-- ^---LI1L~F--
-^.x^l^-
II*II~
II^I~
-)~11.1-^_11 1-L. _LI
ANALYSES OF THE SULFUR SYSTEM IN WATERS
FROM THE GALAPAGOS RIDGE HYDROTHERMAL VENTS
by
SARAH S.
HUESTED
S.B., Massachusetts Institute of Technology
(1978)
SUBMITTED IN PARTIAL FULFILLMENT
OF THE REQUIREMENTS FOR THE
DEGREE OF
MASTER OF SCIENCE
at the
MASSACHUSETTS INSTITUTE OF TECHNOLOGY
SEPTEMBER, 1979
Signature of Author
De
,
of/Tarth & Planetary Sciences
Certified by .....
Thesis Supervisor
Accepted by ..............
....................................
arary
ChairmaYhr
I
MIT
Tr I4US INSTTUIE
I IRRARIES
Sciences Committee
~-~--~LL~II
--
--. 1IIYI
ll111
.(r-._r
~-i.-~L~--_l~l)
I
-2ANALYSES OF THE SULFUR SYSTEM IN WATERS
FROM THE GALAPAGOS RIDGE HYDROTHERMAL VENTS
by
SARAH S. HUESTED
Submitted to the Department of
Earth and Planetary Sciences
on September 1979 in partial fulfillment of the requirements
for the Degree of Master of Science
ABSTRACT
Samples collected from the Galapagos hydrothermal
vents show a decrease in sulfate with temperature resulting
in a flux of 5.3 x 1012 mol/year from the ocean assuming
that these vents are representative of a global process.
Sulfide concentrations increased with temperature (flux of
2.6 x 10 11-2.6 x 1012 mol/year to the ocean).
Reduced
species of sulfur including sulfite, thiosulfate, trithionate
and tetrathionate were found in minimal amounts, if at all.
A substantial portion of the iodine reactive species
(otherwise unaccounted for) was
hypothesized to be elemental
sulfur.
Reduced sulfur species were not found in the overlying
water column.
Name of Supervisor:
Professor John M. Edmond, M.I.T.
..
-
-
0
-3ACKNOWLEDGEMENTS
I am grateful to John M. Edmond for being my advisor
and providing me with insight and suggestions for my research.
Special thanks are due the many people who helped me
accomplish the research work, including Bob Stallard, Barry
Grant, Russ McDuff, and Chris Measures.
Thanks are also due
Loretta Tocio for typing a major part of this thesis.
I was supported by funds from ONR and NSF.
...
r ,.~..r-lll- --
--x~ -;-^-r---~x-I
ll
..._-~---rul-r~--l-~-L-LI~-LIICT
-Y-~L-~--
_~
-4TABLE OF CONTENTS
Page Number
Title Page.........................
.....
Abstract... ........................
000. 0 2
Acknowledgements ...................
.. 0..
3
Table of Contents..................
.0..
4
.....
6
1
List of Figures ....................
List of Tables .....................
CHAPTER 1 - Introduction...........
CHAPTER 2 - Experimental Section...
CHAPTER 3 - Results and Discussion.
References
.......................
..... 15
~1_.
~
---- -I---------------c
--*v---L------~l ---- '-I-----------
----*----II
--- --------ur ~-------------
.1-~IISS131Y+IYLI IIIII~II lrP~~IXI~
I*YU~YIWIYIIIIC liT-~e;~
Y^--~--YFC
-5-
List of Figures
Page Number
Figure
1.1
Hydrothermal Reduction and Oxida-
11
tion of Sulfur
1.2
Oxidation Reaction Pathways -
13
Acidic and Neutral Solutions
2.1
Sulfide Analytical Procedure
2.2
0-3
2.3
Sulfite Analytical Procedure
24
2.4
Thiosulfate, Tetrathionate, and
29
MM. Sulfide versus Absorbance
18
21
Trithionate Analytical Procedure
2.5
Thiocyanate versus Absorbance,
31
Method C, (CuSO4 )
2.6
PRA Preparation
37
3.1
Sulfide versus Silica
50
3.2
Sulfide versus Oxygen
52
3.3
Sulfate versus Silica
55
3.4
Sulfite versus Silica
57
3.5
Thiosulfate versus Silica
59
3.6
Iodine Reactive Species versus Silica
61
3.7
Iodine Reactive Species Minus Sulfide,
64
Thiosulfate, and Sulfite Versus
Silica
.~ , .~....1...
._I~.-,i~~-~----- -3~-~~
------- - --~~
-~---------
I--.
~e
-.-,.r~l-ir~
~ iL
~.~lr..-,~.
.,~x.~-I----------
L-~rre---------- ~. ..~,~x---..~~
-6List of Tables
Page Number
Table
2.1
Hydrothermal Vent Locations
16
2.2
Hydrocast Locations
17
2.3
Sulfide: Reagent Concentrations
20
and Dilutions
2.4
Sulfide Precision
23
2.5
Cyanolysis Reactions Involving
27
Thiosulfate, Tetrathionate, and
Trithionate
2.6
Tetrathionate and Trithionate
33
Precision
2.7
Summary of Method Standard
34
Errors
3.1
Concentrations of Sulfide, Sulfite,
42
Thiosulfate, Trithionate, Tetrathionate, Iodine Reactive Species,
Oxygen, and Silica for the
Galapagos Vents
3.2
Gradients and Fluxes for Sulfide
and Sulfate
47
x l-~-p--lr-.-.
..irr~-- - -...~.._^........~~.~~.I~llc
--- -I .....
~..
--.
-7-
CHAPTER 1
INTRODUCTION
i~-rXCr
-I
IIII^L~---~-
r -----
i
---
-
rsu
ru
I,-,,--, Y-
1~
-1-11-1.
-
Ma
L~L*~*~y~l~--~-~.r--~--
-8Over geologic time, sulfur chemistry in the oceans
is controlled by reactions occurring between water, rock
Under
and sediment phases (Goldhaber and Kaplan, 1974).
surface temperatures and pressures, for instance, it is
thermodynamically favorable for sulfate to be reduced by
organic matter to sulfide in the absence of oxygen.
However, this reaction is not observed, unless there is
biological intervention.
Biologically mediated reduction of sulfate might
occur in the water column in areas of the ocean where
the oxygen concentration approaches zero (e.g., the
eastern tropical Pacific).
This phenomena has not been
observed, however, probably because bacteria preferentially
reduce nitrate and nitrite before sulfate (Brewer, 1975,
Cline and Richards, 1972).
Chemical reduction of sulfate can occur in processes
such as reaction (1) when seawater and basalt react at high
+
-2
11 Fe 2 SiO 4 + SO 4 + 4H = 7 Fe304 + FeS2 + 11 SiO 2 + 2H20
fayalite
(1).
magnetite pyrite
temperature, as in the case of ridge crest hydrothermal
systems
(Bonatti, 1975), (Figure 1.1).
In experiments in
which basalt and seawater were reacted at 300
0 C,
much of
the sulfate present was incorporated in anhydrite, CaSO 4,
and some pyrite with perhaps 10% being reduced to aqueous
sulfide.
In many of these experiments, the total reduced
1I1 I
1
..~-r.,~~~.,~ ~...I-..~
~~
..,..~...~~~.1.~.,.~4
I
_II_
__.....,
.....~.~..~~,~.I~~~_.
U.sl
------- --r~-~-i~~ -- -u~~s~~-hpyuxllllL-L-LIL^~-L1C--
-9sulfur measured exceeded the amount of total sulfur available from seawater, indicating that sulfur was also being
leached from the rock (Mottl, et al., 1979).
Pyrite was
not found in hydrothermally altered pillow basalts from
the Mid-Atlantic Ridge which led to the conclusion that
the reduction of sulfate must be a minor process relative
to the total water flux through the crust (Humphris, et al.,
1978).
Another possible source for reduced sulfur in the
hydrothermal waters are volcanic gases liberated beneath
spreading centers that might be dissolved in the ascending
seawater (Bonatti, 1975).
With these discrepancies in the
literature between laboratory experiments, field observations, and theory, a more detailed look into the sulfur
chemistry of hydrothermal waters is needed.
Seawater, when it reacts with basalt at the ridge
crest, is heated to about 300 oC, decreases in pH, the
result of magnesium fixation, and potential electron
acceptors (oxygen, nitrate, sulfate, etc.) are consumed
resulting in a reducing solution containing hydrogen sulfide
(Edmond, et al., 1979a).
Ambient water entrained along
faults and cracks in the crust would lead to oxygenated
water mixing with these reduced fluids as they rise or
to oxidation of some of the pyrite (FeS2 ).
Large amounts
of sulfur-oxidizing bacteria living in and around the vents
would also contribute to the oxidation of the hydrothermal
-C--"----Y--e~~ -~---c-~-- x.---~..,.....
-III-CL
.I ..l.-a----- ----l-----lsLIL---
-10solution (Corliss, et al.,
1979),
(Figure 1.1).
With the
many redox reactions possible for sulfur under the above
conditions, it is possible that some of the intermediate
species might be measured in the hydrothermal fluids
(Figure 1.2).
Thermodynamic calculations indicate that
sulfate, sulfide, and/or elemental sulfur would be the
dominant species depending on the pH and pE.
However, the
system is not at thermodynamic equilibrium, so that some
of the intermediate thermodynamically unstable species
might be metastable for a sufficient period of time to
be sampled and measured.
Many variables, such as pH,
initial concentration of reduced sulfur, catalysts and
inhibitors, oxygen concentration, and relative reaction
Based on the
rates would influence the products found.
experimental work of Nelson, et al., 1977, Cline and
Richards, 1969, and Chen and Morris, 1972, the
intermediate species that might be expected
would include
elemental sulfur, thiosulfate, polythionate (in particular
tetrathionate and trithionate), and,-possibly, sulfite.
Therefore by measuring the concentrations of these species,
as well as those of sulfide and sulfate, and, as a cross
check, the species that react with iodine
(H2 S, S2 0 3 , SO 3,
polysulfides, and possibly So stabilized by S2 0 3 or polythionates), greater insight into the sulfur cycle in the
Galapagos hydrothermal waters would be obtained.
--~~~~-LI
-
---
~--.
--
----
C-~-----
-11-
Figure 1.1
Hydrothermal Reduction and Oxidation of Sulfur
I
I I
-..EL_^I~I~Y~-~-~-I~-~~ l-lll^-.L~----XI~-l_^.~_
Il^.~l~----- -~---i..
.
-13-
Figure 1.2
Oxidation Reaction Pathways - Acidic
and Neutral Solutions
(Nelson, et al., 1977)
-Z
os
Os
0s
-Z
0
S9.-0
1
9 0 cs
I
EOS k
-
s
-z
c
c os
-~~
0 zS
Sz
z Os
f
zEosle
q
IF
IIIP~--Csll~/ll(lllll~_I
---
--
-
--
_~---- LI^~IIIP
J---_ ~-PI~ ii)- .Il~
P1- ~ll^ls)~tllllr^I~_
1..~.
-15-
CHAPTER 2
EXPERIMENTAL SECTION
.I-~ ~. LI-.I--II~LY
-- --~--~.~._ .. ......~-~s~JI, ~.~~-
.~.^I--II-
~~......,-
^XI-^- III^
~ .~Y-Y~.-X__II-I"LLill'-_.-^II-. ~--"-"----~---I-4~-L I i^-
,
~-~-,
-----
-16Sampling
The Galapagos hydrothermal waters were collected using
the sampling system described by J.B. Corliss, et al., 1979,
mounted on the deep-sea submersible, Alvin.
"Rose Garden,"
areas were sampled;
Three vent
"East of Eden," and
"Mussel Beds" (Table 2.1).
Table 2.1
Dive #
Position
Vent area
Lon.
Lat.
'
Mussel Beds
0047.5
'
N.
8609.0
Rose Garden
0048.3
'
N.
86014.3
East of Eden
0047
'
N.
8602.5
'
W.
'
W.
W.
898, 902, 904
899, 900, 903, 905
901
The samplers were backfilled with nitrogen as samples
were removed.
Samples for sulfur species analyses were
collected in ground glass stoppered reagent bottles using a
plastic tube held at the bottom of the container.
The
sample was then allowed to flow into the bottle and overflow
to avoid oxygen entrainment.
Hydrocast samples from
approximately the same area (Table 2.2) were collected using
five liter OSU bottles.
. ^~
11-1-1---~1^-~--~-.rms1"----~
11
~~~
~ _~ ~_~____ ~~~~~_~-LrmqYL~^
.l.IrrIPlm-.*P
VI~I--~-L-~C~_lj~^L
- liis~sP
-17Table 2.2
Stn #
2012.1
86
Lon.
Lat.
'
N.
135
0038.8' N.
139
0047.7
'
N.
8502.2
'
W.
8604.8' W.
8607.7' W.
Analysis
All colorimetric measurements were made using a PerkinElmer 55E spectrophotometer with a 1 cm. quartz cell, except
for sulfide measurements in the 1 - 3 pM.range which were
made using a 4 cm. quartz cell.
Preparation of reagents is summarized on pages 28-32.
Sulfide
Hydrogen sulfide was determined using the spectrophotometric methylene blue method (Cline, 1969) with slight modifications.
N,N-dimethyl-p-phenylenediamine-sulfate
(Eastman
Kodak No. 1333) and ferric chloride are reacted in 6 N.
hydrochloric acid to produce the reagent.
The sample is added to the reagent.
The procedure is
optimized for a particular concentration range by changing
the reagent concentrations and sample dilution (Table 2.3).
The presence of sulfide is indicated by an intense blue
color, the absorbance of which is measured at 670 nm.
(Figure 2.1).
.~-~
1
_r
^__
II~LL.*II -i.l.i-l.-~L
--rrXII.--Y;I.~.Cr^-.I
r^arn
irtp-.-x-l
;I--*
irrr~ l----lirr~.----.. ..1_ - ~~.~__1___1_~,~.^-~-----r.-.~.---
-18-
Figure 2.1
Sulfide Analytical Procedure
-illi_.
___rlll____~__L_*____3
~-
~W--
.
-L~IIPII~YII~1~ IPII*Ui- I~II~PY~
-xr--~-r;-~----~
^--L"i~-i-i----
-
-19Figure 2.1
Ssample
1 -
I
3
p M.
H2 S
3 -
40 VM. H 2 S
40 -
I
250 lM. H 2 S
add 5 ml.
sample to
0.4 ml.
reagent 1-3
add 5 ml.
sample to
0.4 ml.
reagent 3-40
add 0.4 ml.
sample to
0.4 ml.
reagent 40-250
wait 20 min.
wait 20 min.
wait 20 min.
measure Abs.
at 670 nm.
in 4 cm. cell
-
~----- - -~~
--~1' -"111----1-1-;"'Xp-~~~-"l~^~--~--
-----^-Ir-l-~-r^---~rr~li-1L1--^X-XIXL*~
--I
s(-~------ IIYI
li--~ _I-~1L-CI~I~--
-20Table 2.3
Sulfide conc. Diamine conc.
Umoles/liter
g./500 ml.
Ferric conc.
g./500 ml.
Dil. factor
ml. : ml.
path
length
(cm.)
1 -
3
0.5
0.75
1 : 1
4
3 -
40
2.0
3.0
1 : 1
1
250
8.0
12.0
2 : 25
1
40 -
For sulfide concentrations expected to border on two
ranges (e.g. 40 pM. sulfide), samples were treated by both
procedures.
The concentration of the sample was then
determined by comparing it with the standard curve covering
the correct range.
The standards were prepared from sodium sulfide using
distilled water that had been freshly boiled, cooled, and
then bubbled with nitrogen gas to remove dissolved oxygen
and carbon dioxide.
The sulfide crystals were washed to
remove oxidation products, dried, and dissolved in the water
to give an approximately 0.1 N. concentration of sulfide.
This solution was then standardized iodometrically (Budd &
Bewick, 1952).
Samples
were
compared to standard
curves fitted by linear regression for the concentration
ranges involved.
(Figure 2.2).
The detection limit was 1 pM. sulfide
The precision at the 95% level of confidence
(Table 2.4) was constant over each of the concentration
ranges used.
--
Ilp-l-1 l~-^.
~iU~-L -IL-.^~l^*sryY-ttl~y----------YL
rll_
-21-
Figure 2.2
0-3 pM. Sulfide versus Absorbance
InIIUI~-LIIX
4
h
-22-
3 "e
2
IL.
@0
1
1
0 -a
20
40
60
80
a
-L - -
100
120
ABSORBANCE
140
1
1
160
180
1-
200
220
llll- 1IL Lli L*IYIIII~YLYYrYIPIII~~YYPI*IIPLI-_I~ i-ili~ -~~
~-..Y_.C* ill-~_
(.^. - -- ~- I~- 1_1
LI L----l..
~-_ILI
.~1II.LIBYI
-_j .~i-- ~I.^_II-*-1-I-~(IF-l~
-23Table 2.4
Sulfide
conc. Umol-/liter
Precision
mol/liter
1 - 3
+ 0.05
3 - 40
+ 0.9
+ 29
40 - 250
Sulfate
Sulfate measurements were performed by Russ McDuff
using the polarographic method of G.W. Luther and A.L.
Meyerson (1975).
The standard error was + 0.2% at the 95%
confidence level (A2).
Iodine Reactive Species
The iodine reactive species present were determined
iodometrically by adding an excess of approximately 0.1 N.
iodine and back titrating the excess iodine with 0.102 N.
thiosulfate
(Vogel, 1961).
thiosulfate daily.
The iodine was standardized with
A seawater blank (approximately 46
equiv/liter of iodine) was subtracted from the samples.
standard error was approximately + 65
95% confidence level.
The
eq/l of 12 at the
This large error could have been
reduced by using less concentrated solutions of iodine and
thiosulfate for the titration.
Sulfite
The colorimetric method of Scaringelli, et al., 1967,
with modifications was used for sulfite determinations (Fig 2.3).
--l_~nui~-~--c_---_---rr----~i-Y~-i--~-^.I~
i--_~_-_--_-----r
I
~)~LIII^~LIII-lypm_-o~xr-_~r~-ri
-laa^-- ------r~L_.^
..ll._L-.----..
...
-24-
Figure 2.3
Sulfite Analytical Procedure
~
~~P~-a~~.~d~- -- -I---r-~r,-
. I~ 1
-l
add
-25Figure 2.3
add to
1 ml. HgC1
2
wait 10 min.
wait 30 min.
Y-ill~X--IY-~~
~
~ pCII11.11I-~--.-.-.-I.-l- ..-_--.----
-26Sulfite reacts with formaldehyde and pararosaniline to form
an intense violet color in acid solution.
A 0.2 M. mercuric
chloride solution (stabilizes sulfite from air oxidation by
formation of dichlorosulfitomercurate) was used instead of
0.04 M. HgCl 4 because of the high chloride concentration in
seawater.
practical
The HgCl 2 solution was made as concentrated as
(0.2 M.)
to maximize the sensitivity of the method.
Millipore (0.45 pm.)
filters were used for the filtrations.
Sulfite standards were prepared with sodium sulfite in
distilled water treated as for the sulfide method.
solutions were standardized iodometrically.
The
The standard
error was + 1 pM. at the 95% confidence level.
The absor-
bances of the samples were then compared to the standard
curves which had been fitted by linear regression.
The
detection limit was 1 1M. sulfite.
Thiosulfate, Trithionate, & Tetrathionate
The methods of T. Mizoguchi and T. Okabe (1975) were
followed for thiosulfate, trithionate, and tetrathionate
measurements.
The three procedures are based on the reac-
tions of these species with cyanide under different conditions (Table 2.5).
The first procedure (A) is cupric ion-
catalyzed cyanolysis of thiosulfate at a pH of 4.5.
The
second method (B) is the cyanolysis of tetrathionate at high
pH, in the presence of acetone followed by cupric ion-catalyzed cyanolysis of thiosulfate.
The last method (C)
h
5
-27Table 2.5
Procedure
Reaction
2-
+ CN
-Cu 2
-
2SCN +SO 2-
(1)
S 0
(2)
-Cu
S O 2- + 2CN + 20H-
2 3
Equivalents of SCN
expected (moles)
3
2+
2SCN + SO
2-
SA=
(S2 0 3
S B=
(S2 0 3
22(S
+ SO 4 2-+ H20
(1)
S203
S23
2.
(3) S3062
(1)
+ CN
-Cu +
SCN
+ SO
23
2-
)+
B 20
)
4 6
2-
2-
+ CN + 20H = SCN + SO 3 2- +SO
2+
-3
)
3
-Cu
+ 2CN
4 6- + 20H
-3 2SCN +SO
2+
2-Cu
2S O
+ CN
SCN + SO
(2) S06
2-
2-
2+
2
2-
2-
4 2- +
2-
+SO 4
2-
2-
Sc
=
H20
Sc =
+ H 0
2
(203 2-)+
(S 3 0 6
22) + 2(S4062
3
6
4 6
l~li
.~~-ill__/l__I___I__l II~C^^i~_l_ _PI_____~_1_~
_
I~.._^
~.-C. .-.rl~i~-X-i~
i._i~I-I..
-28-
involves the cyanolysis of trithionate and tetrathionate in
a boiling water bath at high pH, followed by cupric ioncatalyzed cyanolysis of thiosulfate.
Cadmium acetate was added to the samples (Figure 2.4)
to eliminate interferences from sulfide and sulfite (B.
Sorbo, 1957 and P.J. Urban, 1961).
Pre combusted approxi-
mately 1 pm. glass fiber filters were used to remove the
precipitate.
The samples, where noted, were maintained at
190 C. in a thermostatically controlled water bath.
The
absorbances were measured at 460 nm. in a water cooled (190C)
1 cm. quartz cell.
Cupric sulfate was substituted for cupric
chloride in the above reactions (A, B, & C) when the cupric
chloride solution had been used up on board ship.
Standard curves for method A were prepared using a
previously standardized 0.102 N. sodium thiosulfate solution
(Vogel, 1961).
The solution was made up following the sul-
fide procedure with Na2S203 and standardized with potassium
iodate.
The error obtained from replicate standard curves
fitted linearly was + 15 pmoles/liter at the 95% confidence
level.
Thiocyanate standards that had been previously compared
to thiosulfate standards were used for methods B and C.
The
standards plotted in a parabolic curve (Figure 2.5), with a
minimum of absorbance around 15 pmoles/liter of thiocyanate.
Substitution of a less concentrated solution of cupric
x -.rlilP1-1 -----il---1-I-*LIPII~------LIX~
-29-
Figure 2.4
Thiosulfate, Tetrathionate, and
Trithionate Analytical Procedure
*iill
.~
_.._
.~L_ ~........
_~._
CLX
r~Y*III1IIYllllj^_~
------YIU1III~-I I~C--III. -- .I~LI
-30Figure 2.4
150 il.
acetate
I
10 ml. sample
to 19 0 C
cool
III
20 ml. sample
add 3 ml.
acetate buffer
pH adjusted
to 9.3 w/
IM. ammonia water
ad -3ml
mix
add 0.2 ml.
KCN
mix
I
I
B
10 ml.
10 ml.
add 3 ml.
acetone
add
0.4 ml. KCN
cool
to 190C
mix, boil
for 30 min.,
cool to 190 C
0.4 ml. KCN
added
mix & wait
20 min.
add 0.3 ml.
CuC12 or CuSO 4
mix
add 0.3 ml.
CuC12 or CuSO
4
mix
add 1.5 ml.
Fe (NO3 ) 3
mix
measure
Abs. at 460 nm.
in water cooled
(190C) cell
~..... --r-...r
;-;-- lr- r~l^l-ixc~
-ii-r=iI-* u~-iurax~--^.rs~
--
'-i'p~--i- u--- -Ir~;icsl
uurPnlXL+slll---I~~'LI-~i~
-31-
Figure 2.5
Thiocyanate, versus Absorbance,
Method C, (CuSO4 )
4
50
l 40
S
@0@
MII
I-
30
20
-
S
S.
0
10
550
ABSORBANCE
600
6
4
-33sulfate for the cupric chloride in the procedure increased
the standard error (Table 2.6).
The parabolic shape of the
standard curves in these procedures (B and C) may be due to
an interference caused by the seawater medium since this
phenomena is not observed in distilled water.
In general, errors on all the below procedures might
have been further minimized if initial manipulation of the
samples had been carried out under a nitrogen atmosphere.
This, however, was not possible on board ship due to a
shortage of nitrogen.
Table 2.6
Std. error at 95% conf.
(2a) pM. thiocyanate
(CuCl2 )
+
"
(CuSO4 )
+ 10
Method C
(CuCl2 )
+ 4.5
"
(CuSO4 )
+ 4.7
Method B
7
___=~I~ I_
1 _-I-_YI._U--~L---IC~II^ -~---
~_~_I~_~ __ ___.__II~YIX111~11_1C-^~
.-.
-34Table 2.7
Summary of Method Std. Errors
Std. error at 95%
confidence iM./liter
Method
Sulfide (1-3
f"
_L4)
+ 0.05
(3-40 '_4.)
+ 0.9
(40-250 PM
+ 29
Sulfate
+ 0.2%
Iodine Reactive Species
+
Sulfite
+ 1
S203(A)
65
+ 15
SCN (B, CuC1 2 )
+
7
SCN (B, CuSO 4 )
+ 10
SCN (C, CuCl 2 )
+ 4.5
SCN (C, CuSO 4 )
+ 4.7
--.
I~
-I--IIII~L_-EI--II.
^^__
Il---l~b-.I I-I~-~~C-I~XIII)_.-.__.
-35REAGENTS
Sulfide:
N,N- Dimethyl-p-phenylene-diamine Sulfate;
[NH 2 C 6 H 4 N(CH 3 )2 2 H 2 SO 4 , Eastman Kodak Co.,
No. 1333.
Ferric Chloride; FeCl 3
.
6H20, A.C.S. Reagent
grade, Matheson Coleman & Bell.
6N. Hydrochloric Acid; A.C.S. Reagent grade, Fisher
Scientific Co..
Sodium Sulfide; Na2S-9H20, A.C.S. Reagent
grade, Mallinckrodt, Inc..
Iodine Reactive Species:
0.1 N. Iodine; A.C.S. Reagent grade, J.T.
Baker Chem. Co.,
(standardize daily).
Starch soln. - Make a paste of 1.0g. of Soluble
Starch; A.C.S. Reagent grade, Merck and Co., Inc.,
with a little distilled water, and pour the paste
with constant stirring, into 100 ml. of boiling
distilled water, and boil for one minute (A.I.
Vogel, 1961).
_;r.^-~-- .~-^--XI^-~xl--~nr-^-_m.r.rir.nns--.X.1DLX*-~
~L -I~-*~-I~I~L~--~-IIPXI~PU_yCI
r^l--I-Y---IL---
-360.1 N. Sodium Thiosulfate; Na 2 S2 0 3 .5H2 0 )
A.C.S. Reagent grade, Mallinckrodt, Inc..
Potassium Iodate; KIO 3 , A.C.S. Reagent grade,
Allied Chem., Specialty Chemicals Division.
Potassium Iodide; KI, A.C.S. Reagent grade,
J.T. Baker Chem. Co..
2N. Sulfuric Acid, H2S04, A.C.S. Reagent grade,
Mallinckrodt, Inc..
Sulfite:
3M. Phosphoric Acid; H 3PO 4 , A.C.S. Reagent
grade, 85%, Fisher Scientific Co..
0.2 M. Mercuric Chloride; HgC1 2 , A.C.S. Reagent
grade, Mallinckrodt, Inc..
0.6% Sulfamic Acid; NH 2 SO 2OH, Assay 99.90-100.10%,
G. Frederick Smith Chemical Co (Prepare daily).
0.2% Formaldehyde; HCHO, A.C.S. Reagent grade,
Approximately 37%, Mallinckrodt, Inc.,
(Prepare
daily).
IM. Sodium Acetate-Acetic Acid buffer; Na2C2H302
3H20, A.C.S. Reagent grade, Matheson Coleman and
Bell.
Acetic Acid Glacial; CH 3 COOH, A.C.S. Reagent
Grade, 99.7%, Fisher Scientific Co..
..(-1~-~YII~____~_.-^
~-~I~III-^.~L-I._1II-C--ULII~..i~l-~il..~~~~ 1I
l~l~
I~-~~-IIICII
^I~s~ll~*-_I~1L-----I--^11111
-.. ^__i-1L.I
- ilY--~~
-----
-37Pararosaniline Chloride; C 9H18ClN 3 ,
(0.2% + 0.03 in 1 M. hydrochloric acid), Eastman
Kodak Co., A 14051.
P.C. concentration (Figure 2.6)
Pararosaniline Chloride (P.C.),
0.2% stock solution
1 ml.
Dilute to 100 ml. with
distilled water
5 ml.
Add 5 ml. acetateacetic acid buffer.
Dilute to 50 ml.
with distilled HO
Wait 1 hour, measure
absorbance at 540 nm.
r I~-PPsl~LI
--
IYL1~-.~
--- r~l~-lx.
l-.-.-I.~-. I -nr--~x---- ----
UC- --U~L~~
-~-Ili~~ll~CllTI~ 11I.1~~~
*~L----~~iL-ellrii3-_.__~-~~L.rr.l~..^-----.--~ *--~---lyn~l
-38The actual pararosaline concentration as a percentage of
the nominal concentration is determined by the formula:
(1) % P.C. = Abs x k
grams of dye taken
to allow for correction of P.C.
(lA) % P.C.
= (.458)
, where k = 21.3
in reagent.
(21.3)
= 97.5% P.C.
(.1)
Procedure B (Scaringelli, et
al., 1967) was used
because this method covered a wider range of sulfite concentrations.
PRA reagent for method B is prepared by adding
200 ml. of 3 M. phosphoric acid to 20.5 ml. stock P.C. and
diluting to 250 ml. with distilled water (PRA reagent).
An
additional 0.2 ml. of stock P.C. was added for each one
percent that the stock P.C. assays below 100%
(0.2 ml. x
2.5% = 0.5 ml.).
Sodium Sulfite Anhydrous; Na 2 SO 3, A.C.S. Reagent
grade, Merck and Co., Inc..
Thiosulfate, Trithionate, and Tetrathionate:
0.75 M. Potassium Cyanide; KCN, A.C.S. Reagent
grade, Matheson
Coleman and Bell.
0.3 M. Cupric Chloride; CuCl2*2H20 , Analytical
Reagent grade, Mallinckrodt, Inc.
OR 0.2 M. Cupric
Sulfate; CuSO 4 .5H 2 0, A.C.S. Reagent grade, J.T.
Baker Chem. Co..
/~___ X^__I~____PY___U_____LI~_~\
I41~-rllYIIIII~---L1
e
-390.2 M. Cadmium Acetate;
(CH3 CO 2 )2 Cd2H2O,
Reagent grade, Matheson
Coleman and Bell.
2 M. Sodium Acetate - Acetic Acid buffer;
(see
above).
1M.Ammonium Hydroxide; A.C.S. Reagent grade, Assay
28.0-30.0% NH 3 , Fisher Scientific Co.
1.5 M. Ferric Nitrate - Perchloric Acid Reagent303 g. of Ferric Nitrate; Fe(NO 3 )3 9H2 0, A.C.S.
Reagent grade, Mallinckrodt, Inc. was dissolved in
a small volume of distilled water containing 186 ml.
of conc. Perchloric Acid; HC104, A.C.S. Reagent
grade, 70 wt.%, G. Frederick Smith Chemical Co., and
diluted to 500 ml. with distilled water.
0.1 N. Thiosulfate;
(see above).
Potassium Thiocyanate; KCNS, A.C.S. Reagent grade,
Fisher Scientific Co..
^I-iipl~l----l--l-li~i~.
.
-~~111~1
111_
-40-
CHAPTER 3
RESULTS AND DISCUSSION
~-~-_L
.~IP(-.Y-X-~FIP..XI-~-I^--I--I----Yl~ L~.~
n--I~^-^*PII-~--YI~-~
-C~~~---III--IWI
-41Sulfur species in the Galapagos hydrothermal waters
show linear trends versus temperature in the data (passing
through the ambient sea water temperature and composition)
which were interpreted as dilution lines.
Dilution of the
concentration of the species was due to a combination of
mixing below the seafloor, in the vents, and during sampling
with ambient sea water (J. B. Corliss, et al., 1979 and
J. M. Edmond, et al., 1979a,b).
Problems with the sampling
system leaking led to further dilution (0-89% hydrothermal
water collected) and possible oxidation of the reduced sulfur
species present in the samples.
Sulfide concentrations in the vents increased with
increasing temperature (Table 3.1, Figure 3.1).
East of
Eden had the highest gradient (2.6 x 1012 mol/yr) with
Mussel Beds changing the least (Table 3.2).
These trends
indicated that hydrogen sulfide was being produced in the
hydrothermal system, either by reduction of seawater sulfate,
leaching sulfide from the basalt or primary hydrogen sulfide
from magmatic sources (Bonatti, 1975).
Sulfide concentration
versus oxygen concentration (Figure 3.2)showed a decreasing
trend.
Oxygen was absent in the samples with hydrogen
sulfide concentrations above 120 pM, which differs substantially from data previously published
(Edmond, 1979b)
where oxygen disappeared at approximately 50
M sulfide.
This difference is due to the entrainment of oxygenated
__\_YIY__^C___1__L_)- li--~i~~-I.I-._Y1I_
I-r~Y
LIL~-X -..~~II -~Y~-I-I*LL.
-42-
Table 3.1
Concentrations of Sulfide, Sulfite,
Thiosulfate, Trithionate, Tetrathionate,
Iodine Reactive Species, Oxygen,
and Silica for the Galapagos Vents
-43Table 3.1
Dive #
Vent #I
pM.
H 2!S
iP1.
S203
PA -
Si
02
Bottle #2
1M.
2-
SO 3 2-
PM.
12
peq.
SOM4
VM.
2-
2-
s 02-
S 06
A.
898
1
207
275
84
1.4
1
-1
336
11
-5
898
1
204
259
114
1.6
9
-1
312
9
-5
898
1
103
402
48
10.5
2
0
581
898
1
202
400
60
16.7
2
1
589
898
1
203
457
36
19.7
8
1
581
898
1
206
201
106
7
-1
899
2
105
619
0 147.0
5
899
2
106
596
0 144.0
899
2
107
587
0 165.0
899
2
101
513
4
899
2
109
550
0 145.0
899
2
113
479
0
899
2
112
395
899
2
102
303
900
2
202
854
900
2
206
846
12
1
12
-3
12
-8
450
10
-3
1
216
12
-4
11
1
277
6
7
18
1
314
5
-6
9
0
335
27
0
220
1
67.0
8
1
196
10
-6
19
44.0
7
1
167
11
-6
58
49.0
6
1
151
11
2
0 177.0
14
1
453
7
-4
0 285.0
18
1
510
0.3
75.0
28.16
28.24
27.88
27.42
9
0
4g
Cr~
-44Table 3.1
Dive #
Vent #
Si
02
H S
Bottle #2
PM.
PM.
PA.
(contId.)
2-
SO 3 2,_.3
12 3
peq.
s 0622--
2-.
PM.
8
-5
8
-7
379
7
-7
3
188
9
-7
3
183
10
-6
4
198
369
900
2
103
625
0
190.0
14
0
900
2
207
348
34
93.0
14
2
900
2
204
608
0
162.0
17
0
900
2
203
322
54
98.0
13
900
2
205
279
66
80,0
11
901
3
109
161 117
0.5
901
3
102
161 119
0.4
288
901
3
107
160
115
0,9
288
901
3
112
167 109
0,7
300
901
3
106
336
0
142.0
3
316
901
3
105
360
0
182.0
-7
406
902
1
204
605
0
31.5
15
77
902
1
203
626
0
35.8
21
98
902
1
104
490
43
21,9
902
1
207
575
0
31.5
73
13
28.30
27.78
18
-10
8
-7
5
2
12
-6
6
0
-45-
Table 3.1
(cont'd.)
Si
02
H2S
S 20 3
SO3
PM.
pM.
3
12
peq.
2-
SO 4 2-s06
PM.
pA.
S 06
2-
uM.
Vent #
Bottle #2
PM.
PM.
PM.
902
1
202
486
13
22.4
29
0
-20
1
902
1
205
486
10
19.1
14
0
82
8
902
1
103
443
55
12.9
0
1
16
15
903
2
102
558
0 137.0
11
1
143
10
2
903
2
107
837
0 140.0
5
1
384
9
5
903
2
106
927
0 221.0
4
1
633
13
-5
903
2
101
636
0 138.0
5
1
45
12
-4
903
2
109
724
0 169.0
5
1
294
903
2
302
420
26 118.0
6
1
904
1
205
588
0
28.1
28
-1
132
904
1
207
561
3
24.5
28
0
120
904
1
204
507
15
20.7
25
0
904
1
202
595
0
33.0
25
0
904
1
303
574
11
28.6
27
905
2
113
1007
0 336.0
18
Dive #
27.47
27.78
-1
2
-6
12
2
11
0
0
-1
0
0
28.02
2
0
140
27.84
2
-2
0
112
27.87
1
-1
1
446
27.14
4
-1
62
59
27.81
-46-
Si
Table 3.1 (cont'd.)
22HS
0
S03
S yM. 3
PM.
PA. ,,A.
3
12
-peq.
Bottle #2
PM.
905
106
238
33
31.9
15
905
112
372
28
70.0
15
905
101
627
0 190.0
15
905
109
596
0 148.0
15
1
184
905
304
343
317
7
1
-11
165
120
Dive #
Vent #
0
I.:
Mussel Beds; 2:
1.
Vent #:
2.
Bottle #:
3.
Iodine Reactive Species
1 : A; 2 : B; 3 : H;
2-
2-
s
2S0
60
-5
233
0
0.0
Rose Garden; 3:
SOM.
VM.
East of Eden
(e.g. 105 --A 5).
27.64
5
1
-3
28.62
0
-I~LUIEPr~s~-
-47-
Table 3.2
Gradients and Fluxes for Sulfide
and Sulfate
-~r~-rULll^_nyulr~
-48Table 3.2
SO
Range
0.0-336
moles )
Gradient (Cal.
Cal.
(1)
5.14 x 10
(2)
1.83 x 10
(3)
5.24 x 10-8
5.24 x i0
Flux
moles
)
.
M.
11
(1) 2.6 x 1011
(2) 9.2 x 10
2-
27.14-28.62 mM.
Iodine
Reactive
Species
-20-633
(1)
-1.06 x 10-7
-1.06 x i0
- 12
-5.3 x 1012
(3) 2.6 x 1012
peq.
-8
6.08 x 10
(1A) 3.27 x 10-8
(2)
3.65 x 10
(3)
3.13 x 10-8
12
(1) 3.04 x 1012
(LA)1.63 x 10
(2) 1.83 x 1012
(3) 1.56 x 1012
For the above calculations, 5 x 1019 cal/yr. was taken as the global hydrothermal heat transport from accreting plate boundaries.
(1) Mussel Beds
(2) Rose Garden
(3) East of Eden
,
(lA) Mussel Beds, Dive 898
_I--_---
I
_ ..I~IIIIIIYYCI~-
-49-
Symbols Used in Figures in Chap. 3
O
- Mussel Beds
+
- East of Eden
A
- Rose Garden
_
-PLIC-I^~TYI~IP~I~I~-Y)
1.I~1I~--_1(
11III~-~.I^L----Llr*rillllX- ~Li_--il__i(
- ...~.i_^ll.
.---IPI_~- . .
-50-
Figure 3.1
Sulfide versus Silica
S-51-
,wI
a
*6*
M
In
+
UA
NI
100
iA
Oo (")O
AA
w
200
A
A
A
$00
%V
300
400
SILICA
pM.
900
1000 1100
I~-*~--.
lil__~llll*____1_____~
_IPI_; 1)~
-52-
Figure 3.2
Sulfide versus Oxygen
-l~--si-l. .~-1X~
-53-
350
300~
300
A
2501
200
u.u
Q.
,=6
150
100
50
0 c
o
|
A
I
10
20
OXYGEN
pme
80
Ip
I
%
90 100 o110o120
_ I^lr~_i__l~____ I ~ll----- - IIIICIII1*
~
LYIIII--PIXI~I-~~--rl
-54ambient water during sampling with the leaky samplers and
the relatively slow oxidation of sulfide.
(Sulfide
oxidation is inhibited at pH values below 7, presumably
because of the protonation of the HS
molecule, Chen and
Morris, 1972).
Sulfate decreased uniformly in all vent fields
(Figure 3.3, Tables 3.1, 3.2).
Sulfite was constant with temperature at about 1 pM
(the detection limit),
(Figure 3.4), though in East of
Eden and Rose Garden some of the concentrations at low
temperature were above 1
M sulfite.
Thiosulfate, trithionate, and tetrathionate were also
constant with temperature within the limit of resolution
(Figure 3.5, Table 3.1).
Cyanolysis of any elemental
sulfur present (only particles <1 pm as the samples were
filtered) should interfere positively to some extent in
the tetrathionate method and to a greater degree in the
trithionate procedure (Nelson, et al.,
1977), though this
interference has not been quantified.
Species reduced by iodine showed a general increase
with temperature (Figure 3.6).
The vent fields differed
in their gradients (Table 3.2) with Mussel Beds exhibiting
two different trends associated with two different days
of sampling.
In water with a neutral or acidic pH
(hydrothermal waters--pH 6 to 7) only sulfide, sulfite,
s
_I_ ~_I~
I~--~------------"III~arr~--Ll~iPIIY IZL
-55-
Figure 3.3
Sulfate versus Silica
i-'-ivl'lir~
-x~--~1li1--~
-
I~I
~I--- --
-i-(r~_
-56-
29
E
I-
28
-U
27
r
r
I
200 300 400
I
SILICA
I
I
pM.
800
900 1000 1100
1.
~-T
l)i_ ICI-X----.I.-
-57-
Figure 3.4
Sulfite versus Silica
-58-
4
S
U
L
F
I
AA
A
E
iPMo
&A
Ab
A.dM
'AI
i
tj
Ar
J
1
I
v
- MM
A&VV'N
m
T
v
l
i
-
-
I
-1
-2L
100
SIL.ICA pM.
200
300 400
500
600
700 800 900 1000 1100
____l*l__lj~ll*___C__IP_-C.~
-59-
Figure 3.5
Thiosulfate versus Silica
~--I---X~-L
- LLI_(-LYL_.
- liUCI Ill_~l
-~X
L~l L-LI1
-60-
30
20
AA
10
aL
I-
OAA
A
I,
:)
tn
U6
2
zI-
A
AA
0
-10
I-.I
-
:
I
200
o
.
. "/
•
300
400
V
I
!
!
•
500
600
SILICA pM.
700
I
800 900 1000
1----L1-~
IPIPrrPP~-~1~~
-rilr)~a3
~lrr.~--~
Il~-------~_izr
~yll...
--~r.~
p~__..
-61-
Figure 3.6
Iodine Reactive Species versus Silica
4
4
0
46
-62-
U)
600
400
0
0ILl
w11
400
U
z~d~
2k
200
m-
300
A
A
*
+
k
A A
A %
100
-100
A
200
300 '400
500
SILICA
600
700 800
pM.
900 1000 1100
I~-~L--^II
L;~ ___ZILII^__*XIll___-*II~
-63and thiosulfate should be expected to be reactive with
iodine.
The concentrations of these sulfur species leaves
a large amount of reacted iodine unaccounted for (Figure 3.7).
A possible explanation for this is suggested by experiments
by Nelson, et al.
(1977); colloidal sulfur suspensions
stabilized by thiosulfate or polythionates may lead to
iodine reactivity, unlike other sulfur suspensions.
This
reactivity may be due to a surface charge mechanism
resulting from slow hydrolysis or oxidation upon the
surface of the sulfur crystals, whereby charge sulfoxy
groups react.
The hypothesis that these unaccounted for
iodine reactive species are in fact colloidal sulfur is
further supported by the observation that the major
component of suspended material filtered in situ from
the hot springs is native sulfur (Corliss, unpublished
data).
Polysulfides, while also iodine reactive, were not
measured since their presence is unlikely due to the
low pH conditions present in the hydrothermal solution.
In a profile of the overlying waters
135 and 139),
(stations 86,
reduced sulfur species were not found.
The sulfur chemistry of the Galapagos hydrothermal
vents suggest that hydrogen sulfide and possibly elemental
sulfur are the principal reduced sulfur species.
thiosulfate, and polythionates are observed in low
Sulfite,
~_rsr~-iYr~
-- a-~-r~----r*li-----L
_r~.~....~l.itn~l^ir-~-rr~-
-64-
Figure 3.7
Iodine Reactive Species Minus
Sulfide, 2x Thiosulfate, and
Sulfite (12 ) versus Silica
-65-
500
400-
300
0)
+
AA
2001+
C4
A
A
100-
A
A
#9
4
SA
* W
A * A*P
AQ
A9A-
- 100oo
100
200
300
400
500
SILICA
600 700
pM.
800
900
1000 1100
~___j
LI_l_____n_
I__IC_____J_~_I___~IUL___I_
-66concentrations implying that sulfur is not oxidized beyond
the zero valence state to any significant extent.
The flux
of sulfide from the hydrothermal system is large, 2.6 x 1011
2.6 x 1012 mol/yr, which is consistent with previous observations (Edmond, et al., 1979b).
Sulfate is reduced in the vents (flux = -5.3 x 1012
mol/yr) suggesting that this is an important sink in the
oceanic sulfate budget (Edmond, et al., 1979a).
The observed concentrations of sulfur species in the
hydrothermal vent waters suggest that elemental sulfur
and possibly polysulfides should be examined in more detail
as the measurement of these species will give additional
insight into the sulfur chemistry of seafloor hydrothermal
systems.
-67-
References
--lr r^r --r~--rrP-Nr~ilX1*lyr~ IQIL~LI~LICli
LI1^Y---~
~
-68-
References
Bonatti, E. (1975).
Centers.
Metallogenesis at Oceanic Spreading
Annual Review of Earth and Planetary
Sciences,3, 401-431.
Brewer, P.G. (1975).
Minor Elements in Seawater.
In
J.P. Riley and G. Skirrow (eds.), Chemical
Oceanography, (Vol. 1, 2nd ed.), Academic Press,
London, 415-496.
Budd, M.S. and H.A. Bewick (1952).
Photometric Determination
of Sulfide and Reducible Sulfur in Alkalies.
Analytical Chemistry, 24, 1536-1540.
Chen, K.Y. and J.C. Morris (1972).
Kinetics of Oxidation
of Aqueous Sulfide by 02.
Environmental Science
and Technology, 6, 529-537.
Cline, J.D. (1969).
Spectrophotometric Determination of
Hydrogen Sulfide in Natural Waters.
Limnology and
Oceanography, 14, 454-458.
Cline, J.D. and F.A. Richards (1969).
Oxygenation of
Hydrogen Sulfide in Seawater at Constant Salinity,
Temperature and pH.
Environmental Science and
Technology, 3, 838-843.
Cline, J.D. and F.A. Richards (1972).
Oxygen Deficient
Conditions and Nitrate Reduction in the Eastern
Tropical North Pacific Ocean.
Oceanography, 17, 885-900.
Limnology and
~L~---~.
IY*lle;-~~-IVI
______/_LI~_L___PIYYIC_
-69Corliss, J.B., J. Dymond, L.I. Gordon, J.M. Edmond, R.P.
von Herzen, R.D. Ballard, K. Green, D. Williams,
A. Bainbridge, K. Crane, and T.H. van Andel (1979).
Submarine Thermal Springs on the Galapagos Rift.
Science, 203, 1073-1083.
Edmond, J.M., C. Measures, R.E. McDuff, L.H. Chan, B. Grant,
L.I. Gordon, and J.B. Corliss (1979a).
Ridge-
crest Hydrothermal Activity and the Balances of
the Major and Minor Elements in the Ocean:
Galapagos Data.
the
Earth and Planetary Science
Letters, (in press).
Edmond, J.M., C. Measures, B. Mangum,
B. Grant, F.R.
Sclater, R. Collier, A. Hudson, L.I. Gordon, and
J.B. Corliss (1979b).
On the Formation of MetalEarth and Plane-
Rich Deposits at Ridge Crests.
tary Science Letters, (in press).
Goldhaber, M.B. and I.R. Kaplan (1974).
The Sulfur Cycle.
In E.D. Goldberg (ed.), The Sea, (Vol. 5)
John Wiley and Sons, New York, 569-655.
Humphris, S.E. and G. Thompson (1978).
Hydrothermal
Alteration of Oceanic Basalts by Seawater.
Geochimica et Cosmochimica, 42, 107-125.
Luther, G.W. and A.L. Meyerson (1975).
Polarographic Analy-
sis of Sulfate Ion in Seawater.
Chemistry, 47, 2058-2059.
Analytical
LI
LII
_IIC______Ll_____lr___III1I_--L-
-70Mizoguchi, T. and T. Okabe (1975).
The Chemical Behavior
of Low Balance Sulfur Compounds.
IX.
Photometric
Determination of Thiosulfate, Trithionate and
Tetrathionate in Mixture.
Bulletin of the
Chemical Society of Japan, 48, 1799-1805.
Mottl, M.J., H.D. Holland, and R.F. Corr (1979).
Chemical
Exchange During Hydrothermal Alteration of Basalt
by Seawater-II.
Experimental Results for Fe, Mn,
and Sulfur Species.
Geochimica et Cosmochimica
Acta, 43, 869-884.
Nelson, M.B., J.A. Davis, III, M.M. Benjamin, and J.O.
Leckie (1977).
The Role of Iron Sulfides in
Controlling Trace Heavy Metals in Anaerobic
Sediments:
Oxidative Dissolution of Ferrous
Monosulfides and the Behavior of Associated Trace
Metals.
Sorbo, B. (1957).
NTIS Report AD-A049712.
A Colorimetric Method for the Determina-
tion of Thiosulfate.
Biochimica et Biophysica
Acta, 23, 412-416.
Urban, P.J. (1961).
Colorimetry of Sulfur Anions, Part 1,
An Improved Colorimetric Method for the Determination of Thiosulfate.
Chemie, 179, 415-422.
Zeitschrift fuer Analytische
_
i*.~-yllllll^~--II*--
1~-1sl~
11(L11~
--
~-1-1-_1111
-----~C.~- 1-LI~Cl~-.
11~11~3
1L^.-1__I-L
9-I~LIIY.-P^I~-l^
---ll~ ~
_I~-1
__i
-71Vogel, A.I. (1961).
A Textbook of Quantitative Inorganic
Analysis Including Elementary Instrumental
Analysis (third ed.).
New York, 1216 pp.
John Wiley and Sons Inc.,