Mg2+ removal in the system Mg2+—amorphous SiO2—H2O by adsorption and Mg-hydroxysilicate precipitation

Mg2+ removal in the system Mg2+—amorphous SiO2—H2O by adsorption and Mg-hydroxysilicate precipitation

C;n,...
Download PDF
2MB Sizes 0 Downloads 23 Views
Report

C;n,
Mg2+ removal in the system Mg2+ -amorphous Si02-H20 and Mg-hydroxysilicate precipitation

by adsorption

DOUGLAS B. KENT* and MIRIAM KASTNER Scripps

Institution

of Oceanography, (Received

May

University

of California,

3, 1984: accepred

in revised./imn

San Diego. La Jolla. California Y2OY3 Junuur~~ 25. 198.5)

Abstract-Two chemical processes can remove Mg’+ from suspensions containing amorphous silica (amSiOJ at low temperatures: adsorption and precipitation of a Mg-hydroxysilicate resembling sepiolite. Mg*+ removal from am-SiOz suspensions was investigated, and the relative role of the two removal processes evaluated, as a function of: pH, ionic strength, Mg*+ concentration, and temperature. The extent of Mg*+ adsorption onto am-SiO, decreases with increasing NaCl concentration due to displacement of Mg2+ by Na+. At NaCl concentrations of 0.05 M and above. adsorption occurs only at pH values above 8.5, where rapid dissolution of am-SiOz gives rise to high concentrations of dissolved silica, resulting in supersaturation with respect to sepiolite. Removal of Mg2+, at concentrations of 40 to 650 PM, from am-SiOz suspensions in 0.70 M NaCl at 25°C occurs at pH 9.0 and above. Experiments show that under these conditions adsorption and Mg-hydroxysilicate precipitation remove Mg” at similar rates. For 0.05 M Mg*+, at 0.70 M ionic strength and 25”C, measurable Mg’+ removal occurs down to cu. pH 7.5 but is primarily due to Mg-hydroxysilicate precipitation. For the same solution conditions at 5”C, Mg2+ removal occurs above pH 8.0 and is primarily due to adsorption. Assuming that increasing pressure does not greatly enhance adsorption, Mg” adsorption onto am-Si02 is an insignificant process in sea water. The surface charge of pristine am-Si02 in sea water is primarily controlled by interactions with Na+. The principal reaction between Mg” and am-SiO, in marine sediments is sepiolite precipitation. The aae distribution of seniolite in siliceous uelatic . - sediments is influenced by temperatures of bottom waters and by geothermal gradients. INTRODUCTION Mg2+ IS ONE OF THE PRINCIPAL dissolved

constituents in natural water systems, and the silica phases-Xray amorphous silica (am-SiO*) and quartz-are important constituents of suspended solids and sediments. The two significant low-temperature chemical reactions that occur in systems containing dissolved Mg*+ and solid silica are: (1) Mg’+ adsorption, and (2) Mg-hydroxysilicate precipitation. Adsorption onto suspended solids, such as biogenic am&O2 particles, is an important process in the transport and removal of trace metals in the ocean (BALISTRIERI et al., 198 1 and references therein). Previous studies indicate that Mg*’ adsorption onto silica in the pH range of sea water could be extensive (VUCETA, 1976; GISLER, 198 1). If adsorption of Mg2+ on am-Si02 is extensive in sea water, then: (I) adsorption onto biogenic am&O2 could be a significant process in the marine geochemical cycle of Mg*+, (2) the electrical properties of the am-SiO,-sea water interface, which influence the adsorption properties of am-Si02, could be controlled by Mg2+ adsorption (see e.g. DAVIS et al., 1978; DAVIS and LECKIE, 1978; BALISTRIERI and MURRAY, 198 I), and (3) coverage of a large fraction of surface sites by Mg*+ could inhibit the adsorption of trace metals onto am-Si02 (e.g. BALISTRIERI and MURRAY, 1982). Several studies have suggested the importance of * Current Address: Environmental Engineering and Science, Department of Civil Engineering, Stanford University, Stanford, California 94305

chemical reactions between Mg2+ and am-SiOz in sedimentary environments. In some saline lake environments, an Al-poor Mg-smectite replaces diatom frustules (BADAUT and RISACHER, 1983). The fibrous clay minerals sepiolite (Mg2(H4Si04)3(OH)4) and palygorskite (Mg2(Al,Fe)2SiB020(OH)2 8H20) (COUWRE, 1977) grow directly on biogenic am-Si02 particles in some marine sediments (COUTURE. 1977). DONNELLY and MERRILL (1977) reported a correlation between Mg/AI and Si/AI in carbonate-containing siliceous sediments at Deep Sea Drilling Project (DSDP) Site 29, in the Venezuelan Basin, and Site 42, in the equatorial Pacific. No Mg-containing phase could be found; the authors suggested that Mg2+ adsorption onto am-Si02 was responsible for the observed correlation. KASTNER and GIESKES (1983). however. suggested that this correlation could be due to the presence of poorly crystalline sepiolite. The conversion of am-Si02 (= opal-A, JONES and S&NIT, 1971) to quartz in marine sediments often occurs via the disordered intermediate phase opalCT (KASTNER, 1981 and references therein). The experimental work of KASTNER (‘I al. (1977) and KASTNER and GIESKES (I 983) demonstrated that the rate of transformation of opal-A to opal-CT is greatly enhanced by the formation of a compound that contains Mg’+ and OH- in a I:2 molar ratio. No Mg-containing compound could be detected by Xray diffraction (KASTNER et ul.. 1977; KAS~NER and GIESKES, 1983). and brucite (Mg(OH),) solubility was not exceeded in most of the experiments (KENT, 1983). I123

1124

D. B. Kent and M. Kastner

The objective of this experimental investigation is to determine the importance of Mg’+ adsorption and of Mg-hydroxysilicate precipi~tion in sea water and marine sediments. Mg2+ removal from am-Si02 suspensions was investigated as a function of pH, ionic strength, Mg2+ concentration, and temperature. It was determined whether adsorption on am-Si02 and/ or Mg-hydroxysi~icate precipitation was responsible for the observed removal of Mg*+. The results constrain the extent to which Mg2+ adsorbs onto amSiOz in the marine environment.

dissolved silica concentrations in excess of 100 gM at temperatures of 25’C and below (DREVER, 1974: HEMLFY ef al., 1977), the Mg-hydroxysjl~cate that pr~ip~~~ under these conditions is the fibrous clay mineral sepiolite (WOLLAST et al., 1968 and references therein; COUTURE. 1977). There are two disparate estimates of the solubility of sepiolite at 25°C. CHRIST et nl. (1973) extrapolated solubility measurements made at 51 “C, 70°C. and 90°C for a well crystallized natural sepiolite. Their equilibrium constant for the reaction:

Previous investigations

is presented in Table 1.WOLLASTc’t al. ( I968)and ~‘OU’IIJRL (1977) obtained solubiiities for sepiolite precipitated from homogeneous. sea water-Iike solutions at 25°C that are simitar to each other, but higher than that extrapolated by CHRIST d al. (1973) (Table if. The higher solubiiity determined by the latter two authors could be due to insufficient aging of the precipitates, a process that requires several years at 25°C (COUTURE, 1977). As part of this work. we found

Mg-hydroxysili~te precipitates readily from alkaline solutions at temperatures near 2Y’C (WOLLAST et al., 1968 and references therein; HURD, 1973; SANTSCMI and SCHINDLER, 1974; CULBER~ON et al., 1975; COUTURE, 1977; GISLER, 1981). Although talc (Mg3S&010(OH)2) is the ther-

modynamically

stable Mg-silicate in aqueous media with

2Mg*’ + 3H,Si04 + 4HZO = Mgz(H4Si0.&(OH)., t 4H’

Table 1. Equilibrium constants and correspondingexpressions used in this study.

species

log Q'@) 25'C 5'C

Expression

Solution species

Iissio*-

-9.49(b) -22.4fb)

MgOH+

-9.90(c) -_

-12.O7(b)

-12.9(d)

Mglinsio.+

0.64@)

_-

M#(H,SiOs):,

3.82(e)

__

-7.83(f)

__

-17.41(f)

__

Adsorbed Species I SiOMSf

[%@+I {f SiOHd * ((5 SiO)& 1

(" SiO),MgO

[Ms~+I (r SiOH)' Solids am-sio,

IH*SiO*l

-2.82(')

Brucite Ms(OH)l

I!!aQ

17.4(b)

L(i)

17.0(j)

18.6(J)

19.7(k)

21.3(k)

a;+

JMg'+]' [H&Ok]' Sepiolite * ng,(H*SiO~)~(ON)U aH+ ,, 1,

(a) TerminoloSy after Sacs end lLsrr (1976). [I denote concentration in moles da-'; (1 denote surface concentrationin moles kg-' SiO,.

fbf

camsteats from Baes and Mesrser(1976); uH from Haned and Oven (1958).

Cc)

Constant

from Busey and Mesmer (1977); yH from Hsrned and Oven

(1958).

Cd) Calculated from theory of HelSeson (1967) wing date In Helgesan (1969). fe)

From Santschi and Scbindler (1974), for 1 M NsClOr, 25°C.

(f)

From Gisler (l981), for 1 I N&10,,, 25'C.

(8) Equals 1500 j&, which is in middle of several reported values for slxnilarmedia.: 1290 rtw(1 M N&IO,. JorNenoen. 1968): 1460 #Z (0.70 M NaC1, Wlrth, 1980); 1550 UH (0.70 X NsCl. Hurd, 1973); 1800 PM (0.70 M NaCl, Msrshall and Warskowski, 1960). (h) Constant from Hostetler (1963); yns from Harmed and Own

(1958).

(i) Tenpetsture effect ss calculated by Helgeson (1969). yNS from Hsrned end Owen (1958). (3) From Christ S'C.

SSSUBe

s &. (1973). T 1 fr,) not funct%

frontHsmed and Owen (1958). For of temperature.

(k) From Wollsst et al. (1968). in sea water. For 5'C value, assumed ssme temperatGe?ffect SB reported by Christ st. (1973).

(II

1125

Mg adsorption and precipitation that Mg-hydroxysilicate precipitation occurred in experiments where the solutions were undersaturated with respect to the sepiolite solubility of WOLLAST et al. (1968). Therefore, unless otherwise stated, Mg-hydroxysilicate saturations have been determined relative to the extrapolated value of CHRIST et al. (1973) (Table 1). Adsorption of Mgzf onto quartz at low ionic strength, viz. I X 10m3M NaCl, occurs at pH values as low as 6.0 (VUCETA, 1976). Near pH 6.0, even at dissolved silica concentrations as high as the solubility of am-SiOz, Mg’+containing suspensions are undersaturated with respect to sepiolite. Thus, the Mg*+ removal that VUCETA (1976) observed near pH 6.0 is due to adsorption and not Mghydroxysilicate precipitation. Mg*+ removal from am-SiOz suspensions at high ionic strength occurs at pH values above neutral (GISLER, 198 l), where the solubility and dissolution rate of am-SiOz are high (GREENBERGand PRICE, 1957; STUMMet al., 1967; WIRTH and GIESKES, 1979; KENT, 1983) and the solubility of sepiolite is low. High concentrations of dissolved silica, hence supersaturation with respect to sepiolite, are unavoidable under these conditions. GISLER(198I), who determined the amount of Mg2+ removed from dense suspensions of am-SiO, in 1.0 M NaClO.,, used an equilibrium approach to separate the effects of Mg2+ adsorption and Mg-hydroxysilicate precipitation. The suspensions were allowed to react for 18 hours. Dissolved silica concentrations were not measured; the suspensions were assumed to be in equilibrium with am-SiOr. The precipitation reaction was assumed to obey: Mg2+ + H4Si04 = MgH,SiO,,, + 2H+. (2) Adsorption was assumed to be the result of two reactions: the formation of monodentate and bidentate surface complexes. The equilibrium constants for Eqn. (2) and for the two adsorption reactions were chosen to give the best fit to the Mg*+ removal data. The constant for Eqn. (2) so calculated corresponds to a much more soluble Mg-hydroxysilicate than sepiolite. The resulting surface association constants, which are listed in Table I, indicate that the lowest pH value at which Mg*+ adsorption occurs is pH 8.0. SIGG and SCHINDLER(unpub. results, reported in SCHINDLER, 1975, 1981) had previously measured Mg*+ removal from am-Si02 suspensions in 1.0 M NaCIO, at 25°C but because the effect of Mg-hydroxysilicate precipitation was not taken into account, the adsorption constants that they determined are too large (SCHINDLER,pers. commun.). WIRTH and GIESKES (1979), who studied the initial kinetics of am-SiO, dissolution in NaCl solutions at 25°C reported that 0.055 M Mg*+, at 0.70 ionic strength and 25°C inhibits the dissolution rate at pH 9.0 but not at pH 8.0. From their data (tabulated in WIRTH, 1980) we compute that their suspensions were undersaturated with respect to sepiolite. The authors concluded that the observed inhibition at pH 9.0 is due to Mg*+ adsorption, and consequently that adsorption is insignificant at pH 8.0. Criteria for distinguishing between MgZf and Mg-hydroxysilicate precipitation

adsorption

It is necessary to develop criteria for distinguishing between Mg*+ adsorption and Mg-hydroxysilicate precipitation in suspensions where either reaction might be occurring. The following important differences between the two reactions can be identified: 1) Mg2+ removal by adsorption may occur much more rapidly than that by Mg-hydroxysilicate precipitation; 2) Mg-hydroxysilicate precipitation removes dissolved silica but the adsorption reaction does not; 3) the two reactions respond differently to decreases in pH. Perhaps a range of solution conditions exists where Mg2+ removal by Mg-hydroxysilicate precipitation is much slower than that by adsorption. Mg2+ adsorption onto am-SiOz reaches equilibrium rapidly (see below). At 25°C however,

many years are required to reach equilibrium with respect to sepiolite precipitation. The precipitation of Mg-hydroxysilicate must therefore be considered to be a dynamic process that will occur in any suspension that is supersaturated with respect to sepiolite at a rate that needs to be determined experimentally, and depends upon the concentrations of the reactants, i.e. Mg*+, dissolved silica, and OH-, the presence of seed crystals, and the temperature. The maximum rate of Mg-hydroxysilicate precipitation under conditions similar to those encountered during the Mg2+ removal from amSi02 suspensions was therefore measured. For suspensions with pH values near 8.0. monitoring dissolved silica concentrations will aid in distinguishing whether Mg*+ removal is due to Mg-hydroxysilicate precipitation or adsorption. According to the ideal formula for sepiolite precipitation (Eqn. I), 0.67 moles of dissolved silica are removed per mole of Mg2+removed during Mg-hydroxysilicate precipitation. Mg2+ adsorption is not accompanied by the removal of dissolved silica. The inhibition of the dissolution rate of am-Si02 at pH 9.0 (WIRTH and GIESKES, 1979) will give rise to an apparent removal of dissolved silica compared to a Mg’+-free system, but no inhibition occurs at a pH of 8.0 (WIRTH and GIESKES.1979). Dissolved silica concentrations were measured in all experiments. The response of the Mg*+ containing am-SiO, suspensions to small decreases in pH should also prove effective in distinguishing between Mg2+ adsorption and Mg-hydroxysilicate precipitation. If adsorption is the major Mg*+ removal process, Mg2+ will desorb in response to a drop in pH. Conversely, unless a drop in pH forces the system across the sepiolite solubihty boundary, Mg-hydroxysihcate precipitation will continue at the lower pH: the precipitation rate may however be affected. In several experiments, the response of Mg2’-containing am-Si02 suspensions to decreases in pH was determined. MATERIALS

AND METHODS

Materials

The solids used in this study are listed in Table 2, along with their BET surface areas. Estimated values of the point of zero charge (PZC) of Aerosil and Ludox are 2.0 to 4.0 (ABENDROTH,1970) and 2.5 to 3.5 (BOLT, 1957), respectively. All samples were brought to constant weight by heating at IlO”C. Massive sepiolite from Eskihi-Sher, Asia Minor (Wards) was powdered by scraping with a stainless steel spatula, suspended overnight in 0.70 M NaCl, rinsed several times with doubly distilled deionized water (2XDW) and freezedried. The resultant powder fits the description of the “poorly crystalline” sepiolite of CHRIST et al. (1973), which is from the same locality. All solutions were prepared using reagent grade chemicals and 2XDW. Acids and bases were prepared and standardized, using established procedures. Phosphate (6.86 at 25°C and 6.95 at 5°C) and borate (9.18 at 25°C and 9.30 at 5°C) buffers were prepared from Beckman powders. Solutions of (%Mg,Na)Cl were made up volumetrically from more concentrated stock solutions in CO,-free ZXDW. Cl- concentrations were determined by Mohr titrations. Experimental

methods

Two types of experiments were used: pH stat experiments, where the pH was held constant with automatic titration equipment (Radiometer), and bottle experiments, where the pH was allowed to drift to a steady value. The pH stat experiments, which were used for determining Mg*+ removal versus time and Mg*+ concentration, were performed at 25.0 + 0.2”C in 500 cm3 of solution under a Nz atmosphere. The suspensions required less than two hours to reach equilibrium with respect to OH- adsorption,

1126

D. B. Kent and M. Kastner Table 2.

Solids used in adsorption and precipitation their soutces and specific surface areas.

Aeroeil(~~)

am-:~o:(c)

S,e~Ci;C (S@Ce(e) 8 ) 218.9 + 3.0

Aerosil-II

aIU-SiC,(C)

210.5 ? 2.0

Ludox (&So

am-SiC2(c)

Ludox-V

am-sio, (C)

Deslffnation

Yp

61.5 f 1.5(d)

131.2 t 1.0

experiments,

Source(b)

_

Aerosil-200, used as received from Degussa Corp. Aerosil(u) heated st 65’~ overnight in concentrated RNO,. rinsed several times, freeze-dried. Recovered from colloidal Ludox-TM (DuPont) by evaporating excess Water at 11o*c, washing precipitate several times in 0.1 ZD 1.0 N HCl then ZXDW.drying at IlO’C, crushing in agate mtt~~. Recovered from colloidal

Ludox-TM (DuPont) by ultracentrifugation; rinsing with 2XDh’(twice), 1 N NaOH(twice), 1 N HCl (twice), 2XDW(several times); freeze drying. Sepiolite

(a)

Poorly crysta11ine Hg?mSiD4)J(uR)4

N2 adsorption,

335

BET surface

(b)

For more details,

(c)

X-ray amorphous.

(d)

Kasrner and Gieskes

z 5

see text. ----

area.

see Kent (1983). (1983).

0.005 M NaCI, which are the conditions used m this work after which time enough M&l2 stock solution was added to that should yield the maximum pH,,. The Ton- was within bring the suspension to the desired Mg*’ concentration. Samples were removed, filtered through 0.45 pm Millipore the tolerance of experimental errors (0.447 rt 0.020 peq m ‘) filters into linear polyethylene (LPE) bottles, and analyzed in suspensions ranging from 275 to 5440 m2 dm ‘, Consefor Mg2+ and dissolved silica. In the Mg” removal YC~.SUJ quently, for this electrode system, the suspension effect is Mg’+ concentration experiments, a 20-minute equilibration smaller than the other sources of error. period was allowed before sampling and adding more Mg2+ The Mg-hydroxysilicate precipitation rate experiments were performed in acid-soaked LPE bottles. Solutions with Rate experiments demonstrated that this was sufficient time to achieve steady Mg*+ concentrations. the desired initial dissolved silica concentrations and pH Alkalimetric titrations and some Mg2+ removal experivalues were allowed to sit overnight to insure that all of the ments were performed in acid-washed 60-cm3 LPE bottles. dissolved silica was present as monomeric species (STUMM The pH of the stirred suspensions was measured with a o ul., 1967). The appropriate amount of MgClz stock solution, sepiolite seed (160 mZ dm-‘), and standardized combination electrode (Radiometer) and adjusted by adding standardized base with a Gilmont microburet. The sealed NaOH solution were added to achieve the desired initial bottles were agitated in a water bath at either 25.0 i- 05°C solution composition. Sampling was accomplished as described above, except that ca. 10 ~1 of I .O M HCI was added or 5 f 1°C. Sampling was accomplished by inserting the to the filtered solutions to prevent any further precipitation. combination electrode, monitoring the pH of the stirred suspension until it reached a steady value (usually within a After sampling, the solution was agitated at 25.0 ! 0.5”C. few minutes) and withdrawing and filtering 2 to 3 cm3 of suspension. In some experiments, suspensions were sampled two or three times over a 2-day period. The surface excess OH-, To”-, was computed from the The combination electrode was calibrated against standard quantity of OH- added to the suspension by correcting for buffers (BATES, 1973) at 25.0 + 0.5”C or 5 + 1°C. Measured OH consumed by the bulk solution. which was determined experimentally; and that consumed by H&O, from the pH values were corrected for the presence of 0.70 M Na’. dissolution of am-SiO, , which was computed from the using the nomograph supplied by the manufacturer, and for dissolved silica and pH data using the dissociation constants drifts in electrode calibration. for H,SiO., in Table 1 (YATES and HEALY, 1976). Dissolved Mg2+ concentrations in the vicinity of 0.056 M wcrc silica concentrations rarely exceeded 50% of saturation with determined by titrating against EDTA to the Enochrome Black T endpoint, and those in the sub-millimolar range respect to am-Si02. The dissolved silica correction ranged from insignificant near pH 7.0 to greater than 40% of I’,,,, were determined by flame atomic absorption spectroscopy above pH 9. I. (AAS). The precision and accuracy are 2% and 5%. respecSome adsorption experiments were performed at low tively. electrolyte concentrations where the pH of dense am-SiOz Dissolved silica (Sir) concentrations were determined. suspensions may exceed the pH of the equilibrium dialysate using the reduced silica-molybdate blue method (STRICKby up to 0.6 pH units (BOLT, 1957). The suspension effect LAND and PARSONS, 1965, as modified by GIESKES, pers. commun.). Optical densities were read on a Gilford 300-N (PH,.,, = pH (suspension) - pH (equilibrium dialysate)) semi-automatic spectrophotometer. The precision and acshould decrease with decreasing surface potential and surface to volume ratio, but the relationship between pH,,,, and curacy are 0.5% and 2%~respectively. these parameters is not known. The influence of the suspenAlkalinities were too small to measure; they were calculated sion effect was investigated by determining the Pou of from the dissolved silica data, the pH data, and the dissoseveral suspensions of Aerosil(u) (Table 2) at pH 8.40 in ciation constants for H4Si04 listed in Table I.

1127

Mg adsorption and precipitation 1001

,

I

I

I

pH above pH 9.0 (Fig. I). All of the suspensions depicted in Fig. 1 were supersaturated with respect to sepiolite. The Mg*+ removal observed for these suspensions could be due to adsorption, Mg-hydroxysilicate precipitation, or both. A rate of Mg*+ removal experiment was performed at a constant pH of 10.6. Mg*+ was added to a suspension of Aerosil-II, present at 162 m* drn-‘, to give an initial concentration of 42.0 PM. After 20 minutes the concentration reached 29.6 PM, where it remained for 2 hours. The pH was then decreased to 9.96 and, within 20 minutes, the Mg*+ concentration increased to 40.0 PM. The suspension was highly supersaturated with respect to sepiolite at both pH values, yet Mg*+ was released by the am-SiO,. Thus. at least some of the Mg2+ removal that occurs in this pH range is due to adsorption. The results of a Mg*+ removal (designated {Mg}app, for apparent amount adsorbed, in pmol m-*) versus Mg*+ concentration experiment at a constant pH of 10.6 are shown in Fig. 2. Typical metal ion adsorbed ver.sit.7 metal ion concentration isotherms are either linear or exhibit negative curvature (BENJAMIN and LECKIE, 1981a,b). The isotherm in Fig. 2a suggests that precipitation of a solid controls the Mg*+ concentration, which is approximately independent of the total amount of Mg*+ added to the system. The suspension was undersaturated with respect to brucite. The concentration of dissolved silica reached an approximately constant value, which was well below the solubility of am-Si02 at pH 10.6 (Fig. 2b). The suspension was highly supersaturated with respect to sepiolite. Mg-hydroxysilicate precipitation thus appears to be the dominant Mg’+ removal process in this experiment. Examination, by Scanning Electron Microscopy, of the solids recovered after the experi-

-

.

Aerosil-II Mpr=43.3

ptd

070M

NaCl

0.

25.O’C 323

rn2 dm-=

. T) OJ

>

. 60

z lF

t

m = 8

40

I 20

I

Oe.0

‘0

.

1

.. l* .

I

6.5

9.0

loo

9.5

105

Ti5

pl-’

FIG. 1. Percent Mg*+ removed after 3.9 + 0.8 hours in Aerosil-II suspensions (solid circles) and after 4.5 f 0.2 hours during Mg-hydroxysilicate precipitation rate experiments (vertical bars). All suspensions supersaturated with respect to sepiolite.

For additional details of experimental procedures see KENT ( 1983).

and analytical

RESULTS

Mg2+ removal from am-Si02 suspension in 0.70 A4 NaCl at 25°C Mg*’ removal from am-SiO? suspensions in 0.70 M NaCl at 25°C is a sharply increasing function of

I

I

r

’

’

1

25.0

a

E.5(

Aerosol -II 0.70 M NoCl 25.0 ‘C OH = 10.6

. 2.oc

20.0

N E $

.

1.x

E i

-6

a o

15 0

1oc

5

h4g2’ MgZ’

r

Absent

.g

Present

. *. I ..

10.0

r _-

.

. .

0.5c

.

c

0

5.0

1.0

IO,,,,,

MgT (PM)

2.0 TIME

3.0

(hrs.)

FIG. 2. a) Mg’+ removed versus Mg*+ concentration isotherm at pH 10.6. Brucite saturated at Mgr = 165 PM. b) Dissolved silica concentrations from same experiment. Suspension allowed to react for 60 minutes prior to first addition of Mg2+. Subsequently, an aliquot of Mg*+ was added and the suspension allowed ca. 20 minutes to react before sampling and adding next aliquot.

D. B. Kent and M. Kastner

1128

ment failed to reveal the precipitate, which accounted for less than 1% of the weight of the solids. isotherms neai {MgL,, versus Mg’+ concentration pH 10.0 are linear for Mg’+ concentrations in the range 0 to 300 PM, but show pronounced positiw curvature at higher Mg2+ concentrations (see KENT. 1983). The positive curvature must be due to the importance of Mg” removal by Mg-hydroxysilicate precipitation in addition to adsorption. In the pH range where Mg2+ removal from amSi02 suspensions can be measured at 0.70 ionic strength and 2S°C, and the solubility of sepiohte is exceeded, Mg2+ could be removed by Mg-hydroxysilicate precipitation as well as by adsorption. Adsorption can be measured in the absence of precipitation only if one can locate a range of solution conditions where the rate of Mg-hydroxysilicate precipitation is much slower than that of adsorption. This possibility was explored in a series of precipitation rate experiments.

Rate of Mg-hydroxysilicate in 0.70 M NaCl at 25°C’

16.51

AAlk

(pmoles)

precipi:atiw

These experiments were designed to mimic the conditions of the adsorption experiments. The results are not rigorous kinetic data for the seeded growth of sepiolite, but should give the maximum rate of precipitation under conditions similar to those used in the adsorption experiments. The initial solution conditions were: Mg2+ concentrations of 88 PM, 260 FM, 435 PM, and 635 PM; dissolved silica concentrations of 640 PM, 1265 PM, and 1930 PM; and pH values of 9.0, 9.5, and 10.0. The initial log IAP (ion activity product) values ranged from 18.9 to 20.6. The experiments were sampled 2 to 4 times during the first 5 hours and less frequently for 2 to 3 weeks thereafter. Control experiments established that: no precipitation occurred in unseeded solutions, no Mg2’ removal occurred when dissolved silica was not present, and the amount of Mg’+ contributed to the solution from the dissolution of the sepiolite seed was less than the analytical errors. except in the experiments where the initial concentration of Mg’+ was near 88 @M. The data presented in Fig. 3 demonstrate that the chemical composition of the Mg-hydroxysihcate that precipitated in these experiments closely resembles that of sepiolite. The ratio of Mg” to alkalinity removed during the experiments equals the 12 molar ratio of sepiolite (Fig. 3a). The ratio of Mg’+ to dissolved silica removed is 0.75. which differs from the ideal ratio of 0.67 for sepiolite (Fig. 3b), but lies within previously reported Mg:Si ratios for synthetic and natural sepiolites (WOLLAST et al., 1968; CouTURE, 1977). The results presented in Fig. 4 show that the rate of Mg-hydroxysilicate precipitation increases with increasing Mg2+ concentration, dissolved silica concentration, and pH. All of these trends are corroborated by the results at other initial solution conditions

FIG. 3. Composition of precipitates from Mg-hydroxysilicate precipitation rate experiments in 0.70 M NaCl at 2S”C. Dashed lines are ratios from ideal formula for sepiolite: Mg2(H4Si04)3(OH)4. a) Mg’+ v~~su.salkalinity removed; b) Mg2+ vcrs~s dissolved silica removed. Solid line: linear regression. with slope = 0.67.

and by analogous plots for dissolved silica and alkalinity removed against time. The observed precipitation rate increases with increasing degree ofsupersaturation. The rate of Mg-hydroxysilicate precipitation 15 rapid enough to account for the removal of a measurable amount of Mg2+ within 4.5 hours throughout the entire range of solution conditions studied. The percentages of the total available Mg” that had precipitated as Mg-hydroxysilicate after GU.4 5 hours. the approximate time scale of the adsorption experiments, are compared with the percentages of Mg” removed during the “adsorption” edge determination in Fig. 1. Mg-hydroxysilicate precipitation could account for part, if not all. of the observed removal of Mg2+. The compositions of the solutions changed throughout the 2- to 3-week durations of the experiments (Fig. 4). Equilibrium with respect to well crystallized sepiolite was not achieved in any of the experiments. In many of the experiments. Mg-hydroxysilicate precipitation continued after the IAP had dropped below the equilibrium value for freshly precipitated sepiolite determined by WOL LAS 1 ct tri

Mg adsorption and precipitation

IO

There is a large effect of increased NaCl concentrations on Mg’+ adsorption onto am-SiOz (Fig. 6). All of the suspensions represented in Fig. 6 were undersaturated with respect to sepiolite. In 0.05 M NaCl and more concentrated electrolytes. Mg2’ adsorption cannot be measured directly without entering the stability field of sepiolite. The adsorption edge for Mg2+ onto quartz in 0.001 M NaCl determined by VUCETA ( 1976) is consistent with the trend shown in Fig. 6. The ionic strength effect is due to competition between Na+ and Mg2+ for surface sites. This can be shown by comparing the extent of Mg” adsorption in 0.01 M solutions of LiC1, NaCI, and KCI. At a given pH and electrolyte concentration, the extent to which alkali metal ions bind to am-SiO, surface in the series Li’ groups, hence the rOH-, incraws < Na+ < K’ (ABENDROTH, 1970; corroborated by the data in Fig. 7a). Mg2+ adsorption in the presence of these electrolytes dmreasrs in the series Lit > Na’ > K+ (Fig. 7b), demonstrating that Mg*+ and the alkali metal ions compete for the same types of surface sites. The relatively large experimental errors accompanying roHm and Mg2+ adsorption measurcments at low adsorption densities can account for the deviations evident at low pH values in Fig. 7.

10

Tlmc

(hrs)

bYQ,.435/AM Si,= 16.5

1265pU

O.?OM

NoCl

25.O”C

Time

(hrs)

220,

= pH=lOOO &l,=435@4 0.70 16 5-

M NoCl

Si,=

1930

+‘M

25.0-C

Time

1129

(hrs)

FIG. 4. Mg*+ removed as Mg-hydroxysilicate versus time for some of Mg-hydroxysilicate precipitation experiments. all in 0.70 M NaCl, 25”C, cu. 160 mz dm-3 sepiolite. Initial concentrations of reactants listed on figures. Effect of: a) Mg*+ concentration, b) pH, and c) dissolved silica concentration.

(1968) and COUTURE (1977) (Fig. 5). In one of the experiments, with initial pH, dissolved silica, and Mg2+ concentrations of 9.0, 1265 PM and 88 PM, respectively, all of the IAP values fall below this “equilibrium” value (slashless. filled circles in Fig. 5).

Adsorption of Mg2+ onto am-Si02 in 0.70 M NaCl at pH 8.0 cannot be detected at low Mg2+ concentrations (Fig. l), but mass action could lead to some adsorption near pH 8.0 at the higher Mg2+ concentrations characteristic of sea water. This was investigated by measuring the displacement of the titration curve of am-Si02 in the presence of 0.056 M Mg2’. at an ionic strength of 0.70, relative to that in 0.70 M NaCI. The results show that adsorption and/or Mg-hydroxysilicate precipitation occurs at pH values above 7.6 (Fig. 8). The pH of these alkaline suspensions of am-Si02 drifted toward neutral. The suspensions represented in Fig. 8 were sampled after 7. 21. and 45 hours in order to examine their responses to small decreases in pH. The major reaction that takes place on this time scale is a drift of the pH toward lower values at constant I’oH-. This widely reported phenomenon for many anhydrous oxides, including am-SiOz (ABENDROTH, 1970; FLEMING, 198 I), has been attributed to hydration of the surface layers of the particles and titration of the newly-formed hydroxyl groups (ONADA and DE RRUYN, 1966: B~R~JB~ et al., 1967). This pH drift creates an uncertainty as to precisely when the original surface groups are in equilibrium. The data in Fig. 8 were obtained by averaging the pH values for the 7-hour and 2 I-hour samples; the drifts during this time are represented by the hori-

1130

D. B. Kent and M. Kastner

D WOllOSt

0

.

160

.

005

0 01

010

050

10

50 Time

10.0

500

100

500

1000

(hrs)

FIG. 5. Log of ion activity product (IAP) for sepiolite versns time for selected results from hydroxysihcate precipitation experiments, along with equilibrium values of WOLLASI PI al. (1968) CHRIST et al. (1973). Symbols: filled circles, squares, triangles, open circles, for initial Mg, = 88 PM, hM, 435 pM, 645 FM, respectively; no slash, right slash, left slash for initial pH and SiT pairs of 9.0 1265 PM, 9.5 and 1265 PM, 10.0 and 640 PM. respectively.

zontal bars through the points and are the same whether or not the suspensions contain Mg2+. The results so obtained for 0.70 M NaCl are in close agreement with the titration curve interpolated from the extensive data set of BOLT (I 957) on the lYon_ of Ludox in NaCl solutions at 25°C. The Ton- versus pH data of CULBERSON et al. (1975) for am-Si02 (Davison silica gel) in sea water at 25°C converge with our data at pH 8.0 and below

50.0 I

43.3pM

400

-

Mg,

0.050 M NaCl

0

i

77 aJ g

Experlmentol

0

Error

0 0

300-

a

u)

z

0

A

D

$ mo10.0

a 0

0

-

A

A

0 0 0

o-

A

L__I

A

Mg*’ removal ut 0.70 ionic .strength und t”(

DISCUSSION

A 00

(Fig. 8). CULBERSON et al. (1975) point out that sepiolite precipitation probably affects their data at pH values above 8.5. In the pH range 7.5 to X.5. our data for Mg”-containing suspensions are close to their data for sea water. Thus, at 25°C the other sea water ions have little or no effect on the l’,,, ol am-SiO, .

The effect of 0.056 M Mg” on the I’o,i w:,u.\ pH behavior of am-Si02, at 0.70 ionic strength and 5°C was investigated. No Mg” removal occurs at pH 7.9; a small amount of removal occurs at higher pH (Fig. 8). All Mg2+-containing suspensions were supersaturated with respect to sepiolite. The Ton- verSu.y pH trend for am-SiOz in sea water at 2°C reported by CULBERSON c’t aI. i lY75) lies parallel to and slightly below our data at 5°C. This suggests that the other major ions in sea water do not affect the To”- of am-Si02 near pH 8.0. The decrease in temperature from 25°C‘ to j”C causes a small but significant drop in the I‘,,- of am-Si02 in 0.70 M NaCl (Fig. 8). This corroborates the findings of CULBERSON et ul. (1975) that the rou- of am-SiOz in sea water decreases with temperature from 25°C to 2°C. A drop in ran from 8 to 16% relative to the value at 25°C was calculated from these data (Table 3).

Ludox-P

25.0°C 390 m2dm-= 0 0.005 M NoCl A 0.010 M NoCl

Mgand 260 and

00

00

0

00

7.0

i 90

PH FIG. 6. Effect of NaCl concentration on Mg’+ adsorption onto am-SiO*. At lowest adsorption densities, experimental error somewhat greater than shown,

.Msorption versus Mg-hydrox.v.~iticutr precip~tutron at 0.70 ionic strength, 25°C’ Mg2+ removal from am-SiO, suspensions at 0.70 ionic strength, 25°C and pH values above 9.0 (Fig. I), occurs by both adsorption and Mg-hydroxysilicate

1131

Mg adsorption and precipitation

040

0 OOIOMLiCt

-

0 0010 -

0.010

400-

M KCI M FIPCI

00010M

Lict

v O.OiOM

KC1

-0OlOMNoCI

N ‘&

O30-

::

z i.

=I

.

m

.8zoo-

g 020”

=

L

0 ia -

FIG. 7. Effect of cation type on a) roH- and b) Mg2* adsorption onto am-SiOz. At low adsorption densities, experimental error somewhat greater than shown.

precipitation. Two pieces of experimental evidence indicate that some of the Mg2” removed is due to adsorption. Mg2+ is desorbed in response to a decrease in pH in systems that are supersaturated with respect to sepiolite. {Mg}app versus Mg2+ concentration isotherms for low Mg2+ concentrations near pH 10.0 resemble adsorption isotherms, i.e. f MgjaPP is a linear The latter is a necfunction of Mgzt concentration. essary but not sufficient criterion to demonstrate that the removal is due to adso~t~~~ because Mg-hydrctxysilicate precipitation could conceivably give rise

A l

J.oO-

I

1

I 0.056 M Mgr

+ 0.535

M NaCI

25 QC 5 *c

0,7Cl M No&L 25 *c (Bait)

“*

I / : ‘-At

h 25 OC 0 5 OC CT 2.00

-

‘E

lAIdor-Y S= 315 m2 dm-'

)!AI /'A /f .$"

,AA

1.00-

-

a

,I

z ,r 'I !c

Se0 Water ICulbersonl 25 *c

-Af . 24 I *o"

Cl

/

/pi/-@

0

I

I

I

7.0

8.0

9.0 pH

FIG. 8. Apparent roH- W~SIISpH for suspensions at 0.70 ionic strength with or without 0.056 M A,lg2+at 25 k 0.5”C and 5 + I “C. All M$‘-containing suspensions supersaturated with respect to sepiolite. Solid curve interpolated from data of BOLT (1957) for ToH- of Ludox in N&l solutions at 25T. Curve labeled “Culberson”: data of CULBERSON et al.(L975) for am-W& in sea water. Saturation with respect to brucite: pH G 9,3 and 10.1 at 25°C and 5’C, respectively.

to quasi-linear (Mg],,, W’SUS Mg2+ concentration isotherms. Two pieces of experimental evidence indicate that Mg-hydroxysilicate precipitation is also responsible for some of the Mg2+ removal. The rate of Mg-hydroxysilicate precipitation is rapid enough to account for the observed Mg2+ removal (Fig. 1). The shape of the (Mg),,, YU’SUSMg2+ concentration isotherm at pH 10.6 indicates that precipitation is the major Mg” removal process (Fig. 2a); the dissolved silica concentrations show that the precipitate is a Mg-hydroxysilicate (Fig. 2b). Although no Mg*’ removal could be measured at pH values lower than 9.0 in suspensions with low Mg2” concentrations (Fig. 1), in suspensions with 0.054 M Mg’+ a reaction between Mg*+ and am-Si02 can be detected down to pH values near 7.5 (Fig. a)_ At pH values below 9.0, dissolved silica concentrations can be used in conjunction with the rOHm data to show that this reaction is Mu-hydroxysi~~cate precipitation. Dissolved silica concentrations in Mg’+-containing suspensions are lower than those in Mg2+free suspensions at the same pH, ionic strength, surface to volume ratio and reaction period (Fig. 9a). The dissolved silica concentration expected in the Mg”+-containing suspensions, if Mg-hydroxysilicate precipitation is the only reaction occurring, can be computed from: the dissolved silica concentrations in the corresponding Mg2”-free suspensions, the excess OH- consumed in the Mgz+-containing suspensions over those free of Mg”, and the ratio of Si to OHconsumed during Mg-hydroxysilicate precipitation. The calculation is described in Appendix I. Calculated concentrations are in excellent agreement with the measured concentrations (Fig. 9a). Thus. near pH 8.0, Mg-hydroxysiiicate precipitation appears to be the significant reaction in the system. CISLER ( 198 1) had previously calculated equilibrium constants for the adsorption of Mg2+ onto amSiOz (see Table 1) based on an equilibrium approach of distinguishing Mg2+ adsorption from Mg-hydroxy-

1132

D. B. Kent and M. Kastner Table

3.

Effect of tempperature on rOHand Ha adsorption onto Ludox-V

(a) (ueq

TO-d5 + 1’C

PH + t t

.02 .02 .02

0.79 0.85

t * --

.02 .02

0.90 0.92

8.46 8.71 a.03 9.02

f ?r ? t

.02 .02 .02 .02

1.14

? -+ *

.02

1.24

?

.04 .06

1.53 1.71

‘1 .04 + .04

9.09

2

.02

(a)

Values at 25’C interpolated with maxim.m experimental

(b)

AFOH- = rOH-(25”C) Calculated at

?. .03 ? .03 -_

--

(c)

from

given

pH,

rOH_(25’Cr

+ 0.5OC

7.91 7.94 8.33

1.38 1.43

.04

100%

0.70

M N&l

X Mg Adsorbed

0.09 0.09 0.14

-experimental

(5”C)(c)

0% -_

8% -10% 16%

from errors.

+ .03x ? .04% ? .05x -_. __

0.25 results.

All

?

values

,074, presented

rOH-(S”C)

difference

assuming

x

13% 8% -_

--

-

in

hroH-(b)

m-‘)

25.0

of Ludox-V at S’C.

1.5

between + 0.5

TOH- values

moles

precipitation. His constants are tested against our data for Mg2+ removal in am-Si02 suspensions at 0.70 ionic strength and 25°C in Fig. 10. The concentrations of free MgZi, i.~. [Mg”], were computed from the concentrations of total dissolved Mg’ ’ using the constants in Table I. The density of unionized surface silanols as a function of pH, which is necessary to compute the extent of adsorption from Gisler’s constants, was computed as described in Appendix 2. The scatter in our data is probably due to the fact that both adsorption and Mg-hydroxysilicate precipitation remove Mg’+, and the extent ot Mg-hydroxysilicate precipitation varies due to, among other variables. different reaction times silicate

H+ released

with

and per

without

mole

Mg’+

0.056

M MgT

adsorbed.

The calculated extent of’ adsorption exceeds our measured amount of Mg’+ removal. In the vicinit> of pH 8.0. there appears to be reasonable agreement between the data and the extent of adsorption calculated using the lowest estimate for the density 01 silanol groups on Ludox. Mg’+ removal in this pH range, however, is due to Mg-hydroxysilicate precipitation and not adsorption. In computing his constants for adsorption, GISLER ( 198 I ) assumed that his suspensions were in equilibrium with respect to Mg” adsorption, am-Si02 solubility, and the precipitation of a Mg-hydroxysilicate much more soluble than sepiolite. Our results show that Mg-hydroxysilicate

---

--.__l

70

75

80

85

90

PH

FIG. IO. Mg*+ removed ({ Mg},,) divided by concentration of free Mg ([Mg”]) vcr.s~(s pH. from experiments with amSiOz at 0.70 ionic strength. 25°C. All suspensions supersaturated with respect to sepiolite. Line labeled ‘Yiisler” is trend computed using GISLER’S ( 1981) constants for Mg” 75

80 PH

85

90

adsorption onto am-SiOz, in I.0 M NaC104 at 25°C (see Appendix 2). Insert: {Mg},,,/[Mg*‘] for Ludox-V. pH 7.5

to 9.0, computed from titration data assuming I .5 i 0.5 FIG. 9. Dissolved silica concentrations from IToH- VC%SN.Y H+ released per Mg’+ adsorbed (limits denoted by vertical pH determinations at 0.70 ionic strength and at a) 25°C bars). Lines calculated from Gisler’s constants for 3 estiand b) 5°C. Ellipses: calculated concentrations (see text and mates of total density of surface hydroxyls on L udox (ApAppendix I ). pendix 2).

1133

Mg adsorption and precipitation precipitation begins after short reaction periods, but does not reach equilibrium even after reacting for 3 weeks. Gisler’s thermodynamic approach to distinguishing Mg*+ adsorption from Mg-hydroxysilicate precipitation was therefore unsuccessful.

Adsorption versus Mg-hydroxysilicate at 5°C

precipitation

At 5°C and a Mg*+ concentration of 0.056 M, reaction between Mg*+ and am-Si02 can be detected beginning at pH values slightly higher than 8.0 (Fig. 8). Below pH 9.0, dissolved silica concentrations in Mg*+-containing suspensions are not below those in corresponding Mg*+-free suspensions (Fig. 9b). It is not clear why there is more scatter in the dissolved silica data from Mg*+-containing suspensions than those from the Mg’+-free suspensions. Nevertheless, these results show that Mg-hydroxysilicate precipitation is not significant between pH 8.0 and 9.0; hence Mg2+ adsorption must be occurring. The extent of Mg*+ adsorption, which is presented in Table 3, was calculated from the IoH- data based on the assumption that 1.5 + 0.5 H+ ions are released per Mg2+ ion adsorbed (SCHINDLER, 1981; BENJAMIN and LECKIE, 198 1a).

Strength of‘Mg2+-am-Si02 surface complexes Mg*’ forms relatively weak complexes with amSi02 surface groups and binds to similar types of surface sites as do the alkali metal ions. This is demonstrated by the shift in the Mg*+ adsorption edge toward higher pH values with increasing NaCl concentration (Fig. 6) and the inverse correlation between the extent of Mg*+ adsorption and the extent of adsorption of Li+, Na+, and K+ (Fig. 7). A comparison of the positions of Mg*+ adsorption edges onto am-Si02 (this work), a-FeOOH (BALISTRIERI and MURRAY, 1981), and y-A120X (HUANG and STUMM, 1973) shows that the oxides of Fe and Al have a higher affinity for Mg*+ than does am-Si02. BENJAMIN and LECKIE (1980) had previously shown that quartz has a lower affinity for adsorption of Cd*+, Zn*+, CL?+, and Pb*+ than Y-A1203 and yFeOOH.

Sepiolite solubility at 25°C The results from our Mg-hydroxysilicate precipitation rate studies, as presented in Fig. 5, show that the solubility of sepiolite reported by WOLLAST et al. (1968) and COUTURE (1977) is too high; our data therefore support the lower solubility at 25°C that was calculated by CHRIST et al. (1973). Wollast et al. and Couture precipitated sepiolite from homogeneous solutions, a method that requires high degrees of supersaturation. The resultant fine-grained and poorly crystalline precipitates age slowly, hence require many years to reach equilibrium (COUTURE, 1977). It thus

appears that these authors to reach equilibrium. GEOCHEMICAL

allotted

insufficient

time

IMPLICATIONS

Barring an unexpectedly large pressure effect, adsorption of Mg*+ onto am-Si02 is an insignificant process in sea water and in marine sediments. Furthermore, the similarity between the I’ou- versus pH behavior of am-Si02 in suspensions containing Na+ and Mg*+ (this work) to that in sea water (CULBERSON et al., 1975) suggests that Ca”, K’, and SO:- neither adsorb onto am-SiO, to a significant extent nor enhance the adsorption of Mg2+ onto am-SiOz. The electrical surface properties of am-SiOz in sea water should therefore be determined primarily by the extent of interaction between Na+ and silanol groups. This argument, of course, refers to pristine surfaces. The role of organic and inorganic surface coatings has not been investigated. Mg*+ adsorption may affect significantly the surface charge properties of am-Si02 and quartz in fresh waters. A computation by VUCETA and MORGAN (1978) however. indicates that adsorbed Ca2+ dominates silica surfaces in a fresh-water medium. This is due to the higher abundance of Ca2+ than Mg2+ in fresh-water systems and the ease with which Mg2+ is displaced by competing adsorbates. In marine sediments, the important chemical reaction in this system is the precipitation of Mghydroxysilicate, i.e. sepiolite. The solubility of sepiolite increases with decreasing temperature and increasing pressure (CHRIST et al., 1973; COUTURE, 1977: SAYLES, 198 1). Consequently, even highly siliceous surface sediments, with dissolved silica concentrations and pH values as high as 750 PM and 7.8, respectively, apparently are undersaturated with respect to sepiolite and Mg-aluminosilicates (SAYLES,1981). The precipitation of sepiolite is thus only likely to occur at depth in the sediment column, where temperatures are greater than cu. 10°C. Sepiolite precipitation is likely to occur in carbonate-containing siliceous sediments because the CaCO, dissolution helps to maintain pH values near 8.0. The temperature effect on the solubility might explain why sepiolite is rare in Recent marine sediments (KASTNER, 198 1). Sepiolite occurs in older marine sediments, often together with palygorskite, but is apparently detrital in many occurrences (KASTNER, 1981). Only in a few cases, however, there is unequivocal evidence that it is authigenic (COUTURE, 1977; KASTNER, 198 1). These occurrences reflect the importance of rates of competing chemical processes such as nucleation and growth rates of sepiolite, palygorskite, opal-CT and/ or quartz during diagenesis of siliceous sediments. DONNELLY and MERRILL (1977) suggested that Mg2+ adsorption onto am-Si02 accounts for the good correlation between Mg/Al and Si/Al ratios that they observed in some carbonate-bearing siliceous sediments. Both ratios increase with depth in the sediment.

1134

D. B. Kent and M. Kastner

From the pore water data for two adjacent sites studied by DONNELLY and MERRILL (1977) CouTLJRE (1977) calculated that the pore waters arc supersaturated with respect to sepiolite at the site where Mg/Al and Si/AI positively correlate, and undersaturated at the site where they do not (Venezuelan Basin Sites 29 and 149, respectively). All of the evidence available at present thus supports the suggestion of KASTNER and GIESKES (1983) that these MgJAl and Si/Al correlations are due to the precipitation of a disseminated and disordered sepiolite-like Mg-hydroxysilicate, and not due to Mg” adsorption onto am-SiOz The results of this study also support the suggestion of KASTNER and GIESKES ( 1983) that the Mg-OH compound responsible for enhancing the rate of opalA to opal-CT conversion is a Mg-hydroxysilicatc resembling sepiolite. Of brucite. magnesite, and sepiolite, only sepiolite is supersaturated in all of the experiments of KASTNER and GIESKES (1983) as shown by KENT (1983). Although Mg-hydroxysilicate is not the only compound capable of enhancing the rate of this transformation (KASTNER r/ (I/.. 1977). sepiolite is the most likely compound to precipitate and enhance the transformation rate of opal-A to opal-CT at the diagenetic temperatures at which the reaction takes place in siliceous marine sediments. Acknowledgements-We thank Dr. J. M. Gieskes for invaluable discussions and critical review of this manuscript. Stimulating discussions with the late Dr. Jerry Wirth were influential in the early stages of this work. We thank R. T. LaBorde, G. Anderson, D. Burdige, A. Shiller, and P. Baker for helpful discussions and laboratory assistance; M. Delaney. N. Hinman, G. Redden, and K. Gruebel for their constructive comments, and R. Couture and an anonymous reviewer for their detailed reviews. The typing of M. Hitchcox and D. Chilton and the drafting skills of W. Borst and N. Ferguson sped the production of this paper. D. Kent acknowledges the receipt of a fellowship from Kennecott Copper Corporation. The work was made possible by NSF Grant OCE8024630 (M.K.) and ACS/PRF Grant #15262-AC2 with additional support from Chevron Oil Field Research Company. Editorial handling: S. E. Calvert REFERENCES ABENDROTHR. P. (1970) Behavior of pyrogenic silica in simple electrolytes. J. Cobid Inre+ce Sci. 34, 591-596. ALLEN L. H. and MATIJEVI~.E. (1970) Stability of colloidal silica II. Ion exchange. J. Colloid Interfar,c>Pi. 33, 420429. BADAUTD. and RISACHERF. (lY83) Authigenic smectite on diatom frustules in Bolivian saline lakes. Geochim Cosmochim. Acta 47, 363-375. BAES C. F. and MESMER R. E. (1976) The Ifydro!vsis ol Cations, Wiley-Interscience, 489 pp. BALISTRIERIL. S. and MURRAY J. W. (1981) The surface chemistry of goethite (aFeOOH) in major ion seawater. Amer. J. Sri. 281, 788-806. BALISTRIERI L. S. and MURRAYJ. W. (1982) The adsorption of Cu, Pb, Zn and Cd on goethite from major ion seawater. Geochim. Cosmochim. Acta 46, 1253-1265. BALISTRIERIL. S., BREWER P. G. and MURRAY J. W. (198 1) Scavenging residence times of trace metals and

surface chemistry of sinking particles in the deep ocean. Deep-Sea Rex 28A, 10 I- I2 I. BATESR. (1973) Determination yJ’pIf: Theor!, nnd Prac’licc, Wiley-Interscience, 479 pp. BENJAMINM. M. and LECKIEJ. 0. (1980) ‘The adsorption of metals at oxide water interfaces: effects of the concentrations of adsorbate and competing metals. In (bntaminants and Sediments. (ed. R. A. BAKER). Vol. 2. Chap. 16, pp. 305-322, Ann Arbor Sci. Pub. Inc. BENJAMINM. M. and LECKIEJ. 0. (1981a) Multiple site adsorption of Cd, Cu, Zn, and Pb on amorphous Iron oxyhydroxide. J. Colloid Interface Sci 79, 209-22 1. BENJAMINM. M. and LECKIE J. 0. ( 198lb) Competitive adsorption of Cd, Cu, Zn and Pb on amorphous Iron oxyhydroxide. J. Colloid Interfuce Sci. 83, 410-4 19. B~Ru& Y. G.. ONADAG. Y. and DE BRUIN P. 1.. (1967) Proton adsorption at the ferric oxide/aqueous interface II. Analysis of kinetic data. Surf: Sci. 8, 448-46 I BOLT G. H. (1957) Determination of the charge density 01 silica sols. J. Ph.vs. Chem. 61, 1166-l 169. BOURG A. C. M. and SCHINDLERP. W. (1978) Ternary surface complexes 1. Complex formation in the system silica-Cu(II)-ethylenediamine. Chimia 32, 166-168. BOURN;A. C. M., Joss S. and SCHINDLERP. W. (lY7Y) Ternary surface complexes 2. Complex fomration in the system silica-Cu(II)-2,2’bipyridyl. Chimia 33, 19-2 1. BIJSEYR. H. and MESMERR. E. (1Y77) Ionization equilibria of silicic acid and polysilicate formation in aqueous sodium chloride solutions to 300°C. Itzorg (‘hem 16, 2444-2450. CHRISTC. L., HOSTETLERP. B. and SIEBERI R. M. (19731 Studies in the system MgO-SiO*-CO*-HZ0 (111): The activity-product constant of sepiolite. .4mer I .%I 273, 65-83. COUTURER. A. (1977) Synthesis of some clay minerals at 25°C: palygorskite and sepiolite in the oceans. Ph.D. dissertation, Univ. California, San Diego. CULBERSONC. H., PYCTOWITZR. M.-and A-11,~s t. L. (1975) Hydrogen ion exchange on amorphous silica in sea water. Mar. Chem. 3, 43-54. DAVISJ. A. and LECKIEJ. 0. (1978) Surface iomzation and complexation at the oxide/water interface II. Surface properties of amorphous Iron oxyhydroxide and adsorption of metal ions. J. Colloid Intecfke Sci. 67, 90-107. DAVISJ. A.. JAMESR. 0. and LE~KIE J. 0. (I 078) Surface ionization and complexation at the oxide/water interface 1. Computation of electrical double layer properties in simple electrolytes. J. Colloid Inter&e Sci. 63. 480-499. DONNELLYT. W. and MERRILL.I,. (1977) The scavenging of magnesium and other chemical species by hiogenic opal in deep-sea sediments. Chem. Geol 19, 167-186. DREVER J. 1. (1974) The Mg problem. In 77te Sea, (ed. E. D. GOLDBERG).Vol. 5. Chap. IO. pp. 337-757, WileyInterscience. FLEMINGB. A. (I 98 1) Polymerization kinetics and romzation equilibria in aqueous silica solutions. Ph.D. dissertation Princeton Univ. GlSLtR A. (1981) Die Adsorption van Aminosauren an Grenzflachen oxid-Wasscr. Ph.D. dissertation. Univ. Bern. Switzerland. GREENBERGS. A. and PRICE E. W. (1957) The solubility of silica in solutions of electrolytes. J, f/zy.s. Chem. 61, 1539-1541. HARNED H. S. and OWEN B. B. (1958) 7%~ Pi~~:s~c’cri Chemistry of Electrol_vte Solution.s. Third ed.. Reinhold Publ. Corp., 803 pp. HELGESONH. (1967) Thermodynamics of complex dissociation in aqueous solution at elevated temperatures. J Ph.vs. Chem. 10, 3 121-3 136. HELGESONH. ( 1969) Thermodynamics of hydrothermal systems at elevated temperatures and pressures. .jmer .I Sci. 267, 729-804.

Mg adsorption and precipitation HEMLEYJ. J., MONTOYAJ. W., CHRISTC. L. and HOSTETLER P. B. (1977) Mineral equilibria in the MgO-SiO*H*O system: I. Talc-chrysotile-forsterite-brucite stability relations. Amer. J. Sci. 271, 322-35 I. HESTONW. M., ILER R. K. and SEARSG. W. (1960) The adsorption of hydroxyl ion from aqueous solution on the surface of amorphous silica. J. Phys. Chem. 64, 147-150. HOSTETLERP. (1963) The stability and surface energy of brucite in water at 25°C. Amer. J. Sci. 261. 238-258. HUANG C. P. and STUMMW. (1973) Specific’adsorption of cations on hydrous y-Al*03. J. Colloid Interface Sci. 43, 409-420. HURD D. C. (1973) Interactions of biogenic opal, sediment and seawater in the Central Equatorial Pacific. Geochim. Cosmochim. Acta 31, 2257-2282. JONES J. B. and SEGNIT E. R. (1971) The nature of opal. I. Nomenclature and constituent phases. J. Geol. Sot. .4ustr. 18, 56-68. JORGENSENS. S. (1968) Solubility and dissolution kinetics of precipitated amorphous silica in I M NaC104 at 25°C. Acta Chemica Stand. 22, 335-341. KASTNER M. (1981) Authigenic silicates in deep-sea sediments: formation and diagenesis. In The Sea, (ed. C. EMILIANI),Vol. 7, Chap. 13, pp. 915-980, Wiley. KASTNERM. and GIESKESJ. M. (1983) Opal-A to opal-CT transformation: a kinetic study. In Siliceous Deposits in the Pact@ Region, (eds. A. IIJIMA, J. R. HEIN and R. SIEVER),pp. 2 I l-228, Elsevier. KASTNER M., KEENE J. B. and GIESKES J. M. (1977) Diagenesis of siliceous oozes. I. Chemical controls on the rate of opal-A to opal-CT transformation. Geochim. Cosmochim. Acta 41, 1041-1049. KENT D. B. (I 983) On the surface chemical properties of synthetic and biogenic amorphous silica. Ph.D. dissertation, Univ. California, San Diego. MARSHALLW. and WARAKOWSKIJ. M. (1980) Amorphous silica solubilities II. Effect of aqueous salt solutions at 25°C. Geochim. Cosmochim. Acta 44, 915-924. ONADAG. Y. and DE BRUYNP. L. (1966) Proton adsorption at the ferric oxide/aqueous solution interface 1. A kinetic study of adsorption. .Sucj: Sci. 4, 48-63. SANTSCHIP. H. and SCHINDLERP. W. (1974) Complex formation in the ternary systems Cu(II)-H.,Si04-H*O and M&II)-HlSi04-H*O. J. Chem. Sot. Dalton 2, I8 I - 184. SAYLESF. L. (1981) The composition and diagenesis of interstitial solutions-II. Fluxes and diagenesis at the water-sediment interface in the high latitude North and South Atlantic. Geachim. Cosmochim. Acta 45, 10611086. SCHINDLERP. W. (1975) Removal of trace metals from the oceans: a zero order model. Thalassia JugosIavica 11, 101-111. SCHINDLERP. W. (198 I) Surface complexes at oxide-water interfaces. In Adsorption qf Inorganics at Solid/Liquid Interfaces, (eds. M. A. ANDERSONand A. J. RUBIN), Chap. I, pp. l-49, Ann Arbor Sci. Pub. SCHINDLERP. W., FURST B., DICK R. and WOLF P. U. (1976) Ligand properties of surface silanol groups I. Surface complex formation with Fe’+, Cu*+, Cd*+ and Pb*+. J. Colloid Interface Sci. 55, 469-475. STRICKLANDJ. D. H. and PARSONST. R. (1965) A manual for sea water analysis. Bull. Fish. Bd. CQnQdQ 167, l31 I. STUMM W., HUPER H. and CHAMPLINR. L. (1967) Formation of polysilicates as determined by coagulation effects. Environ. Sci. Technol. 1, 221-227. TYLERA. J., TAYLORJ. A. G., PETHICAB. A. and HOCKEY J. A. (197 I) Heat of immersion studies on characterized silicas. Trans. Farad. Sot. 67, 483-492. UYTTERHOEVEN J., HELLINCKXE. and FRIPIATJ. J. (1963) Le f&age des gels de silice. Silicates Industriels 23, 241246. VUCETA J. (1976) Adsorption of Pb(II) and Cu(I1) an OIquartz from aqueous solutions: influence of pH, ionic

1135

strength and complexing ligands. Ph.D. dissertation, California lnstitute of Technology. VUCETA J. and MORGAN J. J. (1978) Chemical modeling of trace metals in fresh waters: role of complexation and adsorption. Environ. Sci. Technol. 12, I302- 1308. WERES O., YEE A. and TSAO L. (1981) The kinetics of silica polymerization. J. Colloid Interface Sri. 84, 379402. WIRTH G. (1980) The heterogeneous kinetics of silica dissolution in aqueous media. Ph.D. dissertation, Univ. California, San Diego. WIRTH G. and GIESKESJ. M. (1979) The initial kinetics of the dissolution of vitreous silica in aqueous media. J. Coiioid Interface Sci. 68, 492-500. WOLLASTR., MACKENZIEF. T. and BRICKER0. P. (1968) Experimental precipitation and genesis of sepiolite at earth-surface conditions. Amer. Minerul. 53, I645- 1662. YATES D. and HEALY T. W. (1976) The structure of the silica/electrolyte interface. J. Colloid Interface Sci. 55, 920. YOUNG J. R. (1981) A study of the adsorption of Ni(II) onto an amorphous silica surface by chemical and NMR methods. Ph.D. dissertation. California Institute Technology. APPENDIX

1

After reacting for 21 hours, one of the Mg*+-containing suspensions had reached a pH of 8.17 and Ton- of 1.52 pmole mm*. The Ton of a suspension in 0.70 M NaCl at pH 8.17, as read off a smooth curve drawn through the data, as in Fig. 8, is I. 13 pmole m-*. Thus, 0.39 pmole m 2 more OH- was consumed in the Mg*+-containing suspension than in the suspension without Mg2+. All suspensions in these experiments had surface to volume ratios that fell in the range 3 I5 ? 10 m* dmm3. The excess concentration of OH- consumed in the Mg*+-containing suspension is 123 pmoles OH- dmm3.The ratio of Si to OH- consumed during Mg2+-hydroxysilicate precipitation is 0.67 (see Fig. 3); therefore, the removal of I23 pM of OH- should be accompanied by the removal of 82 pM of dissolved silica. The concentration of dissolved silica achieved by a suspension with 3 I5 m* dm-’ of am-SiO, in 0.70 M NaCl under identical agitation conditions after 2 I hours is 1090 PM (as read off a smoothed trend in Fig. 9). The dissolved silica concentration in the Mg*+-containing suspension should be 1090 PM - 82 @M. or 1008 WM.if Mg-hydroxysilicate precipitation is the reaction between Mg*+ and am-SiO, evident in Fig. 8. The measured concentration for the Mg*+-containing suspension at pH 8.17 after 21 hours in 1010 PM. APPENDIX Computation af (SiOH)

2

for Aerosil sampkr

(rSiOH} is computed by subtracting the density of ionized silanols, {=SiO-} in moles kg-’ Si02. from the total density of silanol groups, {=SiOH}r in moles kg-’ SiO*. GISLER (1981) determined {=SiOH}, to be 2.6 f 0.1 moles kg-’ Si02, which corresponds to 4.4 k 0.2 sites nm-*, by titrating suspensions of Aerosil-500 to high pH. There is a large experimental uncertainty associated with titrating am-SiO* suspensions to high pH because a large correction for the OH- consumed by dissolved silica must be made. For example, in suspensions of Aerosil-II, at pH 10.6 in 0.70 M NaCl at 25°C the correction increases with time and, after 20 minutes, exceeds 100% of the measured {=SiO-} (KENT, 1983). Extensive work on evaluating the ( =SiOH}T of Aerosil silica demonstrates that variations in the density of micropores, i.e. surface pores and capillaries of molecular dimensions, give rise to variations in the density of silanol groups, with 4.6 sites nm-* being a minimum value (TYLER et al., 197 1). The {=SiO-} in 0.70 M NaCl at 25°C of our Aerosil samples (Table 3) in the pH range 8.0 to 9.0 are 30% lower

1136

D. B. Kent and M. Kastner

than those reported by GISLER (1981) for his Aerosil-500. in 1.0 M NaClO, at 25°C and lower by smaller amounts than the corresponding values reported elsewhere (SCHINDLER et al.. 1976; BOURG and SCHINDLER, 1978: BOURG et al.. 1979; YOUNG, 1981). Despite the discrepancy with the results of TYLER ef al. (1971) it was assumed that our Aerosil samples have silanol densities 30% lower than GrsLER’s (1981). The trend in Fig. 10 thus represents the minimum amount of adsorption predicted by the constants of GISLER(198 I ). The {=SiO-} of our Aerosil samples could not be measured with satisfactory accuracy in the pH range of interest to this calculation because of the large dissolved silica correction. In order to maintain consistency with GISLER( 198 I), we computed {=SiO-} by extrapolating his titration data into the pH range of interest and adjusting them downward by 30% (KENT. 1983).

Computationof {SiOH) for Ludox .samples Various estimates of {=SiOH}T on Ludox am-SiO* have been reported: 2.3 sites nm-*, from a combination of thermogravimetric data, IR absorbance intensities and determinations of the amount of CH., produced by reaction of Ludox with CH3MgI and CHSMg (UYTTERHOEVENPI al., 1963). 3.5. sites nm-‘, by titrating suspensions to high pH in the presence of various electrolytes (HESTON ef a/., 1960; ALLEN and MATIJEVI~, 1970), and 4.6 sites nm -‘, calculated from an analysis of the extensive titration data of BOLT (1957), (FLEMING, 1981). WERESet al. (198 I) estimated a value of 6.84 sites nm-* for Ludox undergoing deposition of dissolved silica from supersaturated solutions. It is unlikely. however, that this value is characteristic of Ludox undergoing dissolution (WERES et al., 1981). (=SiOH} values were calculated from each of these total silanol densities and the IoHm values measured in 0.70 M NaCl at 25°C (Fig. 8).