TECHNICAL REPORTS
Wetlandsand Aquatic Processes
Mineral Precipitation
and Trace Oxyanion Behavior
Evaporation of Saline Waters
during
D. B. Levy,* C. Amrhein, and M. A. Anderson
ABSTRACT
Evaporationbasins used for the disposal of agricultural drainage
waters in central California contain elevated trace elementlevels that
pose hazards to groundwaterquality and wildlife visiting the ponds.
A study was conductedto evaluate the solution dynamics of mineralforming elements and trace oxyanions (U, Mo, and V) during evaporation of saline waters whose chemical compositions evolve according
to two distinct chemical divides, and to characterize the evaporite
minerals formed from the complete evaporation of these waters. The
alkali and alkaline earth metals exhibited nonconservative behavior,
forming evaporite minerals such as bloedite, calcite, aragonite, gypsum, halite, thenardite, and trona. Molybdenum
behaved nonconservatively, while V exhibited conservative behavior that did not differ
whether V was initially added as V(IV) or V(V). Uranium displayed
conservative behavior under conditions of low U concentrations and
high alkalinities. Nonconservative behavior was observed for U, however, under higher U concentrations andlow alkalinities. Weconclude
that V and U in waters with alkalinities
>10 mmol¢ L-~ will not
accumulatein evaporationpondminerals. In ponds with low alkalinity,
U will partition to a solid mineral phase,
2900ha of land in the Central Valley
~ PROXIMATELY
of Californiaare usedfor the containment
and subsequentevaporativedisposal of agricultural drainagewater.
In 1988and 1989,the California Central ValleyRegional
Water Quality Control Board (RWQCB)
conducted water
quality surveys of 28 evaporation basins used for the
disposal of irrigation drainage water in the Tulare Lake
Basinof the SanJoaquinValley (Chilcott et al., 1990).
These investigations were promptedby findings that
elevated Se levels have caused waterfowl deaths and
deformitiesin areas used for the disposal of the drainage
waters (Ohlendorfet al., 1986). Localgroundwatercontamination due to recharge from drainage waters with
elevated concentrations of As, B, Mo,Se, V, and other
trace elements has also been of concern (Deverel and
Millard, 1988; Tanji and Valoppi, 1989).
Results from the RWQCB
water quality survey demonstrated that evaporationgenerallyincreased the salt and
trace element levels of drainage waters in evaporation
basins compared
with the inlet waters. Thetotal dissolved
solids of inflow waters ranged from 450 to 57 000 mg
L-1, while those of the basins rangedfrom800to 395 000
mgL-1 (Chilcott et al., 1990).Inlet watersto the basins
D.B. Levy, Dep. of Soil and Environ. Sci., Univ. of California, Riverside,
CA92521 (currently Dep. of LAWR,Univ. of California, Davis, CA
95616); and C. Amrheinand M.A.Anderson, Dep. of Soil and Environ.
Sci., Univ. of California, Riverside, CA92521. Received 29 Apr. 1993.
*Corresponding author ([email protected]).
Published in J. Environ. Qual. 23:944-954(1994).
944
contained concentrations of As, B, Mo,Se, U, V, and
SO~-that weresignificantly elevated comparedto those
of ocean water, and Bradford et al. (1990) reported
potentially toxic concentrations of U, Mo, and V in
agricultural drainagewaters that wereelevated relative
to both the Salton Sea and MonoLake. Quite often, the
evaporation basins are exposedto cyclic evaporative
salinization fromwhichextensive evaporite deposits are
produced. Tanji and Dahlgren(1990, unpublisheddata)
found that ponds similar in evaporite mineralogywere
also similar in their chemicalcomposition,andthat evaporite mineralogycould be predicted by the brine chemistry model CSALT
(Smith, 1989).
The RWQCB
investigation demonstratedthat V, unlike
the other elementsstudied, exhibited lower concentrations in the pondedbasin water whencomparedwith the
inlet water. In natural waters the chemistry of V is
dominated by V(IV) (e.g., 2+ and VO[OH]÷) and
V(V) (e.g., H2VO~-and HVO42-)(Wehrli and Stumm,
1989). Thefate of V in evaporationbasins is thus governed by the pHand redox status of the water column,
whichultimately control the chemicalbehavior of V in
the basins. Uncertaintiesin the fate of Vand other trace
oxyanionsin pond waters have delayed the development
of an additional 4000to 8100ha of evaporation ponds
in the San JoaquinValley (Chilcott et al., 1990; Tanji
et al., 1992). Thus, an understandingof trace element
behavior and evaporite mineralogyof evaporationponds
is neededto properly managethese drainage water disposalfacilities.
Thebasic principles of the chemicalevolutionof saline
waters during evaporation in simple systems are described by the Hardie-Eugster model(Hardie and Eugster, 1970). Animportant concept of the modelis that
of a chemical divide, a point at which elements are
removedfrom solution whenprecipitation begins, resuiting in a changein relative elementalconcentrations.
The compositionof a water during evaporation maypass
througha successionof chemicaldivides that ultimately
control the compositionof the resulting brine. In most
natural waters, CaCO3is the first mineralto precipitate
and cause a chemicaldivide (Drever, 1988). The chemical evolution of the water followingCaCO3
precipitation
will dependon the ratio, in molesof charge, of Ca2÷/
- + COl-)
(HCOj- + CO~-). If the ratio Ca2÷/(HCOA
exceeds unity, a significant fraction of the carbonate
species will be removedfrom solution due to continued
Abbreviations: RWQCB,
Regional Water Quality Control Board; CF,
concentration factor; SI, saturation index; XRD,x-ray diffraction; IR,
infrared; EC, electrical conductivity.
945
LEVYET AL.: MINERAL
PRECIPITATION
& TRACE
OXYANION
BEHAVIOR
Table1. Selectedchemical
characterist".cs
of the syntheticinlet watersusedin the study.~"
pH
Water
Ca
Na
CI
SO4
Mo
Si
V(IV) V(V)
U
Mg
-t
-1
mm01L
tttnolL
WaterI
10.0
16.5
121 113
25.2 20.2
23.1 3.34
3.53 3.56 8.82§
Water
II
0.668
2.13 145
68.5
26.9
7.81 34.1 3.14 3.14 1.60 8.00
Dataarefortheaverage
of six replications
determined
byanalysis.Vanadium
valuesareaveraged
fromthreereplicates.
Calculated
byCSALT
fromtotal alkalinityandusinga PCO2
of 34.45Pa.
CaCO3precipitation,
and the pH of the solution will
approach neutral. On the other hand, if the ratio Ca2÷/
(HCO~-+ COl-) is less than unity after CaCO3precipitation, essentially all Ca2÷ will be removedfrom solution
as CaCO3,
and the solution will evolveinto a high alkalinity,
high pH brine. The Hardie-Eugster model also predicts
the precipitation of either sepiolite (Mg2Si3OT.5OH"
3H20),
or gypsum (CaSO4"2H20), depending on the composition
of the solution following the first chemical divide
(Drever, 1988).
As presented, the Hardie-Eugster model is primarily
useful for predicting the chemical evolution of waters
whose solution composition is dominated by alkali and
alkaline earth metals during evaporation. The concept
of a chemical divide, however, is also useful when
studying the behavior of dissolved trace oxyanions during
evaporation. Dissolved trace oxyanions mayexhibit conservative behavior during evaporation, with no losses of
the element from solution due to precipitation, coprecipitation, or adsorption. Alternatively, trace oxyanions may
exhibit nonconservative behavior, i.e., concentrations
that deviate from what would be predicted from evapoconcentration alone. Nonconservative behavior may result from precipitation as discrete mineral phases, coprecipitation
with commonevaporites, carbonates, and
aluminosilicates, or adsorption to mineral surfaces. Element concentrations mayalso exceed that predicted from
evapoconcentration alone due to mineral dissolution or
desorption from mineral surfaces in heterogeneous systems.
The objectives of this study were to evaluate the solution dynamics of mineral-forming elements and trace
oxyanions (U, Mo, and V) during evaporation of saline
waters whose chemical compositions evolve according
to two distinct chemical divides, and to characterize the
evaporite minerals formed from the complete evaporation
of these waters. This study is useful as an initial approach
to elucidating the factors responsible for the behavior
of trace oxyanions and their possible interaction with
evaporite minerals in agricultural evaporation basins.
The information obtained will be useful whenestimating
the chemical composition of basin waters from their inlet
water concentrations, predicting the types of evaporite
minerals that will be formed during periodic episodes
of dryness, and assessing techniques for the disposal of
accumulated evaporite minerals.
MATERIALS AND METHODS
Synthetic inlet waters were prepared using Fisher A.C.S.
Certified salts and deionized-distilled water. Fourliters of
each water matching the chemical compositions of the Pond
16 inlet water at Morris Farms(Water I), and the Pond
EC
-I
dSm
13.2
12.2
Alkalinity
-1
mmo~L
11.3
21.6
inlet water at the 4-J Corporation(WaterII), from the San
JoaquinValleyof California (Chilcott et al., 1990)wereprepared in triplicate (Table 1). Uraniumwas added as UO2
(NO3)2"6H20 and Moas Na2MoO4.TwoV treatments were
prepared from either Na3VO4
(V[V])or VOSO4"
3H20(V[IV]),
andthe concentrationfor bothspecies adjustedto a representative value of 4.0 ~tmolL-1. Silicon wasaddedas Na2SiO3
and
adjusted to the level of 100 IxmolL-~ supported by quartz
(Lindsay, 1979). The major elements (e.g., Ca, Mg,Na)
addedas chloride, sulfate, or bicarbonatesalts. Thesynthetic
waters were then placed in 4-L capacity, acid-washedpolyethylene beakers and allowed to evaporate in a laboratory fume
hood. Twenty-fivemilliliter sampleswere periodically withdrawnfor chemicalanalysis, and the solutions wereevaporated
to dryness(approximately30 d). The concentrationfactor (CF)
was recorded as the ratio of the initial mass(corrected for
sampling)to the massat the time of sampling.
All samples were immediatelyfiltered through an acidwashed, deionized-H20-rinsed, 0.2 ttm polycarbonate membranefilter and thenanalyzedfor total alkalinity andelectrical
conductivity(EC). Electrical conductivity wasdeterminedand
corrected to the standardtemperatureof 25°C.Alkalinity was
determinedby titration underair to pH4.4 with standardized
0.025 MH2SO4. The samples were then acidified to a pH <
2 with high purity HNO3
and stored at 4°C prior to analysis
for Ca, Mg,Na, Mo, and Si using inductively coupled plasma
atomic emissionspectrometry(Soltanpour et al., 1982). Uraniumin solution was determinedby inductively coupled plasma
massspectroscopyusing Bi as an internal standard (Toole et
al., 1990; Shiraishi et al., 1991). Vanadium
was determined
by a gallic acid.oxidation technique (Fishmanand Skougstad,
2- were determinedby ion chromatog1964). Chloride and SO~
raphyusing conductivitydetection.
Saturation indices (Sis) were determinedfor manycommon
evaporites by computingthe logarithmof the saturation ratio,
wherethe saturation ratio is defined as the ratio of the ion
activity product to the solubility product. The geochemical
models MINTEQA2
(USEPA, 1991) and CSALT(Smith,
1989) were used to speciate solutions and calculate Sis. The
model MINTEQA2,
which utilizes the Davies equation for
ionic strength corrections and thus is appropriate for waters
of modest ionic strength (_<0.50 mol kg-l), was used
calculate Sis and determinedominanttrace dementspecies of
WaterI for CFsfrom 1.0 to 2.7 and of WaterII for CFsfrom
1.0 to 2.7. The modelCSALT,
whichuses the Pitzer equations,
was used to speciate higher ionic strength solutions (waters
with CFs> 2.7).
The solution pHof water II wasmeasuredwith a combination
glass-reference electrode immediatelyafter sampling.Thesolution pHof the WaterI was both measuredwith a combination
glass-reference electrode and calculated by CSALT
using total
alkalinity and a PCO2 of 34.45 Pa. Calculated pHvalues for
Water I are reported here and were used in all equilibrium
calculations due to inconsistent pHmeasurements
of this water,
most likely a result of the higher ionic strength of WaterI
(Galster, 1991).
Elementaldata wereplotted as log of the relative concentration (C/Co)versus log CF. Becausethe CFis undefinedat
946
J. ENVIRON.
QUAL.,VOL.23, SEPTEMBER-OCTOBER
1994
~-~ 200
7
’---
aaaaa Water I
~ Water
II
Water I
Water II
~
~50
_~I008
o
o
6
>,
2
0
’~ 50
On
A
~
0
4
6
o
~ ~ oo 5 o~oo
oo
0
6
2
10
ConcentraLion facLor
4
6
8
2
10
Concentration f~ctor
9.40
8.80
Water I
8.40
9.00
oooooWater II
o
0
B°
oB
O
0
~o_ 8.60
o
[]
8.00
8.20
°° C
7.60
2
4
6
8
10
Concentrotion foctor
D
2
2
4
6
8
10
Concenf.rotion foctor
Fig. 1. Electrical conductivity,total alkalinity, andpHas a functionof concentration
factor for WatersI andH.
abscissae begin at a CFof 1 such that the plots have an apparent
intercept of 1.0. Thetheoretical line for an elementthat exhibits
conservative behavior is the line on this plot with a slope (b)
equal to 1.0 and a y intercept of 0. The Student’s t was used
to test the null hypothesis that b= 1 (Ho: b= 1) for a line that
passes through the origin (Steel and Torrie, 1980). Regressions
of element relative concentrations were also tested against C1
(H0: b=ba) using the Behrens-Fisher Test (Snedecor
Cochran, 1980).
Samplesof precipitates were collected after the waters evaporated to dryness, and then finely ground in a mortar and
pestle. The samples were prepared for x-ray diffraction (XRD)
as random powder mounts (Jackson, 1979) and analyzed with
a Siemens D-500 diffractometer using Cu-Kct radiation with
a graphite crystal monochromator.Mineral phases were identified with XRDby calculating distances between diffraction
planes from the recorded diffraction peaks and comparing them
with reference standards (JCPDS, 1993).
Precipitates were rinsed by leaching approximately 15 g of
sample with 100 mLof distilled H20to aid in the identification
of the minor mineral components. Single-beam infrared (IR)
spectra were obtained with a Mattson Cygnus 100 FTIR spectrometer equipped with a mercury-cadmium-telluride detector.
Thirty-two spectra were averaged at 4 cm-~ resolution referenced against the empty sample compartment.
RESULTS AND DISCUSSION
Samplesfrom WaterI collected during this experiment
represented CFs from 1.0 to 52 and produced ionic
strength values ranging from 0.22 mol kg-~ to 6.9 mol
kg-1. Samplesfrom Water II represented CFsranging
from 1.0 to 34 and yielded respective ionic strength
values from 0.16 mol k&~ to 5.2 mol kg-1. Electrical
conductivity, total alkalinity, and pHvalues for Waters
I and II as a function of the CFare shownin Fig. 1.
Electrical conductivity values (Fig. 1A) ranged from
15.6 to 143 dS m-1 -~
in WaterI and 12.2 to 160 dS m
in WaterII. The trend in ECis characteristic of the
exponentialconcentrationof water observedduring evaporation. TheECtended to deviate froman ideal exponential curve above CFs of approximately 6, however, as
a result of ion pairing, precipitation, and reduction in
the activity of n20 during evaporation.
Total alkalinity (Fig. 1B) ranged from 3.0 to 11.2
mmolc L-1 in Water I and decreased with increasing
evapoconcentrationup to a CFof approximately6. Alkalinity increased, however,uponfurther evapoconcentration of WaterI from CFsof 6 to 52. Total alkalinity in
LEVY ET AL.:
MINERAL PRECIPITATION
947
& TRACE OXYANION BEHAVIOR
6
ooooo Ca
" " " " " Mg
4
.Y
ooooo No
2
~
o10
o
15
I
15 0
4
2
2
4
6 8
2
~0
Concentration factor
4
2
4
68
2
4
10
Concentration f~ctor
ooooo Mo
" " " " " Si
GlOa
(_~10
C
2
4
6 8
2
~0
Concentration factor
4
2
4
6
8
2
I0
Concentration factor
Fig. 2. Relative concentrations (C/Co) of selected elements as a function of concentration factor for Water I. The solid Hne represents
line with a slope (b) of 1.0 and a intercept of 0.
Water II ranged from 21.2 to 541 mmolc L-~, and
increased steadily duringthe entire evaporativeperiod.
The pHvalues calculated by CSALT
for WaterI (Fig.
1 C) decreasedwith increasing evapoconcentrationas the
solution alkalinity also decreased(Fig. 1B). Themodel
CSALT
failed to converge for the Water I data when
calculation of pHwas attemptedat CFsgreater than or
equal to approximately 20, even though the solution
composition was electrically neutral and the ionic
strength wasless than the limit imposedby the Pitzer
-1).
equations (20 molkg
The measuredpH values for Water II (Fig. 1D) increasedwith increasing evapoconcentration,
as alkalinity
also increased (Fig. 1B), up to CFsof approximately
Uponfurther evaporation, however, the pHof WaterII
beganto decrease (Fig. 1D), although total alkalinity
continuedto increase. Anincrease in pHmust accompany
an increase in alkalinity, unless the partial pressure of
CO2is increased, upon which it is possible for the
pH to decrease with increasing alkalinity (Stummand
Morgan,1981). Thus, the decrease in the measuredpH
a theoretical
values of WaterII with increasing alkalinity beyondCFs
of 5 are indicative of an alkaline error (ApH),in which
the glass electrode respondsto singly chargedcations
that are present at high concentrations,as well as Hions
(Galster, 1991). Thealkaline error increases with both
increasing concentrationsof the interfering cation and
increasing pHabovea pHof 9.0. Themagnitudeof this
effect as influenced by Na ions is consistent with pH
changes that we have observed in the laboratory with
Na-spikedbuffer solutions (e.g., ApH= -0.20 at pH
9.50 and [Na] = 1.2 M).
Calciumloss from solution in WaterI (Fig. 2A) and
WaterII (Fig. 3A) is indicated by values of C/Cothat
lie belowthe theoretical line for conservativebehavior.
Theslopes (b) for regression throughthe origin of log
C/Coas a function of CF for the elements studied are
shownin Table2. Theb values for Ca in Table2 indicate
nonconservativebehavior for Ca that wassignificantly
different than the theoretical line.
Thefirst chemicaldivide caused by CaCO3precipitation (Drever, 1988) is given by the reaction:
948
J.
ENVIRON. QUAL., VOL. 23,
SEPTEMBER-OCTOBER1994
ooooo Co
3
ooooo ~a
//
oi%
n
~ []
121
o
o
A
0.01
2
4
6 8
2
Concentration
o
18!
4
2
2
10
Concentration foctor
f~ctor
4.
o I0~
o 10~
o
4- 6 8
4
I
1
2
Fig. 3. Relative concentrations (C/Co) of selected elements as a function of concentration
line with a slope (b) of 1.0 and a y intercept of
Ca 2+
4--
2HCO~- #
CaCO3
4-
CO2 4-
H20
Water
6 8
2
10
Concentration f~ctor
factor for Water II. The solid line represents
6
¢
a theoretical
second chemicaldivide caused by gypsumprecipitation,
resulting in CaCO3dissolution and subsequentincreases
in HCO~-and CO32-.
Calculated cumulative losses of Ca, Mg, and SO42from solution during evapoconcentration of Waters I
and II (data not shown)were consistent with the trends
discussed above. The calculations showedthat in Water
I, 25 mmolof Ca were precipitated from CFsof 1.0 to
2.5 (first chemicaldivide), after whichCa losses from
solution remained relatively constant up to a CF of
approximately 7. Correspondingcalculations for SO~indicate conservative behavior for SO~-from CFs of
1.0 to 7, suggesting that Ca losses between CFs of
1.0 to 2.5 were due to CaCO3, rather than CaSO4,
[1]
Thesteady.increase in alkalinity observedduringevapoconcentration of WaterII (Fig. 1B) is a consequence
- + CO23
-) being less than unity
the ratio Ca2÷/(HCO~
followingCaCO3precipitation, the first chemicaldivide
occurring betweenCFsof 1.0 to 1.22. The ratio Ca2÷/
-) in Water I exceeded unity as CaCO3
(HCO~-+ CO23
precipitation caused the first chemicaldivide between
CFsof 1.0 to 1.17. Thus,total alkalinity decreased(Fig.
1B) with increasing evapoconcentration up to a CF of
approximately6. Further evapoconcentrationresulted in
an increase in the alkalinity in WaterI from CFsof 6
to 52 (Fig. 1B). Theincrease in alkalinity reflects
Table 2. Slope values (b) for regression
factor. Asterisks denote significantly
4
through the origin of relative
concentration
(C/Co) for the element as a function
different
slopes relative to a theoretical
line for conservative behavior.
of concentration
Ca
Mg
Na
Cl
SO4
Mo
Si
V(IV)
V(V)
U
Water I
0.154"*
0.669**
0.630**
0.829**
0.385**
0.760**
0.940**
1.03"*
0.849**
0.540**
1.08"*
=
0.993
0.831"*
0.025**
0.870**
=
1.01
0.840**
Water II
0.276**
=
0.988
*, ** Significant
at P = 0.05 and P = 0.01, respectively,
ns = nonsignificant.
0.903**
LEVY ET AL.:
Table 3. Saturation indices (SI)~- of various minerals calculated
Water
Conc.factor
Calcite
(CaCO3)
WaterI
1.0
1.6
2.7
3.1
4.2
6.0
8.4
13
20
1.0
1.7
2.7
3.3
4.5
6.5
8.9
16
19
- 0.493
0.845
0.780
0.569
0.551
0.562
0.022
0.398
0.326
0.500
1.56
0.381
0.294
0.162
0.265
0.147
0.231
0.413
WaterII
Aragonite
(CaCO~)
- 0.680
0.658
0.593
0.382
0.364
0.375
- O.165
0.211
0.139
0.313
1.37
0.194
0.100
- 0.025
0.078
- 0.040
0.044
0.226
949
MINERALPRECIPITATION & TRACE OXYANIONBEHAVIOR
Gypsum
(CaSO,-2H20)
-
-
0.351
0.339
0.216
0.170
0.223
0.530
0.209
0.059
0.021
1.39
1.35
2.82
3.10
3.48
3.55
3.96
4.16
3.82
by CSALTfor selected concentration factors for WaterI and Water H.
Nesquehonite
Halite
Mirabilite
Nahcolite
Thenardite
(NaCl)
(Na2SO+-10H20) (NaHCO3)
(Na2SO+)
(MgCO~-3H20)
-
3.64
3.48
2.80
2.75
2.97
2.48
1.88
1.78
1.69
3.90
3.63
3.13
2.99
2.75
2.43
2.18
1.64
1.46
- 3.20
- 2.81
- 2.15
- 2.08
- 2.21
- 1.67
- 1.23
- 1.03
- 0.981
-2.95
-2.57
- 1.87
- 1.90
- 1.66
- 1.38
- 1.23
- 0.922
- 0.762
- 4.27
- 3.41
- 3.16
- 3.25
- 3.30
- 3.15
- 3.24
- 2.87
- 2.85
-2.27
- 1.93
- 1.53
- 1.38
- 1.07
- 0.801
- 0.408
0.105
0.129
-
4.09
3.70
3.03
2.88
3.00
2.50
1.96
1.74
1.66
3.89
3.51
2.81
2.78
2.51
2.20
1.99
1.51
1.30
-
3.46
1.86
1.88
2.09
1.96
1.91
2.55
1.76
1.67
1.29
1.14
0.540
0.960
0.869
0.825
0.873
0.803
0.648
~" SI = log (ion activity product/solubility product).
precipitation. Calciumprecipitation resumedfollowing
evapoconcentrationbeyonda CFof 7, with stoichiomet2- also beinglost, providingfurther
ric equivalentsof SO~
evidence for a second chemicaldivide caused by gypsum
precipitation.
Cumulativeelemental losses calculated for WaterII
showedthat 3 mmolof Ca were precipitated from CFs
of 1.0 to 2.0 (first chemicaldivide) and subsequently
remained constant throughout the CFs sampled. Additionally, trends in SO~-provided no evidence for a
second chemicaldivide causedby gypsumprecipitation.
Thus,gypsum
precipitation is not expectedin evaporating
- + COl-) is less
waters where the ratio Ca2+/(HCO~
than 1, due to the lowsoluble Caconcentrations(Drever,
1988).
Saturation indices for calcite (CaCO3),aragonite
(CaCO3), and gypsum (CaSO4"2H20) computed
CSALTand MINTEQA2
for both waters are given in
Table 3. Positive Sis for calcite and aragonite were
observedduring evapoconcentrationof WaterI, indicating supersaturationof the solution with respect to these
solid phases. TheSis in Table3 showa net dissolution of
CaCO3
at a CFof 8.4 wherethe SI for calcite approaches
equilibrium and the SI for aragonite becomesnegative.
Theseobservations are consistent with the increase in
alkalinity that occurred beyonda CFof 6 (Fig. 1B) and
the positive Sis for gypsumin WaterI beyonda CF of
6. Positive Sis for calcite and aragonite(Table3) indicated that WaterII remainedsupersaturatedwith respect
to these phases throughout the course of evaporation.
WaterII remainedundersaturated, however,with respect
to gypsumduring the entire evaporativeperiod, consistent with the elementallosses discussedabove.It should
be noted that positive Sis do not providedirect evidence
for solid phase formation, but certainly indicate when
precipitation of the phaseis possible. Supersaturation,
as opposedto equilibrium, of the waters with respect to
CaCO3
is attributed to slowcalcite precipitation kinetics
(Amrheinand Suarez, 1987).
X-raydiffraction analysisof evaporitesproduced
diffraction peakscorrespondingto characteristic d-spacingsfor
gypsumin WaterI evaporite samples (Fig. 4), which
in agreementwith positive Sis observedfor gypsumin
Water I beyond CFs of 5 (Table 3). Gypsumwas not
detected in evaporite samples from WaterII, whichis
consistent with negative Sis for gypsumin Water II
(Table3). Calcite and aragonite werenot identified with
XRDin the Water I evaporite samples, although Sis
indicated supersaturationwith respect to these minerals
(Table3). Calcite andaragonite werenot identified with
XRD
in WaterII either, and similarly the Sis for calcite
and aragonite in WaterII werepositive (Table 3).
Theinability to detect calcite or aragonitein the evaporites by XRDindicates that the mineral is present in
amountsthat are belowthe detection limit of the diffractometer (generally 5-10%by weight). It is important
note that peakintensities for a given mineral will vary
as a functionof factorsthat affect diffractionpeakintensities, suchas particle size, crystal imperfections,chemical
composition,self-absorption, crystal orientation, and the
numberof correspondingdiffraction planes in the sample
(Jackson, 1979). Thus, standard and accurate detection
limits for evaporite minerals cannot be established. If
all of the Ca in the waters wereprecipitated as CaCO3,
this mineralwouldconstitute only 6 %of the solid collected fromWaterI, and less than 1%of the solid from
WaterII. Mineralogicalanalyses of actual evaporation
basin precipitates indicated that CaCO3
was present in
small amounts (K. Tanji and R.A. Dahlgren, 1990,
unpublisheddata). Therefore, it seemsreasonable that
no forms of CaCO3were detected with XRDanalysis.
Infrared spectra obtained for evaporite samplesfrom
WaterI producedabsorption bands at 1486, 1094, 872,
712, and 700 cm-1, which have been attributed to the
carbonate componentof aragonite (Doner and Lynn,
1989). Trace amountsof calcite werealso present, indicated by additional bands at 872 and 712 cm-~. Water
II evaporites produced IR absorption bands at 1490,
1085, 880, 712, and 700 cm-1, indicating the presence
of aragonite. There were no IR absorption bands correspondingto calcite in Water
Values of b for Mgin Waters I and II (Table 2)
indicate behaviorthat wassignificantly different than the
theoretical line for conservative behavior. Nonconserva-
950
J. ENVIRON.QUAL., VOL. 23, SEPTEMBER-OCTOBER
1994
H
H
0.326
0.282
A
H
0.200
G
BI
0.763 nm
G
BI
0.178
B
H
Th
0.464 nm
H
0.283
0.200
0.265
Th/Tr
0.308
0.187
H \
mr
0.221
T~ Tr
11.0
31.0
51.0
degrees2e
Fig. 4. X-ray diffractograms of evaporites fromthe waters studied.
(a) Water I. Bl = bloedite,
G = gypsum, H = halite,
Th =
thenardite. (b) Water H. H = halite, Th = thenardite,
Tr =
trona.
five behaviorof Mgcommonly
results fromsolid solution
formationof Mgin calcite to formMg-calciteor adsorption to calcite surfaces (Gac et al., 1978; Stummand
Morgan,1981; Miyamotoand Pingitore, 1992). Changes
in lattice parameters as detected by XRDare useful in
distinguishing betweenMg-calcite formationand adsorption of Mgonto calcite surfaces. There were no forms
of CaCO3detectable with XRD,however, so we could
not makethis distinction. CumulativeMglosses in Water
I paralleled those of Ca up to a CF of 6 (approximate
Ca/Mgloss ratio of 2.5:1), after which negative Mg
losses werecalculated for the remainderof the samples.
This suggests adsorption of Mgonto calcite surfaces or
formation of Mg-calcite, whichbegan to dissolve as
gypsumprecipitation occurred at CFs greater than 6
(previously discussed). CumulativeMglosses paralleled
those of Ca in WaterII also. Magnesium
and Ca losses
in this water increased up to a CFof 2.0 (approximate
Ca/Mgloss ratio of 2.3:l),and remainedconstant for
the remainderof the samples.
Saturation indices for various Mg-containingcompounds were calculated by CSALTand MINTEQA2
for
WatersI and II (Table3). CalculatedSis for the possible
evaporites nesquehonite (MgCO3"3H20),
bischofite
(MgCI2"6H20),bloedite (Na2Mg[SOa]2"4H20),
epsomite (MgSOa’7H20),hexahydrite (MgSO4.6H20),
kieserite (MgSO4H20)indicated that the solutions were
undersaturatedwith respect to these phases, further suggesting coprecipitation with or adsorption to calcite as
the primaryloss mechanism,
at least duringearlier stages
of evaporation.
X-ray diffraction analysis of evaporite samplesfrom
Water I produced diffraction peaks corresponding to
characteristic d-spacingsfor bloedite (Fig. 4), although
Sis for bloedite indicated undersaturationfor all solutions. Apparently, bloedite precipitation occurred between the time whenthe last samplewas collected and
the solutions evaporated to dryness. Bloedite has been
identified as a dominantevaporite component
in several
evaporationpondsin the SanJoaquinValley(K. Tanji and
R.A. Dahlgren, 1990, unpublisheddata). No definitive
evidence for precipitation of any Mg-containingcompounds was found by XRDanalysis of evaporites from
WaterII. Saturation indices for bloedite in WaterII (not
shown)were negative, again suggesting that nonconservative behaviorof Mgin WaterII results fromMg-calcite
precipitation or Mgadsorption to calcite surfaces.
Table 2 showsthat Nain both watersexhibited nonconservative behavior denoted by a b value that deviated
significantly fromthe theoretical line. Solutionlosses of
Naobservedduring the samplingperiod maybe attributed
to incorporation of Nainto CaCO3
(Busenbergand Plummer, 1985). Halite (NaCI), bloedite, and thenardite
(Na2SO4)were identified by XRDin evaporite samples
from WaterI (Fig. 4), although Sis calculated by CSALT
and MINTEQA2
showsthat the water remainedundersaturated with respect to halite and thenardite up to a CF
of 20 (Table3). This indicates that halite andthenardite
precipitation mayhave only occurred betweenthe time
whenthe last sample was collected and the solution
evaporatedto dryness. Themodelresults should be interpreted with caution, however.Onget al. (1992) consistently observed negative Sis calculated by CSALT
for
halite and thenardite in pondwaters that were knownto
be in contact with these minerals.
Halite and trona (NaaH[CO3]2"
2H20)were identified
by XRDin evaporites from WaterII (Fig. 4), although
Sis for halite (Table 3) and trona (not shown)indicated
undersaturation from CFsof 1.0 to 19. Positive Sis for
nahcolite (NaHCO3)
were observed for Water II from
CFsof 16 to 19 (Table3), but nahcolite wasnot identified
with XRD.Nahcolite mayhave been present below XRD
detection, or mayhave not formedat all. Nahcolitewas
ubiquitous in the salt deposits of several evaporation
951
LEVY ET AL.: MINERALPRECIPITATION & TRACEOXYANIONBEHAVIOR
Table 4. Calculated distributionS"
of U, V, and Mospecies for the waters used in this study.
Percentage distn%ution of species
WaterI
Water H
Conc. factor
UO2(CO3)~-
UO2(CO3)~-
I-IV20~
-
H2VO4
-
1.0
1.2
1.3
1.4
1.6
2.1
2.7
1.0
1.2
1.4
1.6
2.0
2.5
2.7
5.3
8.0
14.4
17.3
13.8
22.1
14.5
2.1
0.3
0.2
0.2
0.1
0.1
0.1
94.7
92.0
85.5
82.5
86.1
77.6
85.4
97.9
99.7
99.8
99.8
99.9
99.9
99.9
63.3
65.9
68.3
68.0
72.1
72.3
74.5
67.1
50.3
49.7
47.3
48.3
51.9
45.2
16.3
16.3
19.5
21.6
15.8
18.9
14.0
12.2
4.7
3.4
2.8
2.2
1.9
1.5
20.4
17.7
12.2
10.3
12.1
8.80
11.5
20.7
44.9
46.8
49.8
49.4
46
53.2
100
100
100
100
100
100
100
100
100
100
100
100
100
100
-1.
Distribution was calculated by MINTEQA2
for samples with ionic strengths -<0.50 mol kg
ponds in the San Joaquin Valley (K. Tanji and R.A.
Dahlgren, 1990, unpublisheddata).
The nonconservativebehavior of Moin WatersI and
II is shownin Fig. 2C and 3C. The b values for Mo
(Table 2) indeed indicate nonconservativebehavior and
showsimilar slopes for Moin both waters. ModelMINTEQA2
calculations indicate that 100%of the dissolved
Moin WatersI and II existed as MoO~4(Table 4). This
is expectedin neutral to alkaline waters whereH2MoO4
(pI~ = 4.00) should exist primarily as the dissociated
form (Lindsay, 1979). Positive Sis were identified
MINTEQA2
calculations for CaMoO4
in Water I corresponding to CFs from 1.32 to 2.70, although XRD
evidence for CaMoO4precipitation would not be expecteddue to the low Moconcentrationof WaterI (Table
1). Negative Sis for CaMoO4
were obtained for Water
II, and negative Sis for MgMoO4
were calculated for
both waters. Molybdenum
loss in Waters I and II may
result from adsorption to calcite surfaces (Stummand
Morgan,1981), coprecipitation with other evaporite minerals, or from precipitation as CaMoO4
in WaterI. The
association of Mowith evaporites has been demonstrated
by Tanji and Dahlgren (1990, unpublished data), who
observed a slight enrichment of Moin thenarditedominatedevaporite deposits from evaporation basins.
Silica behavednonconservativelyduring evaporation
of WaterI (Fig. 2C) and WaterII (Fig. 3C), deviating
significantly fromthe theoretical line as indicated by the
b value in Table 2. Silica loss from saline irrigation
waters wasstudied in detail by Eaton et al. (1968) who
concludedthat precipitation as an amorphous
Mgsilicate
wasthe mechanism
of Si loss fromsolution. Precipitation
of Si as amorphous$iO2(Tanji and Valoppi, 1989)
as sepiolite (Drever,1988)are other possible mechanisms
of Si loss fromevaporatingsaline waters.
Thedata in Table2 showthat the slope of the C1line
for WaterI wassignificantly lowerthan the theoretical
line, indicating nonconservativebehavior. This behavior
is most likely due to halite precipitation, even though
Sis for halite indicated undersaturation, as did those
observedby Onget al. (1992) for waters in contact with
evaporitedeposits containinghalite. Furthermore,halite
wasthe only chloride mineral identified in evaporites
fromthree evaporationpondsin the San Joaquin Valley
(Tanji and Dahlgren,1990). Thenonconservativebehavior of SO~-(Table 2) in WaterI is attributed to gypsum
precipitation (previously discussed) and/or SO~-incorporation into calcite (Busenbergand Plummer,1985).
Chloride in WaterII exhibited conservative behavior
(i.e., the slope of the C1data wasnot significantlydifferent than the theoretical line [Table 2], althoughhalite
was eventually produced at very high CFs as the XRD
patterns show(Fig. 4). Variation in SO~4-concentrations
in WaterII resulted in a significant b (t = 2.83) value
greater than a theoretical b of 1.0 (Table2). Anyvalue
of b greater than 1.0 for regression throughthe origin
designatesan increase in concentrationgreater than can
be attributed to evapoconcentrationalone. The slight
increase in the slope of the SO~4-above that of the
theoretical line (Table2) can only be attributed to experi- data for Water
mentalerror. Thus, the fit of the SO24
II (Fig. 3D) is regarded as behaving conservatively,
although in a mannersimilar to CI-, SO~-eventually
precipitated as thenardite at high CFs(Fig. 4).
Compared
with the major elementsstudied, the overall
trend for V behavior in WaterI (Fig. 5A)and Water
(Fig 5B) appears to be conservative. Examinationof the
b values (Table 2), however,showthat in somecases the
V data wasclassified as nonconservative.Furthermore,
these values indicate behaviorthat is inconsistent with
the water type, the initial oxidationstate of V, andthe
null hypothesisthat wastested.
Vanadium
concentrations for CFsof 1.0 to 2.7 listed
in Table 4 ranged from 170 to 360 ~tg L-~ for Water
I, and 160 to 440 Ixg L-~ for WaterII. Table 4 shows
the dominant V species in solution as calculated by
MINTEQA2
should be HV20~-.Saturation indices for
V205,Ca3(VO4)2,Ca2V2OT,
and other metal vanadates
calculated by MINTEQA2
indicated that the solutions
wereundersaturatedwith respect to these mineralphases.
Sadiq (1988) proposed that calcium vanadate,
(VO3)2(c), in equilibrium with CaCO3(c) controls the
solubility of V in oxidizing marine environments(pc
pH> 9), although the solubility product wasnot met in
our study.
Vanadium(IV)
should initially exist as the species
952
J.
ENVIRON. QUAL., VOL. 23,
10
WaterII
V(V)
2
[] o Do [] V(IV)
n
c.? o
2
2
,r
6 8
2
10
ConcentroLion factor
~,,aa~ Water I
ooooo Water II
~
/o
o~o°°
2
~le~
2
~ ~-’~
~ .......
,
....
2
10
Concentration fuctor
k
6 8
#
6
Fig. 5. Relative concentrations (C/Co) of V initially
added as V(IV),
V initially
added as V(V), and U for Waters I and 11. The solid
line represents a theoretical line with a slope (b) of 1.0 and a
intercept of 0.
SEPTEMBER-OCTOBER1994
H2V20~+ and V(OH)3+ in both of the waters studied as
calculated by the MINTEQA2
model. For the full evaporative period of approximately30 d, however,the behavior of V between the two treatments, V(IV)vs. V(V),
wouldnot be expected to differ significantly, because
the V(IV)wouldbe oxidized to V(V). Wehrli and Stumm
(1988) presented an empirical rate law for the homogenous oxygenationof V(IV)with a half-life of 4.3 d for
dissolved V(IV).
Basedon the qualitative observation of conservative
V behavior in this study, the lower concentrations of
V in ponded basin waters comparedwith agricultural
drainage inflow waters in the San Joaquin Valley noted
by Chilcott et al. (1990) are most likely a result
incorporation of V into pondsediments via adsorption
to pond sedimentsor complexationwith and settling of
humicsubstances,rather than precipitation or coprecipitation with evaporite minerals. Vanadium
has long been
recognizedas being associated with bituminousmaterials
and carbonaceousrocks (Premovi6et al., 1986; Breit
and Wanty, 1991; Wantyand Goldhaber, 1992), humic
substances (Szalay and Szil~gyi,1967; Goodmanand
Cheshire, 1975; McBride,1980; Templetonand Chasteen, 1980), and mineral surfaces (McBride,1979; Shieh
and Duedall, 1988; Wehrli and Stumm,1989).
Uraniumclearly exhibited nonconservative behavior
in WaterI, yet conservative behavior wasobservedfor
U in WaterII (Fig. 5C; Table 2). Modelcalculations
indicated undersaturated conditions in both waters for
UO3(c), gummite(an amorphousweathering product
UO2), I}-UO2(OH)2,rutherfordine (UO2CO3),uranophane, and schoepite. Uraniumwas shownto exist as
either UO2(CO3)~-or UO2(CO3)~-as determined
MINTEQA2
calculations.
In Water I, the species
UO2(CO3)34becameless abundant (Table 4) as CaCO3
precipitation causeda reduction in the solution pH(Eq.
[1]; Fig. 1C). In Water II, the species UO2(CO3)~becamemore abundant whenthe solution pH increased
(Fig. 1D), due to evapoconcentrationof alkalinity that
occurred when Ca/(HCOf+ CO]-) exceeded unity.
Duff and Amrhein(1992) studied U adsorption
soils andclay mineralsin solutions of synthetic drainage
waters similar to those of the San Joaquin Valley of
California and Showedthat maximum
adsorption of U
occurred in solutions with the lowest alkalinities and
pH values of approximately 8.6. Model(MINTEQA2)
calculations showedthe dominantsolution species of U
to be UO2(HPO4)22-,
undoubtedly due to the modest
concentrations found in agricultural drainage waters in
this region. Other soluble species were UO2(CO3)~and
UO2(CO3)~-.
Consideringthese results, in addition
the higher U concentrations and lower alkalinities of
WaterI, the nonconservativebehavior of U in WaterI
(Fig. 5C)maybe attributed to coprecipitationwith evaporite minerals. LowerUconcentrationsandhigher alkalinities contributedsignificantlyto the conservativebehavior
observedfor U in WaterII (Fig. 5C).
Regressions of element relative concentrations were
also tested against C1(H0: b = bcl) (data not shown).
Chlorideis traditionally regarded as behavingconservatively becauseit is not precipitated as a salt until very
LEVY ET AL.: MINERAL PRECIPITATION & TRACE OXYANION BEHAVIOR
high salinities are reached (Drever, 1988). In Water II,
Cl behavior was classified as conservative, so the results
from this statistical test were identical to those in Table
2. In Water I where Cl was shown to exhibit nonconservative behavior, V in the V(V) treatment 2and Mo were
not significantly different than Cl, and SO ." was significantly different only at P = 0.05.
SUMMARY AND CONCLUSIONS
This study has described the behavior of the dominant
alkali metals, alkaline earth metals, and trace oxyanions
(U, V, and Mo) in synthetic agricultural drainage waters
that followed two distinct pathways of the Hardie-Eugster
model. The chemical parameters of the waters studied
during evapoconcentration verified the concept of a
chemical divide outlined in the Hardie-Eugster model.
Thus, evaporating drainage waters in disposal basins in
the San Joaquin Valley of California are expected to
evolve into either: (i) high alkalinity, high
pH brines
when the inlet waters have ratios of Ca2+/(HCO3~ +
COl~) less than unity, or (ii) low alkalinity, nearly neutral
brines when the Ca2+2+/(HCOf + CO2")
ratio is greater
than unity. For Ca /(HCO3~ + CO2,') ratios slightly
less than 1.0, the precipitation of Mg becomes important
in controlling alkalinity levels in evaporating waters.
The initial precipitation of CaCOs during evapoconcentration caused the first chemical divide. Gypsum
precipitation
followed CaCOa precipitation in waters with
a Ca2+/(HCO3- + CO2") ratio greater than unity (Water
I). Evaporite minerals such as bloedite, thenardite, and
halite were found to precipitate at higher CFs. Additionally, carbonate minerals such as nahcolite and/or trona
formed during
evapoconcentration of waters that possessed a Ca2+/(HCO3~ + CO2.") less than unity.
In pure systems Mo exhibited nonconservative behavior resulting from precipitation as CaMoCu and/or coprecipitation with evaporite minerals. Vanadium behaved
conservatively during evaporation of the waters in the
absence of a sediment sink, such that V concentrations
could be predicted based on evapoconcentration alone.
Thus, observed reductions in V concentrations within San
Joaquin Valley evaporation pond basin waters compared
with inlet waters appear to be due to incorporation of V
into bottom sediments via adsorption to mineral surfaces,
settling humic substances, and biological uptake. The
fate of U in these waters depended on alkalinity and U
concentration. High alkalinities contributed favorably to
conservative behavior, while low alkalinities allowed for
significant incorporation of U into evaporite minerals.
Further studies would be useful to determine the effect
of pond acidification on trace element concentrations in
the basin waters. Acidification to reduce alkalinities below Ca concentrations would prevent pond waters from
evolving into high alkalinity brines, and reduce soluble
U levels in pond waters.
ACKNOWLEDGMENTS
The authors thank Gordon Bradford, Paula Bosserman, and
James Strong for their analytical assistance. This research was
953
partially supported by a grant from the U.C. Salinity Drainage
Task Force (Project 90-11).
954
J. ENVIRON. QUAL., VOL. 23, SEPTEMBER-OCTOBER 1994