Colloids and Surfaces, 8 (1983) 103-l 19 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands

103

THE ADSORPTION OF HYDROXAMATE ON SEMI-SOLUBLE MINERALS. PART I: ADSORPTION ON BARITE, CALCITE AND BASTNAESITE

PRADIP and D.W. FUERSTENAU*

Deportment of Materiots Science and Mineral Engineering, University of Cafifornio, Berkeley, CA 94 720 (U.S. A.)

(Received 3 December 1982;acrrpted in final form 8 April 1933)

AB3TRACT

The uptake of potassium octyl hydroxamatc on baritc, calcite and baslnaesite has been determined as a function of time, reagent concentration and pH. The results of this ad- sorption study indicate a surface rcwtion/chemisorption mechanism whereby the cation& on the mineral surface form hydrowy complex- 3n solution, readsorb at the interface, and then interact with the hydroxamate. The hydroramate polar group is highty specific to rare-earlh cations as compared to alkaline-Earth cations, CaH and Ban, giving rise to very extensive and rapid uptake on bastnaerite in comparison to calcite and barite, Multilayers of adsorbed hydroxamate form on hastmwsitc and calcite.

1NTRODUCTlON

Qnc of the primafy requirements for a flotation collect.or is that it achieve u high degree of selectivity between the mh~crals being separated. There are a num&r of ore systems whcrc, as a consequence of the presence of various minerals having very similar properties, conventional collectors arc not very effective. Chclatlng agents, which can form stable complcxcs wit.11 sl)ecifir: metallic cations on the mineral surface, appear to bc promising alternative -reagents in such cases. Hydroxamates belong to this category of flotation col- lectors, and in recent years their use has been attempted on a number of ores [ 11. The Ixrcsent paper deals with tho adsorption of aqueous octyl hydroxamate, a chelating agent, on three semi-solub!e minerals, namely bar&o, calcite atlcl

bastnacsite. The latter, bast.nae&.e, a rare-earth fluocarbonatc, is one of the principal mineral sources for rare-earth met&. One of the largest deposits of bastnaesite occurs at Mountain Pass, California, as a carbonaceous ore body with alkaline-earth carbonates and sulfates as the major ganguo. Hydroxamates were found to be quite selective in the flotation separation of bastnaessite from the associated gangue minerals, mainly calcite and barite [ 2). The l>resent investigation was undertaken in order to understand the mechanism of hydroxamatc interrtctlon with these minerals.

*To whom all correspondence should be addressed.

0166.6622/831$03.00 Q 1983 Elsevicr Science Publishers B.V.

HYDROXAMIC ACIDS AND THEIR METAL COMJ’LRXES

Iiydroxamic acids exist in two tautomeric forms [3] :

0 OH R C/

1) Hydroxyamide, and 2) Hydroxyoxime, R- II H-N

‘OH Nk,lH

Structure 1 has one rcplaccable hydrogen atom and would behave as a mono- basic acid while structure 2 has two replaceable hydrogens. That hydroxamic acid forms metal complexes through the hydroxyamido functional group 1 and not through the hydroxyoxime structure 2 has been proved by 1R studies and UV speckal investigations [ 31. Complexes are formecl via the substitu- tion of a hydrogen atom of hydroxyamide by a metal cation and ring cIosure via t.he carbonyl oxygen atom [4 J :

‘ihc pKa of the usud hydroxamic acids, such as aceto and benzo hydroxamtc acid, is close to 9. Most of the reported vdues of the pK of hydroxamk acids are in the range of 7-9 [3]. The ovedl stability conslants of several metals have been evaluated by Bjerrum’s technique, pH titration, distribution, and potentiometric methods. Schwarzenbach [ 5) has reported on the values of stability constants of several metal hydroxamates (chains containing two carbon atoms). Recently Dutta and Scshadiri 16) and Liu aud Sun [7] deter- mined the stability constants for various rare-earth elements when reacted with bcnzohydroxamic acids. Agrawal et al. [ 8 J us~4 phenylhydroxamic acid to synthesize rnebl comptcxes of La, Cc, Pr, Nd, etc, The complexes were found to be fairly soluble in ethanol and of general composition hSRl for rare-earth elements, where R represents thr! hydroxamati ligand. Infrared study of the complexes registered no band around 3260 cm”, indicating t,hc absence of an OH group in the complex molecules, and the l*(C+O) stretch which occurs at around 1660 cm’ I in the uncomplexed hydroxamic acids wns shifted to lS90 cm” I. Their results strongly suggest that the coordination of the l&and is through carbonyl oxygen and nitrogen. Barium hydroxamati has also been synthesized by the same procedure 191, The stability constants for the compIe~es of octyl hydroxamatc with calcium, barium and rare-ear&h clcmcnts arc not available in the literature but, based on the studies thus far reported, it is reasonable to assume that they are of the same order of mag- nitude as aceto, benzo and phenyl hydroxamates of these cations.

According to Bogdanov et al. [lo, 111, hydroxamic acids form comylexes of different strength depending upon the metal cation position in the periodic table. The weakest complexes are those formed wit-h alkaline-earth metal

105

cations (Ca, Ba and others), whereas complexes formed with the transition elements (Nb, Ti, V, Mn, Zr, Hf. Ta) have somewhat greater stability. Rather strong complexes are formed with rare-earth elements and aluminum, as they are highly charged cations. The strongest complexes arc formed with Fe’*, and probably with Ta’* and NbS+.

HYDROLYSlS OF METAL JONS

Metal ions undergo hydrolysis forming a number of hydroxy complexes in aqueous solutions. lf the reaction scheme for hydroxide formation can be set. out as shown below for a trivalent rare-earth metal ion, Ml+,

M3+ * MOH’+ * M(OH)z+ * M(OH),* * M(OH)4- C= hl(OIi)52-... . . . 1C Jr

M2(0H),“’ 11

Ihl(OH),* n~~,Ol mjid

hlp(oIQpp-q)+ IM,O~~mll~O,,&

For the reaction

MS+ + i H20 G+ hl(OIi)i(s”)’ + i H+

the strtbility constant *& can be exlrrcssed as

The solubility product of the corresponding hydroxides can be tvritten as

[hl”‘] *KS = -

[H’]”

Table 1 summarizes the values of the foregoing constants taken from Kragtcn 1121.

TABLE 1

Stability condan& for the fiydrowy complexes of .;ome raw-earth and alkatine-carth metal ions IlZJ

__-

Metal Ion lo&t 8, log 8* lok! 0, lot3 P. log ‘K, Additional complexes

Lam -8.6 -17.2 -26.9 -36.9 19.7 B = .* 71.2 Ce” -8.1 -16.3 - 26.0 -38.0 20.1 81, = -32.8 Ca” -12.6 23.0 SP -13.4 24 .o Ba* -13.2 24.0

EXPRRIMRNTAL MATERIALS AND h~l3’lWODS

Baritc samples from Kings Creek, South Carolina, and calcite samples from Kansas, were obtained through Ward’s Natural Science Establishment., New York. Thcsc samples were washed in triply distilled water and ground to the desired size. A handpicked sample of pure bastnaesite ore from Mountain Pass, California, was obtained from the Molybdenum Corporation of America. The “pure” bastnaesite grains, sorted in unfiltered UV light (bast-naesitc fluor- esces with a characteristic green color) were g-round to minus 400-mesh size for the adsorption studies. The analyzed poM*der yielded 57.4% REO (rarc- earth oxide) as compared to 76% REO for pure bastnacsitc. The other impurities were 8.8% BaO, 1.6% CaO and 0,4A SrO, indicating the presence of alkaline-earth carbonates in the sample. Because of the very fine grain size of the natural mineral, it was impossible to obtain natural bastnacsik of higher purity, The surface areas of Lhe ground mineral samples, as mez,surecl by the BET adsorption met-hod using nitrogen in a hlicromcritics Orr Surface Area/Pore Volume Analyzer, wcrc found to be 1.73, 2.15 and 2.4 m’/g for the bastnaesito, barih atild calcite powders, respcctivcly.

All the rcngents used in the study wcrc analytical grade chcmlcal reagents. The hydroxnmatc collector was synthcsizcd in the laboratory by the well known method of preparation by reacting the methyl ester (methyl octanoatc) with hydroxylaminc hydrochloride and KOIi in t-he prescncr! of mct.hyl alcohol I13], The hydroxamate thus prepared was purifiecl by recrystallization from hot nwthyl alcohol solution a number of times.

‘The synthcsizcd product was chemically analyzed and characterized with infrared spectroscopy. The chemical analysis suggcstcd that the product is 50M potwium octyl hydroxamatc? salt., the rest being the corresponding acid [ 141. For adsoq,tion studies, aqueous hydroxamate solutions were annlyzccl sl~eotroptlotoa~ctric~ly by the well known ferric hydroxamatc method [ 15,161. To 5 cc of hydroxamaLe solution, 10 cm3 of ferric perchlorate solu- tion was added and the absorbance of the purple-colored complex was deter- mined wit.h a Gary 17 sllcctrollhotollleter at, 510 nm, with ferric perchlorate solulion and water in the same ratio, as a standard reference. The ferric per- chlorate solution used in this analytical procedure was prcparcd in the laboratory [ 15). To start, 1.G g of F&l, was mixed with 6 cm3 of 60% per- chl~rir: acid and the mbture was evallorated to dryness. Artcr extract,ing the rcsiduc with 100 cc nf water, 10 cm3 of this extract was transferred to a 500 cm’ volumetric flask. Ethanol and 60% pcrchloric acid were added alternately until 35 cm3 of percl~loric acid had ber?n added. The solution was then made up to 500 cm’ with etlian~l. In making a calibration curve, ab- sorbance at 610 mn was found to increase linearly with concentration until about 2 X 1W3 MI above which the rise in absorbance is no longer linear. Since measurements art? accurate for solutions up to about 6 X fW3 1’11 con- centration, the absorbance can be measured above this concentration only after dilution. The lower limit due to the sensit-ivity of this method is lo-* M.

107

The validity of this procedure was tested under various experimental con- ditions. Blank tests were run on all the three minerals used in the study. Bar&e, calcite, and bastnaesite supcrnatants did not show any absorbance with ferric perchlorate. Hydroxamatc solutions aged in glass and polyethylene bott.lcs evon after 10 days did not show any change in absorbance values with ferric perchlorate, indicating no appreciable adsorption on glass and poly- &hylene. Furthermore, the calibration curve is reproducible over the whole pH range 3-10. This impiies that the adsorption experiments carried out at different plls need not be buffered for analysis. Perhaps, perchlorate solu- tions act as a buffer since a purple-coIored complex is obtained all through these experiments and this is only stable at low pHs (1 to 2). The complex formation reaction with iron in the bulk have been discus& by Haghavan and Fuerstenau [ 151. At higher pll ranges, 1 : 2 and 1 : 3 complexes are formed and hence the peak in the absorbance shifts to lower wavelengths [ 17 ] .

The geometrical area per molecule for potassium octyl hydroxamatc was calculated to be 20.6 AZ considering the end group only and 55 A2 including the chain (lying down flat position).

Adsorptiorl cxperimen ts

The adsorption of potassium octyl hydroxamnte under various conditions : was evaluated by determining the difference in the concentration of the sur- factant in solution before and aftir adding the solid mineral powders. The experiments were conduct& in 50 cm3 polyethylene bottles at a solid : liquid ratio (by weight) of 1 : 20 for barite, 1: 10 for calcite and 1 : 100 for bastnaesite, unless otherwise indicated. The conditioning (adsorption) times were selected for each mineral system according to the kinetics of adsorption for the different minerals. Conditioning was carried out in polyethylene hottles that WCM rotated in a shaker for the predetermined t.imcs required for attaining equilibrium. After the conditioning operation, to separate the liquid from the solids, t.hc supcmatant was centrifuged twice in a SORVALL supcrspeed centrifuge for 20 min at 12,500 rpm. The concentration of hydroxamatc in solution was then measured by the spectrophotomet.ric method described in the previous section. The plI adjustmenti were made with NaOli and HN03 additions.

Kinetics o/adsorption

Figure 1 shows the results of the uptake of K-octyl hydroxamate on barite, calcite and bastnaasib as a function of time. The initial concentration was 10-3M in all three cases. Note, that barite being the least soluble mineral, takes as long as 10 days to reach equilibrium whereas the more soluble calcite reaches equilibrium in about 30 h. Bastnaesite exhibits bot.h the greatest ad-

Fig. 1. The kinetics of hydroxamle {HXm) ad-orption on baritc, calcite and bastnacssite.

sorption and fastest approach to equilibrium, which may be related to its solubility. For subsequent experiments, an attempt was ma& to use 8 con- ditioning time that would achieve equilibrium in the system, but for baritc 85 h was assumed to br! sufficient for equilibkam.

The equilibrium adsorption density on bastnacsitu is pIott& &s a function of the equilibrium concentration of hydroxamati in solution in Pig. 2. The pll in this cab was 0.3 f 0.6, whfch also is the natural pH of b&nae&e In aqueous medium. The adsorption isotherm pm cs the typical S-type shape. l’he monolayer adsorption density irf 8.1 ymoljm2, assuming the cross- sectional arca of the hydroxamab group to be 20.6 Al, is also shown in the figure. It appears that the monolayer of hydroxamatc on bastnacsite is ob- tained at much lower concentration than the range of concentration that could be used in this study and that multilayer physical adsorption must occur on tne initial chemisorbcd hydroxamate layer. The plateau at 60 pm jm* indicates that as ~;~any as six layers of hydroxamate may exist at the bast- nacsitc/watct interface.

Figure 3 shows similar adsorption isotherms for bariti and calcite, together with bastnacsitc for comparison (the equilibrium pH for all three mineraIs WBS

.--- VtCITftU MOMNAY ECI I---__--------_ _ _ 2o5 ;2,__

1 1 I I I I I I I I

IO 20 30 40 EOUILIBFIIUM CONCENTRATION, mol&lmtaIQ4

so

109

Ffg. 2. Adsorption isotherm for bartnae~ite/hydroxamate system at 20% and pH 9.3 * 0.1.

:40 z G

B 0 O A

0 5 k 20 8

8 kTr.rL-1 08 0

0 0

---1-- VERTICAL ----c-c--

0 y--- I t I I 0 10 20 J 30 40 SCI

I ~QUILIBflUM CONCENTRAYION. mde/liWr a ;6* * -_

Fig. 3. Hydroxamate adsorption isotherms for barite, calcite, and bostnaesite at 21°C and pH 9.3 i 0.1.

around 9.4 f 0.1). Clearly, hydroxamate adsorption on barite and calcite is significantly lower than on bastnaesite (except for the last calcite point). Adsorption on bar&e seems to approach a monolayer corresponding to a con- figuration as If all chains are lying flat on the surface. In this position; the

110

geometrical area per adsorbate molecule is 65 A*, corresponding to sn adsorp- tion density of 3 pm/m*. The relatively weaker interaction of hydroxamato with barium, perhaps, requires a higher driving force; that is, higher concentra- tions are required to achieve the close-packed monolayer tit.h all the molecules being vertically oriented, This observation is certainly consistent with tho minimal flotation response of ba&e to hydroxamatc collectors 11.2~41~

Calcite behaves in a strange fashion with respect to adsorption, similar to that observed In flotation experiments as reported earlier t21, There arc two distinct discontinuit-ies in the adsorption isotherm corresponding; to two possible monolayer dcnsitics of hydroxamate on the surface. The first occurs at 3 ~mol/m’, indicating a horizontal orientation as discussed for barito, With an increase in t-he concentration of the adsorbate iu soiution, the next plateau tends to occur at 8 pmol/m2 when the area occupfed by hydroxamate

10 -

I 1 t I. I 2 4 6

Pi IO I2 I4

Fig. 4. Effect of yU on the adsmption density (top) and permnt uptakb (bottom) of hydroxamate eo~lector on barite at 8 x lW*Af and 1O”Af initial concentrations.

111

is 20.5 Af per molecule, indicating a vertically oriented monolayer. Instead of an expected plateau with further increase in concentration, a linear rise in the uptake of hydroxamate from solution is observed. This is, perhaps, due to the precipitation of calcium hydroxamate in the bulk or alternatively on the surface. Because of t.he inherent limitation of the Whnique of measuring adsorption by difference in the bulk concentration, it is not easy to differ- entiate between these two postulates. No visible 6igns of such precipitation in bulk, however, were observed during the experiments. The two plateaus corresponding to a change in conformation of the adsorbed hydroxamate species were further borne out in experiments carried out at higher tempera- tures [ 141.

Effect of 311 on adsorption Figure 4 shows the adsorption behavior of hydroxamate on barite as a

function of pH for two selected values of the initial hydroxamate concentra-

1 I I 1 1 I

Fig. 6. Effect of pH on the rrdcorplion density (top) and percent uptake (bottom) of hydroxamnte collector on calcite at 0 x 10’. AI and IO’ ‘AI inith! cowenlrat ions:

112

tion. There is a distinct peak around pH 85-9 in tie adsorption curve and a change irl pH on either side of this peak results in a decrease in adsorption. This is very similar to the flotation behavior of barite in t.he presence of hydroxamate collector 121. The drop in adsorption density (both as a perct-ntagc of the added material or in terms of the adsorbed hydroxamate at th surface) is more or less symmetric but less sharp than the results of flotacian tests on the same mineral [ 2 1. Figure 8 summarizes the results of some experiments at the cakite/water interface. As was observed In earlier not&ion experiments [2], the adso@ion of hydroxamatc on calcite seems to increase at lower pHs, besides having a peak at the charactetistic pIJ 9. The results for bastnaesite are shown In Fig. 6 which differ from the &sults for the other two minerals. There is a drop in adsorption density at pH 8 for 6 X 10e4 M initial concentration and at p1J 7.6 for 1O”Af init.ial hydroxamate concentration, The experiments are difficult to carry out at pHs below 7 because of the basic nature of both calcite and bastnaesitc.

I I I I I N

BASTMESTE 4574 9b RIO)

I I I I 1 6 7 8

P’H 10 II 12

Fig. 6. Effect of ~1% on the adsorption dencity (lop) and percent uplake (bottom) of hydroxamate cotleclor on bastnaesile at 6 x fO-*AI and 1O’JAl initial concentrationr.

114

position. A hydroxylation can result in the formation of a hydroxylated metal ion which is pulled out of its lattice position, as designated by moving

Hydroxamates can interact dm 11arly, it to the right of the vertical dotted line, as illustrated with these equations.

(Me*) t + Hco;+

(MOU) i

i (nleHco,)+

+ HXm 3 i (MeXm)+ + H’

$MeOH)+ + HXm -+ i (MeXm)’ + Ha0

(Me”) i + Xm’ -+ i (MeXm)+

i(hleOH)+ + Xm- + i (MeXm)’ + OH-

BastnaesiW contains lantianum and cerium as its major cation constituents. The a&cm&ion results indicate a plateau below pH 9, with t-ha adsorption

it3 _

I sharply decreasing at highr pHs. Figures 7 a&d 8 show the conccn-

/

/- IO-%& Ce

7 9 9 IO II 12 13 6

Fig. 7. Aqueous solution ?H

equilibria for Cs at IO-‘Al total solution concttntmtion.

115

t-

?-

Fig. 6. Aqueous solution equilibria for La at 1 O-a AI total aolutlon cmcentration.

tration of various solution species for cerium and lanthanum, respectively. The diagram is based on data compiled by Kragten [ 121 (see Table 1).

At lWsbf total concentration, in the range of the observed pfatmu, &I*, M(OH): and MOW predominate in the bulk. It appears that these species are responsible for hydroxamate adsorption. In the alkaline pH range, a sharp drop in the concentration of these hydroxy complexes is observed, which also coincides with the drop in ‘adsorption. The uncharged hydroxylated species Ce(OlI), or La(OI& may not assist flotation and adsorption and neither do the negatively charged Ce(OH)i and La(Oll)i+ The reactions at the inter+ face can be illustrated schematically as follows if Xm represents the hydrox- amate ligand and CeXma (surf) the cerium chelate adsorbed at the surface:

Ce(OH): (surf) c 3 HXm = Ce XmS(surf) + 2 Ha0 + H*

Ce OH* (surf) + 3 HXm = Ce Xrna (surf) + Ha0 + 2H+

Ce OH+, (surf) + 3 Xm’ = CeXm, (surf) + OH’

Ce(OH)~ (surf) + 3 Xm’ = CcXm, (surf) + 2 OH’

11G

Similar reactions can also be written for other rare-earth cations, especially La* which is also abundant in bastnaessite. All the above-mentioned reactions are possible depending upon the pIi of the solution. In the alkaline pIi range, hydroxamate anions will predominate whereas at acidic pHs neutral hydrox- amic acid molecules are expected to play an important role. The trihydrox- am& complexes of ccrium and lanthanum have been reported in the literature 161.

The formation of drihydroxamato complexes by reaction with neutral hydroxamic species releases Ii+ ions, causing the pH to decrease as a conse- quence of adsorption. A not-&able decrease in pII after adsorption was obsetverl at higher temperatures [ 141, which also substantiates the fomguing adsorption mechanism. Itaghavan and IQerstenau [ 16) also observed a decrease in pII after adsorption in t-hc hcmat.ite/hydroxamat.e system, further confirming the role of the undissociatid hydroxamic acid species In adsorp- tion.

Multilaycr adsorption, once the first chcmisorhcd layer has been com- ]&ted, may occur through hydrogen bonding as shown in (i) below or more likely may result from hydrophobk bmmling through flat-lying hydrocarbon chains (ii), as su~csted earlier by Raghavan and Fuerstenau (161 for t-ho hcmaiite[hydroxamatc system.

CCOy 0:: c --A _,..._.. A- = c 0

I I (ii)

----O--N---H 14 - h - OH

In the case of bastnncsite, due to the very high specific free energies of ad- sorpliotl, there is conslderabb multilayer adsorption.

‘I‘he rr?action mechanism for adsorption on barite and calcite is expected to bc the same as that suggested for bastnacssite with an important difference that Ra+’ and Ca- are divalent ions. There are, at most, only two hydroxamic acid molecules which may he interacting with one Ba* or Ca” site. Incidcntal- ly, this may be another reason why tho adsorption on the trivalent rare-earth fluocarl,onatc (bastnacsitc) at room tamperat-ure is many fold higher than on divalent atkaline-earth carbonates and sulfatis (calcite and bariCo),

In the case of baritc, t-he monolayer of hydroxamate corresponds to a cotlfonnntion of the kind where the adsorbed collector species lie parallel to the solid surface, From an energy standpoint, t.hc most favorable condition may be having coiled hydrocarbon chains lying flat on the surracc, at least initially at lower concentraCions. Once the free energy of adsorption is high cnuugll so that the chains can be oriented verticalty to make space for the adsorbing collectors, the vertical orientation can be a&sled by hydrophobic bonding between the chains. This is what is observed with calcite at higher concentrations but apparently the interactions of the hydroxamate head

117

group with the Ba* sites are not strong enough to do the same in the case of barite at the concentrations used in this study. Plitt and Kim [ 25) ob- served similar results with oleate adsorption on barite.

The sharp peak observed for barite in the adsorption experiments is not easily explainable, The limited data that are available for bulk Ba’, aqueous chemistry indicate that at pH 9, the &OH+ concentration is negligible. Perhaps in the presence of CO,, some other barium species is formed and is responsible for hydroxamato adsorption.

The characteristic peak at pH 9 for calcite was also olwenwd by Bogdanov et al, for fluorite ilO], The solution equilibria for Ca” indicate that the CaOW peak occurs only at pi1 13 or so, in the absence of CO2 [ 23 ] . How- ever, CO2 plays a significant role in controlling the aqueous chemistry of calcite [ 26 J. The presence of such species as CaHCO; bar; been reported and possibly complex ions of this type may participate in adsorption on calcite. An increase in adsorption and flotation at. acidic pHs tends to substantiate this claim since the CaHCOf concentration increases at lower pHs. An invcsti- gation by Sanpatkumar et al. 1271 on i he effect of CO, in calcite flotation indicated that the improved flotation ~ecovcry of calcite is due to the adsorp- tion of species CaOH+, CaHCO;, etc. These species are produced in the presence of CO2 and thus a positive charge on the surr’ncc of calcite may bc enhanced by a CO2 environment.

It is also apparent from these adsorption results that hydroxamatc collec- tors are very specific to rare-earth cations as compared to alkaline-earth cations, Ba” and Ca +‘. The adsorption is ahnost 6 to 10 t-imcs more on bastnaesite as compared to barite and calcite, this dearly being the reason why our flotation results reported in ~11. earlier paper indicate the highly selective nature of hydroxamatc reagents in bastnaesite flotation [ l]*

Free energy of udsor/ltim

Information about the thermodynamic aspects of collector adsorption can be obtained from the experimentally measured adsorption isotherms. Specifically adsorbing collcctom will be closer to thl? surface, that is adsorp- tion takes place in the Stern plane. The free energy of adsorption, A<J& can bc calculated using the following form of the Stcrn-Grahame equation [ZS] :

9 = & cxp(-AC&,/W”)

1-O .

where @ is the fraction of Stern plane sites covered with adsorbed hydrox- amates, C is the equilibrium concentration of the collector in solution, II is the gas constant, and 7’ is the temperature. To use this equation for estima- tion, our results on the adsorption of hydroxamate were here plotted on log-log paper to determine the solution concentration at which 60 percent (lJ3 = 0.5) coverage occurs.

118

For the btitelwater, calcite/water and bastnaesiteiwater systems at room temperature, our resukts show that the free energies of adsorption of hydroxamate are - 26, -28 and - 57 kJ/mol, respectively. The order of these values is consistent with all our observations, that is hydroxamate has the greatest affinity for bastnaesite and the lowest for barite.

SUMMARY AND CONCLUSIONS

The adsorption of potassium octyl hydroxamate on barite, calcite and bastnaesitc has been determind as a fun&ion of time, concentration and PH. The results of this adsorption study indicate a surface reaction (chemi- sorpt.ion) mechanism whereby the cations on the mineral surface form hydroxy complexes in solution, readsorb at the interface, and then interact with the hydroxamate, The hydroxamate polar group is highly specMic tu rare-earth cations as compared to alkaline-earth cations, Ca’, and Da”, givjng rise to vaw extensive uptake on bastnacsite in comparison to calcite and barite, The standard free energy of adsorption has been estimated to be -26, -28 and --57 kJ/mol for hydroxamate adsorption on barite, calcite and bastnacsik, respectively,

ACKNOWLEDOEMENTS

\Vc wish to acknowledge the financial support provided initially by the Xtolybdcnum Corporation of America, Mountain Pass Operations and then by the National Science California Regents for a of his program.

Foundation, Yradlp is grateful to the University of fellowship grant awarded to him during the course

1

2

3 4 5 6

;: 9 10

11

12

Prsdip and D.W. Fucrstcnau, Mineral Flotation with Hydroxamate Colleclors, AtME Annual Meeting, Dallas, 1982. U.W. Fuerstenau, Pradip, L,A, Khan and S. R8ghauan. An Alternate Reagent Scheme for Ihe Ptotation oIMounlain PaM Rare-Estth Or@, Xl11 Int. Mineral Processfng Congrrar, Toronto, 1982. El. ChmttcrJec, Coord. Chcm. Rev., 26 (1978) 281. Y.K. Agrawal, Russ. Cham. Rev., -18 (1979) 948. G. Schwancnbach and K. Schwarzenharh. Helv. Chim. Acta, 46 (1963) J 390. N.K. Dutta and T. ScshadM, J. Inorg. Nucl. Chcm., 31 (1975) 3337. C.Y_ Liu and P.J. Sun, J. Chin. Chem. Sot., 22 (1975) 317. Y.K. Agrawal and T.P.Shukls, J. Ind. Chem. SOP., 61 (1974) 373. Y .K. Agrawal and 14-P. Kapoor, J. Ind. Chcm. Sm., 51 (1976) 174. OS. Bogdanov ct al., In Ii lntcrnationsl Mineral Processing Congress, Procccdlngs, London, P8$M 7,1973. OS. flogdenov ef al.. in: XII International Mineral Procslsrtng Congress Proceedings, S8o Paula, Brazil, 1978. J. Kragtm, Atlas of hfeta&Ligand Equilibrium in Aqueous SoIutions. John Witty and Sons, New York, 1978.

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13 H.D. Peterson, MC. Fuerstenau, R.S. Rickard and J.D. Miller, Trans. AIME, 232 (1965) 338.

14 Pradip, Ph.D. Ditsertation, University of California, 1981. 16 S. Raghaven and D.W. Fuerstenau, J. Colloid Interface Sci., 69 (1975) 319. 16 J.B. Neilands, Struck Bonding, 1 (1966) 69. 17 S. Seifter, P.M. Gallop, Susan Michaek and E. Meilman, J, Biol. Cham.. 235 (1960)

2613. 13 R. Natarajan and D.W. Fuerrtensu, Int. J. Miner. Process., in prets. 19 J. Le Normand, R.H, Yoon and T, Salman, presented at the 14th Annual CIM Con-

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