO Box U 1987, Perth 6845, WA, Australia c Interactive Virtual Environments Centre (IVEC), Technology Park, Kensington WA 6151, Australia Available online 25 November 2006 Abstract Calcium sulphate scaling is a serious problem encountered in many industrial and domestic applications. Supersaturation has been proven to be the major driving force of scale formation, but the solubility of calcium sulphate changes with temperature. The main purpose of this work is to investigate the effects of temperature on the formation of calcium sulphate scales in pipes, using a pipe flow system. Various levels of supersaturation of the calcium sulphate solution have been employed at different temperatures. Results indicated that higher temperature produced a large increase of scale amounts and a significant decrease of induction periods. Many forms of hydrated calcium sulphate were created at high temperature. The relationship between deposited scale mass and temperature was deduced from experimental data. From the relationship between induction period and temperature activation energies of the surface nucleation were estimated to be in the range of 42 to 48 kJ mol 1 . © 2006 Elsevier B.V. All rights reserved. Keywords: Calcium sulphate; Scale deposit; Pipe flow system; Temperature effects 1. Introduction Scaling or the accumulation of materials depositing on the surface of equipment is a complicated phenomenon, which significantly affects a wide range of industrial processes, with serious technical and economic consequences [14]. Calcium sulphate is frequently encountered both in nature and in industry [59]. It is also the most unwelcome scalant in the production of oil and gas, in water cooling systems and in hydrometallurgical processes [10]. Early studies on gypsum scaling control mainly focussed on the kinetics of scale formation [11,12] but later investigations emphasized the influence of external factors [8,9,1322].A systematic study of the effects of various process parameters and the efficacy of some inorganic and organic additives in controlling the formation of calcium sulphate in pipes was first undertaken by researchers of the Department of Chemical Engineering and A.J. Parker CRC for Hydrometallurgy at Curtin University of Technology, WA. [23,24]. Using a self designed pipe flow system and systematically altering the process parameters they found that the scaling of calcium sulphate was affected significantly by the supersaturation of the solution, run time and flow rate [24]. Calcium sulphate precipitates in many different solid phases: dihydrate (or gypsum), hemihydrate and anhydrous although gypsum is the most common one at ambient temperature [25,26]. The solubility of all forms of calcium sulphate changes with increasing temperature (Fig. 1). The driving force for crystallisation depends on the supersaturation level of the solution. However, activation energy is another important factor that needs to be considered. 2. Experimental 2.1. Equipment setup The pipe flow system consisted of two glass vessels, one containing CaCl 2 solution and the other Na 2 SO 4 solution. The vessels were placed in a water bath with a heating unit to control temperature. The solutions were transported through the test section at the same flow rate by using a double-line peristaltic pump (Fig. 2). The two solutions were mixed at the inlet of the test section, which consisted of two stainless steel tubular units. Eight pairs of stainless steel semi-annular coupons were inserted into the tubular units, serving as scaling surface. Each semi Powder Technology 179 (2007) 31 37 www.elsevier.com/locate/powtec Corresponding author. E-mail address: [email protected] (T.A. Hoang). 0032-5910/$ - see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.powtec.2006.11.013

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Page 1: pembentukan kerak

9 (2007) 31–37www.elsevier.com/locate/powtec

Powder Technology 17

Effects of temperature on the scaling of calcium sulphate in pipes

Tung A. Hoang a,b,⁎, H. Ming Ang a, Andrew L. Rohl b,c

a Department of Chemical Engineering, Curtin University of Technology, PO Box U 1987, Perth 6845, WA, Australiab Nanochemistry Research Institute and A.J. Parker CRC for Hydrometallurgy, Curtin University of Technology, PO Box U 1987, Perth 6845, WA, Australia

c Interactive Virtual Environments Centre (IVEC), Technology Park, Kensington WA 6151, Australia

Available online 25 November 2006

Abstract

Calcium sulphate scaling is a serious problem encountered in many industrial and domestic applications. Supersaturation has been proven to bethe major driving force of scale formation, but the solubility of calcium sulphate changes with temperature. The main purpose of this work is toinvestigate the effects of temperature on the formation of calcium sulphate scales in pipes, using a pipe flow system. Various levels ofsupersaturation of the calcium sulphate solution have been employed at different temperatures. Results indicated that higher temperature produceda large increase of scale amounts and a significant decrease of induction periods. Many forms of hydrated calcium sulphate were created at hightemperature. The relationship between deposited scale mass and temperature was deduced from experimental data. From the relationship betweeninduction period and temperature activation energies of the surface nucleation were estimated to be in the range of 42 to 48 kJ mol−1.© 2006 Elsevier B.V. All rights reserved.

Keywords: Calcium sulphate; Scale deposit; Pipe flow system; Temperature effects

1. Introduction

Scaling or the accumulation of materials depositing on thesurface of equipment is a complicated phenomenon, whichsignificantly affects a wide range of industrial processes, withserious technical and economic consequences [1–4]. Calciumsulphate is frequently encountered both in nature and in industry[5–9]. It is also the most unwelcome scalant in the production ofoil and gas, in water cooling systems and in hydrometallurgicalprocesses [10].

Early studies on gypsum scaling control mainly focussed onthe kinetics of scale formation [11,12] but later investigationsemphasized the influence of external factors [8,9,13–22]. Asystematic study of the effects of various process parametersand the efficacy of some inorganic and organic additives incontrolling the formation of calcium sulphate in pipes was firstundertaken by researchers of the Department of ChemicalEngineering and A.J. Parker CRC for Hydrometallurgy atCurtin University of Technology, WA. [23,24]. Using a selfdesigned pipe flow system and systematically altering theprocess parameters they found that the scaling of calcium

⁎ Corresponding author.E-mail address: [email protected] (T.A. Hoang).

0032-5910/$ - see front matter © 2006 Elsevier B.V. All rights reserved.doi:10.1016/j.powtec.2006.11.013

sulphate was affected significantly by the supersaturation of thesolution, run time and flow rate [24].

Calcium sulphate precipitates in many different solid phases:dihydrate (or gypsum), hemihydrate and anhydrous althoughgypsum is the most common one at ambient temperature[25,26]. The solubility of all forms of calcium sulphate changeswith increasing temperature (Fig. 1). The driving force forcrystallisation depends on the supersaturation level of thesolution. However, activation energy is another important factorthat needs to be considered.

2. Experimental

2.1. Equipment setup

The pipe flow system consisted of two glass vessels, onecontaining CaCl2 solution and the other Na2SO4 solution. Thevessels were placed in a water bath with a heating unit to controltemperature. The solutions were transported through the testsection at the same flow rate by using a double-line peristalticpump (Fig. 2). The two solutions were mixed at the inlet of thetest section, which consisted of two stainless steel tubular units.Eight pairs of stainless steel semi-annular coupons were insertedinto the tubular units, serving as scaling surface. Each semi

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Fig. 1. Solubility of calcium sulphate in water as a function of temperature. (Datafrom Linke, Marshall and Slusher, Silcock, [27–29]).

32 T.A. Hoang et al. / Powder Technology 179 (2007) 31–37

annular coupon had a length of 3 cm and an internal diameter of1.3 cm. The mixed solution coming out from the test sectionwas collected into a waste container. The test section andconnection hoses were covered with insulating material to keepthe temperature constant.

The conductivity of the output solution was measured usinga Yokogawa Model SC82 conductivity meter, which has twooperational ranges from 0 to 20 μS/cm and from 0 to 200 mS/cm. The meter was equipped with an auto range function and anautomatic temperature compensation.

2.2. Chemicals

Calcium chloride dihydrate (CaCl2U2H2O), AR grade, andsodium sulphate (Na2SO4), AR grade, were from Chem SupplyPty Ltd., South Australia.

Fig. 2. Diagram of a pipe flow system.

2.3. Procedure

Equimolar solutions of calcium chloride and sodiumsulphate in predetermined concentrations were placed in awater bath until they reached the required temperature. Thesesolutions were pumped and mixed together before going intothe test section. After a predetermined time, the pump wasswitched off and the test section disconnected. The semi-annular coupons were withdrawn out of the units and placed inan oven at 60 °C overnight, then cooled down to roomtemperature and weighed. The difference between the massesbefore and after the experiment was the mass of scale deposit.The concentration of the solutions, the run time and thetemperature were altered to investigate their effects. To estimatethe induction period, a 50 mL aliquot of the existing solutionwas taken every 2 min during the run and its conductivity wasmeasured immediately. The induction period could be deter-mined by drawing the best fit lines through the two sections ofthe conductivity-time curve and reading the time correspondingto their intersection.

3. Results and discussion

3.1. Effect of temperature on scale formation

The effect of temperature is shown in Fig. 3. Generally,within the investigated temperature range from 20 °C to 60 °C,the higher the temperature, the more scales formed. Tempera-ture seemed to significantly promote the scaling of calciumsulphate. For a 0.075 M solution of calcium sulphate, a smallrise of temperature from 20 °C to 30 °C tripled the scaleamounts formed after 3 h from 0.0212 kg/m2 to 0.0603 kg/m2.At 40 °C, scale mass increased to 0.1599 kg/m2, almost 8 timesas much as that at 20 °C. At 50 °C and 60 °C, the scales wereformed so abundantly that it started blocking the flow after 2 h.In contrast, the solubility of gypsum does not change muchwithin this temperature range (Fig. 1). It slightly increases from2.02 g/L at 20 °C to 2.08 g/L at 30 °C and 2.10 g/L at 40 °C,

Fig. 3. Effect of temperature on the scale formation. Flow rate=30 mL/min., Ca2+

concentration=0.075 M, run time=3 h.

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Fig. 4. Correlation between log (scale mass) and temperature. Ca2+

concentration=0.075 M, flow rate=30 mL/min., run time=3 h. Fig. 6. Comparison of scaling rates at different temperatures. Ca2+

concentration=0.05 M. Flow rate=30 mL/min: a. 20 °C, b. 25 °C, c. 30 °C,d. 40 °C.

33T.A. Hoang et al. / Powder Technology 179 (2007) 31–37

then reduces to 2.07 g/L at 50 °C and 2.01 g/L at 60 °C. Theseresults indicate that solubility change was not the only cause ofthe increase in scaling rate. Higher temperatures providedenough energy to the molecules to overcome the activationenergy of the precipitation reaction and sped up the transport ofscale components from the bulk solution to the surface.

To determine the correlation between scaling and temperature,logarithm of scale mass is plotted against inverse temperature(Fig. 4). The curve is linear, indicating an exponential rela-tionship between the scaling mass and temperature. An equationcould be calculated from the graph using Microsoft Excel.

logðscale massÞ ¼ −4047:11T

� �þ 12:141 ð1Þ

with the Pearson product moment correlation coefficientR2 =0.9995.

Fig. 5. Plots of scale amounts against time at various temperatures. Flowrate=30 mL/min. Ca2+ concentration=0.05 M. Temperature: a. 20 °C, b. 25 °C,c. 30 °C, d. 40 °C.

From Eq. (1) it can be seen that the higher the temperatureor the lower 1 /T, the larger log (scale mass), and accordinglyscale mass increased by an exponential factor when temper-ature rose.

scale mass ¼ CT :10−4047:11T ð2Þ

where CT is a constant.Since secondary nucleation is not significantly affected by

temperature, it could be concluded that the primary nucleationdominated the scaling mechanism. This is consistent withresults from other researchers, who reported that in the precipi-tation of sparingly soluble substances, secondary nucleationeither did not occur [30] or did occur only to a small extent[31,32]. The number of crystals formed by secondary nucleationduring precipitation was substantially lower than those resultingfrom primary nucleation [33].

3.2. Relationship between scaling rate and temperature

Plotting the amounts of calcium sulphate formed after theinduction period at different temperatures against time showsthat the curves are parabolic indicating scale amount is aquadratic function of run time (Fig. 5).

To obtain the value of average scaling rate, the scaleamount deposited after every hour was divided by time.Plotting average scaling rate against run time produces a

Table 1Rate constant at different temperatures Ca2+ concentration=0.05 M, flowrate=30 mL/min

Temperature k(kg/m2.min2) R2

40 °C 1.4215×10−6 0.998630 °C 1.1973×10−6 0.999325 °C 1.0037×10−6 0.991020 °C 0.7850×10−6 0.9940

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Fig. 7. Dependence of induction period on inverse temperature. Flow rate=30 mLmin−1. Coupon diameter=1.3 cm. CaSO4 concentration: a, 0.05 M; b, 0.075 M.

Table 3Activation energy for calcium sulphate dihydrate crystallisation from literature

Mode CCaSO4 (mol L−1) Ea (kJ mol−1) References

Unseeded batch 0.04 160 [36]Unseeded batch 0.03–0.1 53–61 [6]Unseeded batch 0.03–0.1 51–71.5 [37]Seeded batch 0.0414 63±2 [38]Seeded batch 0.0148 59±8 [39]Seeded batch 0.04 75 [40]Seeded batch 0.01–0.02 60 [41]Seeded batch 0.0155 60 [42]Seeded batch 0.05 60 [43]

34 T.A. Hoang et al. / Powder Technology 179 (2007) 31–37

straight line (Fig. 6) with the slope k as indicated in Eq. (3)below:

1Ar

dmdt

¼ kt þ C ð3Þ

where,

dm / dt the scaling rate (kg/min)Ar the scaling area (m2)t run time (min)C constantk constant, referring to the change rate of the scaling rate

(kg/m2.min2)

The k constants derived from the curves in Fig. 6 are listed inTable 1.

It is indicated that the rate constant also increases whentemperature rises.

3.3. Effect of temperature on induction period

Induction period is the period of time that elapses betweenthe achievement of supersaturation and the first detection ofcrystal deposit onto the pipe walls.

tind ¼ tn þ tg ð4Þwhere tn is the time necessary for the critical nucleus to be formedand tg represents the time for this nucleus to grow to visible size.

Table 2Activation energy for different supersaturation ratios estimated from the slopesof the curves

CaSO4 conc. (mol L−1) Slope=E /2.303R E (kJ mol−1)

0.050 −2517.4 (R2=0.9986) 48.20.075 −2215.1 (R2=0.9993) 42.4

From the statistical concept of nucleation, the mean time ofcritical nucleus formation is proportional to J −1 [34]

tn~1J

ð5Þ

where J is the steady-state rate of homogeneous or heteroge-neous primary nucleation. If the time period necessary for thecritical nucleus to be formed is much longer than the time for itto grow to a measurable size, the induction period is determinedprimarily by the former.

Thus,

tind ¼ tn ð6ÞThe rate of solid phase nucleation from a supersaturated

solution is [35]

J ¼ X exp −DG⁎

kT

� �ð7Þ

where,

X the pre-exponential factor depending on diffusioncoefficient in solution and interplanar distance in thecrystal lattice (s−1 m−3)

ΔG⁎ the energy barrier to nucleation, which corresponds tothe change of Gibbs energy accompanying formationof the critical nucleus (J)

k Boltzmann constantT absolute temperature (°K).

Combining these three equations and performing a few mathe-matical transformations gives a relationship between the inductionperiod and temperature, which is similar to the Arrhenius equationfor the dependence of rate constant on temperature

logðt−1indÞ ¼ A−E

2:303RTð8Þ

where,

E the molar activation energy for nucleation (J mol−1)R universal gas constant=8.31451 J K−1 mol−1

A constant

Table 4Water content of the calcium sulphate scales deposited at different temperature

Temperature (°C) 20 30 40 50 60Water % 20.21 20.07 19.40 13.52 6.31

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35T.A. Hoang et al. / Powder Technology 179 (2007) 31–37

Eq. (8) indicates that the induction period of the scalingdepends strongly on temperature. The molar activation ener-gy corresponding to any supersaturation ratio can be esti-mated from the slope of the curve plotting log (tind

−1) against1 /T. (Fig. 7). Results of these estimations are given inTable 2.

The activation energy values estimated here correspondingto CaSO4 concentrations of 0.050 and 0.075 mol L−1 are 48.2and 42.4 kJ mol−1, respectively, a little lower than those found

Fig. 8. SEM images of calcium sulphate scale deposited on the pipe walls at20 °C. From top: a, ML=75; b, ML=300; c, ML=1000.

Fig. 9. SEM images of calcium sulphate scale deposited on the pipe walls at50 °C. From top: a, ML=75; b, ML=300; c, ML=1200.

by other researchers (Table 3). These values suggest that themechanism of scaling in pipes might be different from scaleprecipitation in stirred crystallizers. The wall of a pipe can act asa nucleating site similar to foreign substances causingheterogeneous nucleation, which requires less energy becausethe foreign substances reduce the surface energy of the nucleus[44].

Also, the mechanism could be partly contact nucleation,since the change of free energy for contact nucleation is lower

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Fig. 10. SEM images of calcium sulphate scale deposited on the pipe walls at60 °C. From top: a, ML=75; b, ML=300; c, ML=1200.

36 T.A. Hoang et al. / Powder Technology 179 (2007) 31–37

than that of homogeneous or heterogeneous nucleation [45].Contact nucleation occurs when a growing particle contacts thepipe wall. However, the deposition of calcium sulphate isgenerally a surface controlled process, since the activationenergy for surface controlled precipitation is typically greaterthan 40 kJ mol-1, while for a diffusion controlled process it isexpected to be ≈ 10 – 20 kJ mol-1 [34].

3.4. Forms of scale deposit

The solubility of calcium sulphate as a function of tempera-ture is illustrated in Fig. 1. Within the range from 0 to 40 °C,calcium sulphate dihydrate has the lowest solubility. Therefore,it is expected that the scale will contain mainly the calciumsulphate dihydrate form. At temperatures higher than 40 °C thesolubility of other forms decreases significantly, resulting in thedeposition of many forms at the same time. Table 4 presents thewater content in the scales determined by TGA (Thermalgravimetric analysis).

The water content in pure calcium sulphate dihydrate is20.9%. Therefore, more than 90% of the scales depositing at atemperature lower than 40 °C would be in the dihydrate form.At 50 °C and 60 °C, the water content decreased significantly.Results from X-Ray Diffraction analysis indicated that thescales formed at 30 °C or 40 °C contain only calcium sulphatedihydrate while there are other hydrate and anhydrous formsexist in the scales formed at higher temperatures. This led to theextremely fast deposition of scales at 50 °C and 60 °C, causingthe blockage of the flow in the test units.

SEM images of the scales reveal different crystalline shapesformed in low and high temperatures. At 20 °C there are scales inthe floral structures created by numerous plate-like and needle-like crystals. The needles are orthorhombic or hexagonalprismatic, which are typical forms of gypsum crystals (Fig. 8,ML =magnification level). On the other hand the floral structuresdisappear at 50 °C, the number of plate-like crystals reduces andneedles with thinner and longer shape dominate (Fig. 9). Thetrend continues as temperature rises to 60 °C (Fig. 10).

4. Conclusion

Solution temperature proved to have important effects onthe formation of calcium sulphate scales. High temperaturespeeds up the scaling process by significantly reducing theinduction time and enormously increases the scaling rate.Results indicate the scaling rate is a linear function of time andthe scale mass is an exponential function of inverse absolutetemperature. The activation energies for scaling processestimated from logarithmical relationship between inductionperiod and inverse temperature are in the range from 42 to48 kJ mol−1, indicates that the deposition of calcium sulphateis generally a surface controlled process, but contact nucle-ation partly affects the scaling mechanism. High temperaturealso changes the hydrate forms of the scale deposit. At tem-peratures lower than 40 °C, only calcium sulphate dihydrateis formed, but beyond 40 °C, many hydrated and anhydrousforms precipitate simultaneously.

References

[1] D. Hasson, Precipitation fouling, in: E. Somerscales, J. Knudsen (Eds.),Fouling of Heat Transfer Equipment, Hemisphere, New York, 1981,pp. 527–568.

[2] A. Pritchard, The economics of fouling, in: L. Melo, T. Bott, C. Bernado(Eds.), Fouling Science and Technology, Kluwer Academics, Dordrecht,1988, pp. 31–45.

Page 7: pembentukan kerak

37T.A. Hoang et al. / Powder Technology 179 (2007) 31–37

[3] L. Legrand, P. Leroy, Prevention of Corrosion and Scaling in Water SupplySystems, Ellis Horwood Series in Water and Waste Water Technology,New York, 1990.

[4] R. Rosset, S. Douville, M. Ben Amor, K. Walha, Inhibition of scaleformation by Southern Tunisia geothermal water field experiments, Revuedes Sciences de l'Eau 12 (4) (1999) 753–764.

[5] M. Tadros, I. Mayes, Linear growth rates of calcium sulfate dihydratecrystals in the presence of additives, Journal of Colloid and InterfaceScience 72 (2) (1979) 245–254.

[6] S. He, J. Oddo, M. Tomson, The seeded growth of calcium sulfatedihydrate crystals in NaCl solutions up to 6 m and 90 °C, Journal ofColloid and Interface Science 163 (1994) 372–378.

[7] P. Klepetsanis, E. Dalas, P. Koutsoukos, Role of temperature in thespontaneous precipitation of calcium sulphate dihydrate, Langmuir 15(1999) 1534–1540.

[8] O. Linnikov, Investigation of the initial period of sulphatescale formation.Part 1-Kinetics and mechanism of calcium sulphate surface nucleation atits crystallization on a heat-exchange surface, Desalination 122 (1) (1999)1–14.

[9] J.F. Adams, V.G. Papangelakis, Gypsum scale formation in continuousneutralization reactors, Canadian Metallurgical Quarterly 39 (4) (2000)421–431.

[10] G. Van Rosmalen, P. Daudey, W. Marchee, An analysis of growthexperiments of gypsum crystals in suspension, Journal of Crystal Growth52 (1981) 801–811.

[11] D. Hasson, J. Zahavi, Mechanism of calcium sulfate deposition on heattransfer surfaces, Industrial and Engineering Chemistry Fundamentals 9(1) (1970) 1–10.

[12] S. Liu, G. Nancollas, The crystal growth of calcium sulfate dihydrate in thepresence of additives, Journal of Colloid and Interface Science 44 (3)(1973) 422–429.

[13] H. Mori, M. Nakamura, S. Toyama, Crystallization fouling of calciumsulfate dihydrate on heat transfer surfaces, Journal of ChemicalEngineering of Japan 29 (1) (1996) 166–173.

[14] A. Lancia, D. Musmarra, M. Prisciandaro, Measuring induction period forcalcium sulfate dihydrate precipitation, Journal of American Institute ofChemical Engineering 45 (2) (1999) 390–397.

[15] H. Muller-Steinhagen, Q. Zhao, A. Helali-Zadeh, The effect of surfaceproperties on CaSO4 scale formation during convective heat transfer andsubcooled flow boiling, The Canadian Journal of Chemical Engineering 78(1) (2000) 12–20.

[16] O. Linnikov, Investigation of the initial period of sulphatescale formation.Part 2—Kinetics of calcium sulphate crystal growth at its crystallization ona heat-exchange surface, Desalination 128 (1) (2000) 35–46.

[17] O. Linnikov, Investigation of the initial period of sulphatescale formation.Part 3—Variations of calcium sulphate crystal growth at its crystallizationon a heat-exchange surface, Desalination 128 (1) (2000) 47–55.

[18] M. Sudmalis, R. Sheikholeslami, Coprecipitation of CaCO3 and CaSO4,The Canadian Journal of Chemical Engineering 78 (1) (2000) 21–31.

[19] S. Lee, C. Lee, Effect of operating conditions on CaSO4 scale formationmechanism in nanofiltration for water softening, Water Research 34 (15)(2000) 3854–3866.

[20] R. Sheikholeslami, M. Ng, Calcium sulfate precipitation in the presence ofnondominant calcium carbonate: thermodynamics and kinetics, Industrialand Engineering Chemistry Research 40 (16) (2001) 3570–3578.

[21] T.A. Hoang, H.M. Ang, A.L. Rohl, Development of a laboratoryexperimental system for studying gypsum scale formation in pipes, 10thAsian Pacific Confederation of Chemical Engineering, Kitakyushu, Japan,2004.

[22] T.A. Hoang, H.M. Ang, A.L. Rohl, M.I. Jeffrey, A study of gypsum scaleformation using quartz crystal microbalance, Developments in ChemicalEngineering and Mineral Processing 14 (1/2) (2006) 313–321.

[23] S. Muryanto, H. Ang, G. Parkinson, Crystallisation kinetics of calciumsulphate dihydrate in the presence of additives, 2nd World EngineeringCongress, Sarawak, Malaysia, 2002.

[24] T.A. Hoang, H.M. Ang, A.L. Rohl, Effects of an organic substrate andprocess parameters on gypsum scale formation in pipes, 31st AustralasianChemical Engineering Conference, Adelaide, Australia, 2003.

[25] P. Dydo, M. Turek, J. Ciba, Scaling analysis of nanofiltration systems fedwith saturated calcium sulfate solutions in the presence of carbonate ions,Desalination 159 (2003) 245–251.

[26] A. Helalizadeh, H. Muller-Steinhagen, M. Jamialahmadi, Mixed saltcrystallisation fouling, Chemical Engineering and Processing 39 (2000)29–43.

[27] W.F. Linke, Solubilities of Inorganic and Metal-Organic Compounds,D. van Nostrand Co, New Jersey, 1958.

[28] W.L. Marshall, R. Slusher, Thermodymamics of calcium sulfate dihydratein aqueous sodium chloride solutions, 0-110°, Journal of PhysicalChemistry 70 (12) (1966) 4015–4027.

[29] H.L. Silcock, Solubilities of Inorganic and Organic Compounds, vol. 3,Part 3, Pergamon Press, Oxford, 1979.

[30] J. Wey, J. Terwillinger, A. Gingello, Analysis of AgBr precipitation in acontinuous suspension crystallizer, AIChE Symposium Series 193 (76)(1980) 34–42.

[31] L. Breèeviæ, J. Garside, On the measurement of crystal size distributions inthe micrometer size range, Chemical Engineering Science 36 (5) (1981)867–869.

[32] L. Blomen, E. Will, O. Bijvoet, H. van der Linden, Growth kinetics ofcalcium oxalate monohydrate: II. The variation of seed concentration,Journal of Crystal Growth 64 (2) (1983) 306–315.

[33] H.A. Van Straten, M.A.A. Schoonen, P.L. de Bruyn, Precipitation fromsupersaturated aluminate solutions. III. Influence of alkali ions with specialreference to Li+, Journal of Colloid and Interface Science 103 (2) (1985)493–507.

[34] J.W. Mullin, Crystallization, Fourth edition, Butterworth-Heinemann,Oxford, 2001.

[35] O. Söhnel, J. Garside, Precipitation: Basic Principles and IndustrialApplications, Butterworth-Heinemann, Oxford, 1992.

[36] M. Ching, E.R. McCartney, The use of a membrane electrode to study thecrystallisation of calcium sulphate from aqueous solution. I. The relationbetween calcium ion activity and crystallisation rate, Journal of AppliedChemistry and Biotechnology 23 (1973) 441–450.

[37] F. Alimi, A. Gadri, Kinetics and morphology of formed gypsum,Desalination 166 (2004) 427–434.

[38] S. Liu, G. Nancollas, The kinetics of crystal growth of calcium sulfatedihydrate, Journal of Crystal Growth 6 (1970) 281–289.

[39] G.H. Nancollas, M.M. Reddy, F. Tsai, Calcium sulfate dihydrate crystalgrowth in aqueous solution at elevated temperatures, Journal of CrystalGrowth 20 (1973) 125–134.

[40] B. Mile, A.T. Vincent, C.R. Wilding, Studies of the effects of electrolyteson the rates of precipitation of calcium sulphate dihydrate using an ion-selective electrode, Journal of Chemical Technology and Biotechnology 32(1982) 975–987.

[41] G.J. Witkamp, J.P. Van der Eerden, G.M. Van Rosmalen, Growth ofgypsum, Journal of Crystal Growth 102 (1990) 281–289.

[42] J. Kontrec, D. Kralj, L. Breèeviæ, Transformation of anhydrous calciumsulphate into calcium sulphate dihydrate in aqueous solutions, Journal ofCrystal Growth 240 (2002) 203–211.

[43] Muryanto, S. The role of impurities and additives in the crystallisation ofgypsum. PhD Thesis, Curtin University of Technology, Bentley, W.A.(2002).

[44] J. Cowan, D. Weintritt, Water-formed Scale Deposits, Gulf Publishing,Houston, Texas, 1976.

[45] J.A. Dirksen, T.A. Ring, Fundamentals of crystallization: Kinetic effectson particle size distributions and morphology, Chemical EngineeringScience 46 (10) (1991) 2389–2427.