Indian Journal of ChemistryVol. 26A, April 1987, pp. 30\-303

Kinetics of Dissolution of Alkaline Earth Fluorides

S M HAMZA& F AALYChemistry Department, Faculty of Science, Menoufia University

andMAEL-RIES*

National Organization for Drug Control & Research, Giza, Egypt

Received 20 May 1986; revised and accepted 3 November 1986

Constant composition technique has been used to study the dissolution of the fluorides of magnesium, calcium, stron-tium and barium in aqueous solution at an ionic strength of 0.15 M, as a function of concentration, temperature and fluiddynamics. The results have been interpreted in the light of the stability constants, lattice energy, thermodynamic parame-ters and dehydration of cations.

Studies carried out so far on the crystal growth anddissolution of alkaline earth fluorides have not pro-vided much insight into the rate and mechanism ofdissolution of alkaline earth fluorides. In order tolearn more about the rate and mechanism, we have in-vestigated the kinetics of dissolution of alkaline earthfluorides employing the constant composition meth-odl-3• This method is particularly useful at very lowundersaturation.

Materials and MethodsUndersaturated solutions of"the alkaline earth fluo-

rides were prepared in triply distilled, deionised wa-ter, using both ultrapure (Alfa Chemicals) and rea-gent grade (J T Baker) chemicals. Metal ion concentr-ations were determined by atomic absorption spec-troscopy.

Seed crystals of the alkaline earth fluorides wereprepared by precipitation from a mixed solution ofpotassium fluoride and metal nitrate at 25°C. Theseeds were washed with saturated solutions of the me-tal fluorides and allowed to age for at least one monthat '25°C, until the specific surface area (SSA) reacheda constant value.

Kinetic measurementsDissolution experiments were made at 10.1°C in a

double-walled reaction cell (300 ml capacity) fittedwith a teflon lid. The inside of the cell had a lining ofpolyethylene. Nitrogen gas, saturated in the electro-lyte of the medium at the temperature of the reactionvessel, was bubbled continuously through the solu-tions during the experiments. At the beginning of eachexperiment, the fluoride electrode was standardizedby adding aliquots of potassium nitrate solution in thecell. Subsequently, undersaturated solutions of de-sired concentrations were prepared by the slow addi-

tion of metal nitrate solution. Following the introduc-tion of metal fluoride seed crystals, the activities ofionic species were maintained constant (monitoredby means of a pH meter) by the addition of potassiumnitrate as diluent. Aliquots of reaction mixture werewithdrawn at regular time intervals and filteredthrough 0.22!lm millipore. The filtrate was analysedfor metals by atomic absorption spectroscopy using aPerkin-Elmer model 503 spectrophotometer in or-der to verify the constancy of the concentrations(± 1.0%). The solid phases collected during the ex-periment were investigated by X-ray diffraction, andby scanning electron microscopy (ISI, model N).

Results and DiscussionIn order to study the kinetics of dissolution reac-

tions of alkaline earth fluorides, it is necessary to takeinto account electroneutrality and the thermodynam-ic equilibrium constants (K) for the various associat-ed species+>.

The computations were made by successive ap-proximations for the ionic strength (I), as describedpreviously 7 using activity coefficients calculated fromthe extended form of the Debye-Hiickel equationproposed by Davies". The thermodynamic solubilityproducts, Ksp for the alkaline earth metal fluorides aregiven in Table 1, together with specific surface areas(SSA) of the seed crystals.

Concentrations of ionic species in the undersatu-rated solutions were calculated as described previ-ously. The undersatnration (0) may be expressed byEq. (1)

{([M2+]O [F-]~t3 - ([M2+][F-]2)1I3}0= [M2+]o [F-]~ ... (1)

where [M2+] and [F-] are the concentrations oflattice

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INDIAN J CHEM .• VOL. 26A. APRIL 1987

ions at time t, and [M2 + loand [F-lo those at equilibri-um, respectively. The latter values were calculatedfrom the solubility products at the ionic strengths ofthe experiments.

The results of dissolution experiments are summa-rised in Table 2 together with the initial concentr-ations of metal ion, ionic strength, the ranges of un-der-saturation (0) used in experiments and the ratesof dissolution as given by the Eq. (2)1°

Table '1-Mean Diameter and Specific Surface Area ofMetal Fluorides Seed Crystals

Mean Specific surfacediameter area

(,urn) (rn? g-I)

0.5 7.20.8 2.30.1 36.51.5 0.5

Salt

Magnesium fluorideCalcium fluorideStrontium fluorideBarium fluoride

Solubility products at 25°C

EquilibriumMgF2 (S) •••Mg2+ + 2F-CaF2 (S) •••Ca2+ + 2F-SrF2 (S) •••Sr2+ + 2F-BaF2 (S) •••Ba2+ + 2F-

K,p8.52 X 10-9

2.85 X 10-112.45 X 10-9

1.40 X 10-2

ref.33

1920

Rate=R=d(MaA,)/dt=KoSKs~V=K'Sa" ... (2)

where K' is a constant, Ko is the rate constant for dis-solution, S is a function of the initial seed surface areaKsp is the solubility product at the ionic strength of theexperiment, v is the number of ions of electrolyte andn is the effective order of reaction.

The orders of reaction (n) for each of the salts were:MgF2: 3.5 ± 0.15 (for 0.25 < 0< 0.75); CaF2,1 t 0.02 (for 0.24 < a< 0.4) and 1.93 t 0.07 (for0.075 < a< 0.15);SrF2, 1 t 0.05 (for 0.06 < a< 0.2)1.93 t 0.1 (for 0.0075 < a< 0.015); and for BaF2

1 to.2 (for 0.15 < a< 0.2) and 2 to.2 (for0.08 < a< 0.12).

The rate of dissolution of crystals in an aqueous su-spension is, in generalll-16, either controlled by sur-face processes, or by the transport of substance be-tween the volume adjacent to the dissolving surfaceand the bulk solution or by a combination of such pro-cesses. In order to compare the dissolution rates fordifferent electrolytes, the comparisons must be limit-ed to substances which dissolve by similar mechan-isms at a given temperature. The experimental disso-lution parameters should be normalized with respectto undersaturation and with the active unit surface'area. The normalized rate of dissolution (K ) may be

Table 2-Dissolution Kinetics of Alkaline Earth Fluorides at 25°Ca TM rate K K

10i (103 mol-I) (mol rnin -1m -2) (rate/ an) (m S-I)MgF2 50 1.04 6.30 x 10 - 8 1.26 x 10 - 11 4.13 x 10 - 18

60 0.83 1.18 X 10-7 1.95 X 10-11 6.40 X 10-1865 0.73 1.59 X 10-7 2.45xlO-11 8.04 X 10-1870 0.63 2.04 X 10-7 2.91 X 10-11 9.55xlO-1875 0.52 3.30 X 10-7 4.40 X 10-11 1.44 X 10-1780 0.42 6.51xlO-7 8.14xlO-1I 2.67xlO-·17

7.5101520

0.751.001.251.50

0.1780.1730.1640.154

1.4641.4601.4571.453

81012

13.96813.66513.361

Av. 3.52 X 10-11 1.15 X 10-17

8.98 X 10-8 1.20 X 10-8 4.91 X 10-153.86 X 10-7 3.86 X 10-8 1.58 X 10-140.02 X 10-7 5.35 X 10-8 2.19xlO-141.47 X 10-6 7.35 X 10-8 3.00 X 10-14

Av. 4.44 X 10-8 1.52 X 10-14

7.30x 10-7 9.73 X 10- 7 4.82 X 10-131.45 X 10-6 1.45 X 10-6 7.18 X 10-132.13 X 10-6 1.70 X 10-6 8.42 X 10-132.91 X 10-6 1.94 X 10-6 9.61 X 10-13

Av. 1.52 X 10-6 7.51xlO-13

4.78x 10-4 5.98xlO-5 3.57 X 10-117.28 X 10-4 7.28 X 10-5 4.34 X 10-111.07 X 10-3 8.92 X 10-5 5.32 X 10-11

Av. 7.39x 10-5 4.41 X 10-11

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HAMZA et al.: KINETICS OF DISSOLUTION OF ALKALINE EARTH FLUORIDES

represented by the relation K = rate/ So"; the valuesof n should also be the same within experimental er-ror. Normalization of seed surfaces is made by com-paring rates of dissolution per unit seed surface areaand expressed as mol min - 1m --2. As shown in Table 2the rates of dissolution change dramatically fromBaF2 to MgFz in the order similar to that of solubilityorderfi.e.Mgf', < CaF2 < BaF2).Thismaybebecauseof the small size of the F- relative to the M2 + ion.Many physical properties of the fluorides of the alka-line earth metals foUow this trend, thus supporting thekinetic results. The lattice energies decrease unusual-ly rapidly (Table 3) probably because the large cationsmake contact with one another without at the sametime making contact with the F- ions. Due to theabove assumption the hydrated ions on the surfacegive off some of their associated water molecules inleaving the surface to enter the solution phase. Thedissolution rate may be controlled by the partial de-hydration of the cations. Activation energies for thedehydration of cations 17.18 are given in Table 3.

The adsorption coefficients for the adsorption ofpartially dehydrated ions at the crystal surface havebeen taken into consideration 19 in order to calculatethe absolute linear rate constants, K (mol s - I). Thecalculated values of the rate constant are summarizedin Table 2. The experimental effective linear dissolu-tion rate constants K (mol s - 1 ), may be calculated bymeans of Eq. (3 )20:

Table 3- Lattice Energy and Activation Energy of Dehy-dration of Cations

MgF2 CaF,77 43

SrF2

39'BaFz

34G;110-21 JLattice energy(kcal/rnol) 695 624 588Ea (kJ/mol) 71.13 02.24 25.59

Table 4-Experimental Rate Constant (K:xp) for Dehy-dration, and Calculated (Kcalc) and Experimental (Kexp)

Dissolution RateSalt K;'r(s : )

5.0 X 104

1.6 X lOH

5.0 X lOH1.6 X 109

Kcak(ms-I)

3.8 x 10-6

9.9 X 10-5

1.6 x 10- 3

3.9 X 10-2

Kexp(ms-I)

1.2xlO-4~1.8 x 10-14

7.5 X 10-13

4.4 x 10-11

MgF2CaF2s-r,8aFz

"Rate constant for dehydration

K=KAM/p(6X107) ... (3)

In Eq. (3) P is the crystal density expressed in g ern - 3

and M represents the molecular weight in g mol- I.

Equation (3) is valid only for cubic or sphe~cal cry~-tals which dissolve uniformly" . Microscopic exanu-nation of the cubic crystals during the growth experi-ments confirmed a predominantly uniform dissolu-tion.

With the exception of magnesium fluoride, it is evi-dent from the results in Table 4 that there is a generalparaUelism between Kexp and Kcak·. .

The higher values of activation energies for dISSO-lution activation energies for release of a water mole-cule from the hydration shell and lattice energy formagnesium fluoride as compared to those of other al-kaline earth fluorides may explain the much lower ex-perimental dissolution rate of magnesium fluoride.

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