Hydrometallurgy, 22 (1989) 249-258 249 Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands

Technical Note

Precipitation of Magnesium Carbonate

D. SHEILA and P.R. KHANGAONKAR

National Metallurgical Laboratory Unit, C.S.I.R. Madras Complex, Madras-600113 (India)

(Received November 28, 1986; revised and accepted August 23, 1988)

ABSTRACT

Sheila, D. and Khangaonkar, P.R., 1989. Precipitation of magnesium carbonate. Hydrometal- lurgy, 22: 249-258.

The Mg (HCO:~)2 solution obtained by pressure carbonation of hydrated MgO is decomposed by heating to form MgCQ of high purity. Factors which influence the precipitation of MgCO:~, such as temperature, concentration, seeding, aeration and stirring, have been studied. Results of tests carried out with A.R. grade MgC12 and Na2CO~ at various temperatures are reported. The activation energies calculated from the rate equations were 63-85 kJ mol- ' , which suggests a chemically controlled reaction mechanism.

INTRODUCTION

The thermal decomposition of Mg (HCO3)2 solution is an important step in the preparation of high-purity MgC03 via the pressure carbonation route. The chemical reaction encountered in the decomposition has been considered by many investigators.

Shukla [1] suggested that MgC03 could be prepared by several methods: (1) reaction between soluble magnesium salts and NaHCQ, Na2CO~ or

(NH4)2C03; (2) carbonation of a slurry of washed Mg(OH)2 to form crystalline MgCO3;

or (3) thermal decomposition of a supersaturated solution of Mg(HCO3)2 ob-

tained from carbonation of Mg(OH)2. The last method serves as a source of high-purity MgCQ, practically free of

calcium ions. Hepburn [ 2 ] reacted MgS04 solution with Na2CO3 to precipitate MgCQ. Schcegrov et al. [3] mixed solutions of Mg(NO3)2 and Na2CO3 at room temperature with continuous stirring to reproduce MgCO~. Chumaovski and coworkers [4 ] carried out precipitation by mixing a solution of MgSO4 or MgC12 with a slurry containing 0.5-10 kg of MgC03. To this a solution con- taining Na2CO3 was added to effect the precipitation.

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Shukla and Datar [5] prepared MgC03 from a solution of Mg(HCO~)2 ob- tained by mixing solutions of MgC12 and N a H C Q , at 25 ° C.

Belyaev and coworkers [6] observed that a metastable solution of Mg (HCQ)2 prepared by carbonation of MgO gradually evolved CO2 and pre- cipitated MgCQ-3H20 on standing in air. When crystallization was carried out at higher temperatures, the resulting precipitate was basic MgCQ. Shukla and Datar [7] prepared basic MgCO3 by raising the temperature of the Mg(HCQ)2 solution to 60 ° C. Belyaev and coworkers [6, 8] reported the ac- tivation energy for thermal decomposition of Mg (HCO3)2 to be 51 kJ m o l - ' at 25-30 ° C and 84.5 kJ mol- 1 at 35-40 ° C. The activation energy for the ther- mal decomposition of Mg(HC03)2 solution prepared by mixing MgC12 and N a H C Q solution was reported as 54 kJ mol-1 [8].

Reddy [9] proposed the following rate expression for the kinetics of crys- tallization of C a C Q

d N ~dr = - K S N 2 ( 1 )

where: N is the concentration of at time t (min) in solution before equilibrium is reached (g 1-1); K is the crystal growth rate constant (1 mol- 1 min- 1 ) / (g seed l - ~ ); and S is the constant initial seed crystal concentration (proportional to the surface area available for growth) (g seed 1-~). The dependence of rate on N 2 indicates second-order kinetics. N is calculated from measured total Mg concentration:

N = TMg--TMg(~) (2)

where: TMg is the total concentration of all dissolved species containing Mg (g l - i ) , and Tig(~) is the equilibrium total Mg concentration (g l-1), obtained from solubility product of MgC03.

The integrated form of Eq. 1 is used to analyse kinetic data

N - l - N o 1 = K S t (3)

where No is the concentration of Mg (g 1 --1) calculated from the Mg to be precipitated from supersaturated solution at zero time.

The present work involves the study of the effects of temperature, concen- tration, seeding, aeration and stirring on the thermal decomposition of Mg ( H C Q ) 2 solutions.

EXPERIMENTAL PROCEDURE

The experiments were carried out in a thermostatic bath which could be controlled to within + 0.1 ° C at the desired temperature. The experiments were

251

carried out with Mg (HC03)e solutions obtained by the pressure carbonation of Mg (OH)e. All the experiments were conducted in a 250-ml flask. A stirring speed of 600 r.p.m, and a flow rate of air of 30 1 h - 1 were used in this study.

Another set of experiments involved A.R. grade MgCle and NaHCO~ solu- tions of known concentration which were allowed to attain the experimental temperature separately. The solutions were then mixed and kept at constant temperature with occasional agitation. At selected time intervals samples were withdrawn and analysed for Mg content by the EDTA method.

RESULTS AND DISCUSSION

The effects of temperature, initial concentration of crystal seeding, aeration and stirring on the thermal decomposition of M g ( H C Q ) 2 are illustrated in Figs. 1-6 and 8 by plotting fraction converted vs. time.

The fraction converted, F ( t ) , is calculated from

Ci - Ct F ( t ) - (4)

C i - C s

where: Ci is the initial concentration of Mg (g 1- ' ) ; Ct is the concentration of Mg at time t (g 1-1); and Cs is the concentration of Mg (g 1- ' ) calculated at saturation on the basis of solubility product of M g C Q .

Figure 1 indicates the effect of temperature on fraction transformed with a starting concentration of ~ 10 g 1-1 Mg. It can be seen that temperature has a marked influence on both the rate and F ( t ) . The maximum conversion (0.93) is obtained at a reaction temperature of 77°C in 46 min. Figure 2 shows the effect of temperature on fraction converted with a starting concentration of

5 g 1-1 Mg. It is observed that the reaction rate is higher when the difference between the initial concentration and the solubility limit is greater, because of the greater driving force. Figure 3 shows the effect of seeding on F ( t ) at 73 and 43 ° C with an initial concentration of ~ 10 g 1-1 Mg and at 43 ° C with an initial concentration of ~ 5 g 1-1 Mg. A comparison of Figs. 1-3 indicates that the crystal seeding has a positive influence on the fraction transformed.

Figure 4 shows the effect of aeration at various temperatures on the fraction converted. The maximum conversion was found to be 99%. It was observed that aeration has a greater influence on the reaction rate than the other vari- ables studied. Figure 5 indicates the effect of stirring on F ( t ) at various temperatures.

The stirring seemed to have the same effect as crystal seeding on the rate. Both aeration and stirring influence the precipitation by carrying away the CO2 evolved from the reaction zone.

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O of 0"8 0 77+C

0.6 X '-4.5"C

'~o"-VI / / ~ • ~ , . , . c

0 . 1 ~ o ~

0 20 LO 60 80 100 120 140 160 100 200 TIME (rain)

Fig. 1. Effect of t empera tu re on f rac t ion t r ans fo rmed F(t) at init ial concen t ra t ion of ~ l0 g 1- Mg.

1.0

0 - 9 ~ 0

0.8

0.7

0'6--

0.5

h 0"/',

0"3

0"2

0'1

ot/

0

!

0 79.4" C

o 5 8 C

I I t I 0 20 /,,0 60 80 100 120 1/.,0 160 180 200

TIME (rain) Fig. 2. Ef fec t o f t empera tu re on f rac t ion t r ans fo rmed F (t) at ini t ia l concen t ra t ion of ~ 5 g I 1 Mg.

253

LL

1.0

0.9

0-8

0,7

0 .6

0"5

0"4.

0.3

0.2

0,1

0

®

I I I I 1 I 20 40 60 80 100 120

T I M E ( r a i n )

I I I I 1'-0 160 180 200

0 73°C

43°C

® 43°C (5 g [ -1Mg)

Fig. 3. Effect of seeding on fraction t ransformed F(t) at 43 and 73°C at ~ 10 g 1-1 Mg, and at 4 3 ° C a t ~ 5 g l 1Mg.

1 " 0

0 . 9 0 3 2 - 5 C

0 , 8 X 5 2 - 5 " C

0.7 ~ 69" C

0 , 6 A 5 8 " C

.1.-, . 0.5 lit ~.3 C b.

O,Z.

0-3

0.2 0

0,1

0 I I I I I I 0 20 40 60 80 100 120 140 160

T I M E ( m i n )

Fig. 4. Effect of aerat ion on fraction t ransformed F ( t ) at various temperatures.

254

0 . 9

0.8 0 60 ' C

0"7 [ ] / . 1 " C

0 , 6 El /" 51"5" C

A •

0 . 5 X 71"5 C

0. / *

0 . 3

0 " 2

0 , 1

0 I I I I I I I

0 20 4,0 60 80 100 120 1/.0 160

T I M E ( m l n )

Fig. 5. Effect of stirring on fraction transformed F(t) at various temperatures.

The kinetic data were tested according to Eq. 3. The linear plots of N - 1 _ N o 1 as a function of time confirm the validity of this equation. Figure 6 shows such a plot for different temperatures with an initial concentration ~ 10 g l - 1 Mg. The activation energy (83.7 kJ mol - 1 ) obtained from the slopes of Fig. 7 indicates a surface-controlled mechanism. The value reported by Bel- yaev and coworkers [6] for the thermal decomposition of Mg(HC03)2 ob- tained by carbonation of MgO under pressure was 84.5 kJ mol- 1.

Reactions involving MgCl2 and NaHC03

Figure 8 shows the plot N - 1 N o 1 vs. t at four temperatures. The linear plot confirms the validity of Eqn. 3. The activation energy of 63-84 kJ mol- 1 obtained here is again in close agreement with Belyaev's findings of 54 kJ mo1-1 (Fig. 9).

Induction period

As observed from the figures, the induction period depends on the experi- mental conditions. Increased crystal seeding, high temperatures, increased aeration and stirring considerably reduce the induction time.

255

,.TE / 1"6

1-5 1

1.4

1"3

1,2

1.1

1.0

-1 o'9 I

~ l z O ' 8

0.7

10'8

O 77°C

I-I 5 1 3 ° C

A Z,0 3°C

• 2 8 6 ° C

X M., 5°C

---J 0-6 - 20

0 '5 I I

0"~ 11

Ii 0 '/~

0,3 II ~ y , 0.2 , J - ,' ~I°'16 / ,"

o . , . . . . .

0 t" . . . . . . . . . . . . . $ l,o

0 20 &O 60 80 100 120 1/-0 160 180 2'00 220 2LO 260 280

05

TIME(rain)

Fig. 6. Plot of N - 1 _ N o 1 vs. t at various temperatures at initial concentration of ~ 10 g l -1 Mg.

Effect of NaC1 on rate constants

Shukla and Datar [5] found the values of rate constants obtained at differ- ent concentrations of M g ( H C Q ) 2 prepared by the reaction between MgC12 and N a H C Q to be half of those found with samples obtained by the carbona- tion of Mg(OH)2. This was attr ibuted to the presence of NaC1 formed in the reaction between MgC12 and N a H C Q . NaC1 reduces the activity coefficient of HCOy ions, thereby increasing the solubility of MgC03. However, in the pres- ent case, the rate constants obtained for the thermal decomposition of

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13 AERATION ; AE: 88 kJ mot -'~

0

-1

- 2

- 3

" * STIRRING ; a E = 82.9 kJ tool "1

O T E M P ( C i = 1 0 g L - 1 ) ; a ' E = 81 6 kJ tool'1

L. . . . . . . . . .

_,~ ~...._ ~ I I ~ ~ "

L . . . . . ,; :,; ~2---'~"~

I I I I

2 " 8 2 " 9 3 ' 0 3.1 3"2

1 - 3 - - x l O

T Fig. 7. P lo t of log K vs. T - ~ X 1 0 - 3 ( ac t iva t ion energy ).

I

3"3

5"5

5,0

4.5

4.0

3.5

o ~ ] Z 3 " 0

i

~I z 2.5

2.0

1.5

1.0

0.5

/ 0 71°C

o / / oss:c / / A6~.c • " t 3 / . % o / /

/',' " / i I i

/ !"" / !

0 - - ~ - -

,,,,

0 O, ),i~...~-, ~--~_.,r_.--?~-----~--- ;- , , , ,

i, 8 12 16 20 2& 28 32 36 l.O &/* 48 52 56 60

T I M E ( r a i n )

Fig. 8. P lo t o f N - 1 _ N o I vs. t ( m i n ) a t va r ious t e m p e r a t u r e s .

257

0 -

~ CORRELATION COEFFICIENT = 0 7072

- 2

-3 $ I I 1 2 . 8 2 . 9 3 , 0 3 '~ 3 . 2

l l T x l 0 -3

Fig. 9. Plot of log Kvs. T × 10 -a (activation energy).

Mg (HC03) 2 prepared by reaction between MgC12 and NaHCO~ are higher than those obtained for the samples prepared by carbonation of Mg(OH)2. This is due to the different experimental conditions, as the mixing of the MgC12 and NaHC03 was carried out at elevated reaction temperatures. However, the Mg(HC03)2 obtained by the carbonation of Mg(OH)2 has to attain the reac- tion temperature before the reaction commences, and as a consequence the rates are slower than those of the former reaction.

CONCLUSIONS

The parameters which influence precipitation of MgC03 from Mg(HCQ)2 solutions as well as by reacting MgC12 and Na2C03 have been studied. The precipitation has been found to be enhanced in the order aeration > temper- ature > stirring. The higher activation energy of 84 kJ mol- 1 obtained in this study suggests that the precipitation is a surface-controlled reaction. The lin- ear plot of N - 1 - N ~ ~ vs. t indicates second-order kinetics.

REFERENCES

1 Shukla, B.K., 1972. Studies on the formation, precipitation, kinetics, methods of manufacture and surface properties of magnesium carbonate, Ph.D. Thesis, Bombay University.

2 Hepburn, J.R.J., 1940. Chemical nature of precipitated basic magnesium carbonate. J. Chem. Soc., 96-99.

3 Schcegrov, L.N., Skrobotun, V.N. and Ryadchenko, A.G., 1965. Some characteristics of for- mation of solid phase from liquid. Akad. Nauk SSSR, Sb State, pp. 59-61 (in Russian).

4 Gotovtsev, V.V., Kudrlyavtseva, V.P., Shvetsov, Yu.A., Bulat, A.E. and Chumaovski, V.A., 1979. Preparation of basic carbonate, Otkrytiya, Izobret Prom obraytsy, Tovarays Znaki, pp. 45-91,

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5 Shukla, B.K. and Datar, D.S., 1972. Studies on the carbonates of magnesium - - Part II. Ki- netics of precipitation of magnesium carbonate from a solution of magnesium bicarbonate. Ind. J. Appl. Chem., 35 (1): 30-34.

6 Tkach, G.A., Telitchenko, V.A., Seryi, B.G. and Belyaev, E.K., 1975. Thermal decomposition of magnesium bicarbonate solutions. J. Appl. Chem. USSR, 48: 2191-2944.

7 Shukla, B.K. and Datar, D.S., 1971. Studies on the carbonation of magnesium - - Part I. Car- bonation of slurry of Mg(OH)2. Ind. J. Appl. Chem., 34 (3): 00.

8 Belyaev, E.K. and Telitchenko, V.A., 1976. Activation energy and order of reaction. Zh. Prikl. Khim., 49 (1): 193-195.

9 Reddy, M.M., 1975. Kinetics of calcium carbonate formation. Verh. Int. Verein. Limnol., 19: 429-438.