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OCEANOGRAPHY M ADMETEOROLOGY ~.

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CARBONATEEG SEDIMENTS TOEOLUIO

THE AGRICULTURAL AND MECHANICAL COLLEGE OF TEADepartment of Oceanography & Meteorology

College Station, Texas

Texas A. & M. Research Foundation

Office of Naval Research National Science roundationContract Nonr 2 119 (04) Grant 11591A & M Project 286-A A & M Project 249

THE RESISTANCE OF RECENT MARINE CARBONATE SEDIMENTS TO

SOLUTION

Technical Report

Ref. 62-18T

Prepared by

John F. Jansen, Yasushi Kitano, Donald W. Hood & Louis S. Kornicker

ACKNOWLEDGEMENTS

This research was supported by grants from Texas A. &. M Research

Foundation Project 249, sponsored by the National Science Foundation

grant NSF-G1 1591 and Texas A. & M. Research-Foundation Project 286-A,

sponsored by the Office of Naval Research contract Nonr 2119(04), Project

NrO 3-0 3 6.

Suggestions and criticisms of Professor George Kunze and Mr. Edward

R. Ibert are gratefully acknowledged.

This report is based on a thesis submitted by John Frederick Jansen

to the Graduate School of the Agricultural and Mechanical College of

Texas in partial fulfillment of the requirements for the degree of Master

of Science.

TABLE OF CONTENTS

LIST Of FIGURES. vi

LIST Of TABLES..................viii

LIST OF APPENDICES ... . ... .. .. .. ix

CHAPTER

I INTRO-DUCTION.......................Objectives....................Summary of Past Research..

II EXPERIMENTAL.... ................. 3Reagents Used it Experimental Procedures3Equipment 'Used in Experimental Procedures. 3Natural Carbonate Samples Used...........4Solution Experiment..................9

Introduction ............. 9Procedure.....................9

III RESULTS AND DISCUSSION.... ........... 12NaHCO Solution Experiment...............12Distihed Water Solution Experiment........19NaCl Solution Experiment. ............ 23MgCL2 Solution Experiment,...........27NaCi + MgC12 Solution Experiment. ....... 31The Effects of Particle Size and pH. . 36

IV CONCLUSIONS.......................40

BIBLI-OGRAPHY .. .. .. .. .. .. .. .. .. 42

v

LIST OF FIGURES

The Relationship Between Aragonite and MgCO 3in Recent Marine Carbonate Sediment Samples. , 10

Dissolved Ca And Mg++ in 0.00,23 M NaHC0 3with Respect to Percent MgC0 3 in SedimentSamples. .................... 13

Dissolved Ca and Mg++ in Distilled Waterwith Respect to Percent MgCO3 in SedimentSamples............................ . •.. 13

4 Carbonate Alkalinity Resulting from theSolution of Sediment Samples in NaHCO3,Distilled Water, and NaCl with Respect toPercent MgCO 3 in Sediment Samples. ........ 14

5 Carbonate Ion Concentration of the Solution ofSediment Samples with Respect to PercentMgCO 3 in Sediment Samples. .. ......... .. 16

6 Ion Products of the Solution of SedimentSamples in NaHC03 , Distilled Water, and NaCIwith Respect to Percent MgCO in Sediment

3Samples ..... . . . ........... ..... . 17

7 Dissolved Ca++ and Mg+ in 0..405 M NaCl withRespect to Percent MgCO 3 in Sediment Samples 24

8 Dissolved Ca+ + in MgCl 2 and NaCl + MgCl 2with Respect to Percent MgCO 3 in SedimentSamples

And

Carbonate Alkalinity Resulting from the Solutionof Sediment Samples in MgCI2 and NaCl + M+Cl 2with Respect to Percent MgCO 3 in SedimentSamples .. . .............. .... ....... 28

9 Ion Products of the Solution of Sediment

Samples in MgCl2 and NaC! + MgCl 2 withRespect to Percent MgCO 3 in Sediment Samples . 30

vi

LIST OF PtG~uRES Continued

Fig ur ag

10 Dissolved Am~ounts of Ca+ with Respect toPercent MgCO,3 in Sediment Samp Ilea in theIot Pairs of Sythketic Sea ;Water...........34

1t Dissolved Amounts of Mg+ with Respect toPercent MgCO-3 in, Sedimfeftt Samtples in th e IonPairs OfSynlthetic Sea Water............35

12 Typical Variations in pH for Various CrainSize 'Carbodate Sedimients.............37

vii

LIST OF TABLES

Table Pg

1 Size Analyses. .................. . 5

2 Ca+ anid Mg ++Content of iUnground NaturalCarbonate Sedimnents.....................6

3 Ca++ and Mg+ Content of rounid NaturalCarbonate Sedimhentsi.................. . 7

4 Mineralogical Composition of Ground NaturalCarboniate Sediments........ ........... 8

viii

LIST OF APPENDICES

APPENDIX I

Sec-tion Pag

A Preparation of Calcium Ion Standard Solution .. 44

B Standardization of EDTA for Ca++ Analyss 45

C Standardizationi of Mg-EDTA for Total Ca+ + Mg+Analysis . 46

D The Determination of Ca++ and Mg++ in NaturalCarbonate Sediment Samples.............47

9 The Determination of the chloride Lon in theCmponients of Synthetic Sea Water.~.. 5

F The Determination of Dissolved Amounts of Ca+and Mg++ in Various Solutionse. ......... 5

APPENDIX It

Table

1 Quantities of Calcium and Magnesium Disgsolvedin 0.0023 M NaHC03 . . . . . . . . . . . 57

2 Quantities of Calcium and Magnesium inDistilled Water......................58

3 Quantities of Calcium and Magnesium in0.405 M NaCl . . 59

4 Quantities of Calcium and Magnesium in0.052 M MgC 2. ?... . . . . . ... . . .. . .. .. . . 604

5 Quantities of Calcium and MagneimI0.405 M NaCl + 0.052 14 MgCl 2 . . . .... 61

ix

LIST OF APPENDICES Continued

Tab-1e Pg

6 Carbonate loft Concentration, [Ca+I+ C and(M" (o) Ion Products Resulting 3from~

the Solution of Sedim'ents in MIICO3 atidDistilled Water ..... ................. 62

7 Carbonate Ioft Cotcentratiot, [C+ (COI&. and[Hg++J [CO; ] Ion Products Resulting from theSolution of Sediments in NaCl, MgCl2, andNaC. + MgCI2 4.. .. . . . .. . . .. . . .. . . . .. . . . .. . . .. 63

APPENDIX III

I PH Valucs of Sediments Dissolving in NaHCO03 - 65

2 PH Values of Sediments Dissolving in DistilledWater,...............................66

3 PH Values of Sedim-ents Dissolving in -0.405 M NAC1 . . . .... j. .. ... .... 67

4 pH Values of Sediments ]Dissolving in0.052 M M-gCl 2 .. . . . . . . . . . . .. . . . .. . . . . 6 8

5 PH Values of Sediments Dissolving in0.405 M NaCl + 0.052 H MgCl 2. . . . . . . . . . . . 69

6 pH Values of Sediments Dissolving in Sea Water .70

x

CHAPTER I

INTRODUCTION

Objectives (Statement of the Problem)

The objective of this study is to determine the effect

of the particle size, mineralogy and chemical com position of

recent marine carbonate sediments on their resistance to soluw

tion in the ion pairs of sea water and distilled water.

Summary of Past Research

The polymorphs of calcium carbonate to be investigated

in this research are calcite and aragonite. Seidel! (1940)

showed that the difference in solubility between those poly-

morphs is associated with their difference in crystal form.

Much of the published literature on the solubility of car-

bonates up to the time of Seidell's work is confusing and

contradictory and is based on insufficient data and little

control.

The most stable polymorph of calcium carbonate is calcite.

Calcite is known to be more resistant to solution than arago-

nite; however, coexisting material or impurities in the cfys-

tal structures of calcium carbonate greatly affects its solu-

bility. Chave (1954) and Kitano and Furutsu (1959) have shown

2

that magnesium, the major impurity in marine carbonates,

exists in the calcite crystal structure as a solid solution

Series between calcite and dolomite, and that as the mag-

nesium content increases in the solid solution series, the

solubility of calcite eventually becomes greater than arago

nite. As magnesium exists in the calcite crystal form and

is incompatible with aragonite (Chave, 1954), the crystal

form of calcium carbonate, as well as the availability of

magnesium) exerts a major control on the solution of car-

bonates in nature.

The majority of the published research in the past two

decades, on the stability of carbonates has been carried out

under higher than normal temperature and pCO2 conditions.

An excellent study of the stability of a few carbonates is

presented by Garrels, et al. (1960), near room temperature

conditions. This study was, however, carried out under one

atmosphere pressure of CO with the major emphasis placed2

on changes in pH during equilibration.

Other investigations on the solution of carbonates have

also emphasized pH changes (Holland, et al., 1960; Weyl, 1961).

The effect of magnesium on the solution of carbonates in

natural samples and the effects of solutions other than dis-

tilled water have not been investigated.

C HAPT ER II

EXPERIMENTAL

Reagents Used in 9Expgrimental Procedures.

1) Buffer Solution: Beckman pH 7 buffer solution was

prepared.

2) Synthetic Sea Water: Solutions of the following ion

pairs of sea water were prepared from reagent grade

chemicals.

NaCl(g/l) MgCl 2 (g/1) NaHCO3 (g/1)

Sodium Chloride 23.70

Magnesium Chloride 4i975

Sodium Hydrogen

Carbonate 0.194

Sodium Chloride +Magnesium Chloride 23.70 4.975

Equipment Used in Experimental Procedures.

1) Electrodes: Beckman calomel electrodes and glass

electrodes were used as furnished with the "Zeromatic".

2) Glassware: All glassware used was of the standard Ilaboratory variety as no special apparatus was needed.

Several gross of 250 ml erlenmeyer flasks and at

3

4

least a dozen long stem funnels and 150 ml beakers

were needed.

3) Mortar and Pestle: A motor driven agate mortar and

pestle was used to reduce sediment samples to approxi-

mately 325 mesh.

4) pH Meter: A Beckman Zeromatic pH Meter was employed

to establish the pH of the suspensions.

5) Sieves: U. S. Standard Sieves of 8, 10, 18, 35, 60,

120, 230, and 325 mesh were used to subdivide several

of the sediment samples.

6) Stirrer: To aid the Solution of sediment samples, a

magnetic stirrer and teflon covered stirring bar were

used.

7) Stoppers: Cotton, covered with cotton gauze, Stoppers

were used to permit the experiment to be conducted

open to the atmosphere.

Njatura Carbonate Samples Used. The natural carbonate samples

used in this study were collected on and about the reefs on

the Campeche Bank in the Gulf of Mexico. The samples include:

#1 - Montastrea annularis plus encrusting foraminifera

2 Coral and shell sand from Alacran Reef

3 - Montastrea annularis

4 - Bottom sediment

5 -Strombus gigas Linne'

6P - Luci na pegnslniain' (fresh and unweathered)

6W - Lucinta p eA syl!vAn-ica Linne' (worn and weather-ed)

7 Olycy-mgris p -tinata Gmelin

13 mBttomn Sediment containing algal material.,

Sample numbers 1 and 3 were collected by Dr. Brian Logan

on Campeche Banik in the Gulf of Mexico. Sample numnbers 2, 5,

6F)6Wand 7 were collected on cruise 61-H-5 at Isla Desertora,I acran Reef, C ampeche Bank, in the gulf of Mexico. Sample

numbers 4 and 13 were bottom samples collected off the Campeche

Bank on cruise 59-H-7.I Sample numbers 2, 4, and 13 Were Subdivided by Sieving

into 10, 18, 35, 60, 120, 230, and pass 230 mlesh fractions.

Sample number 13 was further subdivided into 1/2", 3/8",

1/4"1, and 8 mesh fractions. See Table 1.

4 TABLE 1

Size Analyses

Mesh, 1 Diameter Sample #2 Sample #4 Sample #13

p'er i n ch mm % %.7 . -

10 >2 4.92 61.1 73.1

18 1-2 4.85 3.82 16.4

35 1/2-1 75.5 11.5 8.32

60 1/4-1/2 14.5 11.9 1.10

120 1/8-1/4 0.22 5.94 0.27

230 1/16-1/8 trace 4.61 0.414

<230 _ 11 9--10 0.44-

/U.S. Standard Sieve Sizes

6

Each of the above samples or fractions of a sample were

Split and a portion ground in an agate mortar and pestle for

two hours. The ground portions were used for the determina-

tion of the Ca and Mg++ in each samrple, and also for the

Solution experiments.

A summary of the analyses of the recent marine carbonate

sediments used in this study may be found in Tables 2, 3, and

4. Table 2 shows the percentage of CaCO3 and MgCO 3 present

in each of the unground natural sediment samples.

TABLE 2I++ +CA and Mg Content of Unground Natural

Carbonate Sediments-2/

-a-mpe -e. #. %ape# 76Mesh Size CaCO 3 MgCO 3 Mesh Size CaCO 3 MgCO 3 .

2-18 X 94.50 4.39 4-18 X 90.27 6.910

2-35 X 93.82 2.03 4-35 X 91.96 5.27

260 X 93.48 2.28 4-60 X 95.57 2.91

4-10 X 90.16 6.09 4-230 X 94.07 2.28

Table 3 reports the percentage of CaCO and MgCO3 present33

in each of the ground sediment samples. All analyses are

reported by mesh size fractions whether it has undergone |

grinding or not.

I!Average of three analysesk!

7

A comparison between Tables 2 and 3 indicate different

values of MgCO 3 content exist for ground and unground samples3

of the same size fraction. It is felt that this difference

is due to the heterogeneous nature of the sample. Binocular

examination of the samples reveal a number of echinoid spines

which are known to be very high in magnesium. It is felt

that the ground sample gives a better indication of the

average MgCO 3 present in each site fraction of each sample.

This suggestion is borne out as the difference in MgCO3 con-

tent decreases with decreasing sediment size.

TABLE 3

Ca and Mg Content of Ground Natural

Carbonate Sediments 3/

S atmp~ 7I 7a Sa#l % % .Mesh Size CaCO 3 MgC03 Mesh Size CaCO3 MsCO 3

1 87.32 6.95 5 96.23 0.00

2 18 93.39 2.87 6F 97.76 0.00

2-35 93.02 2.72 6W 96.66 0.00

2-60 92.64 2.93 7 96.12 0.00

3 95.19 0.00 13-1/2" 84.38 9.910

4-10 90.16 6.09 13 3/8" 84.88 9.52

4-18 89.70 6.53 13-1/4" 85.57 12.1

4-35 91.25 5.36 13-8 85.13 12.0

4=60 93.34 3.39 13"18 85.23 11.8

4-230 94.07 2.28

3/Average of two analyses

8

Table 4 shows the mineralogical composition of the

samples as determined by X-ray diffraction. X-ray analyses

of sample number 4 at the onset of grinding and after 2

hours of grinding show no significant alteration of calcite

to aragonite.

TABLE 4

Mineralogical Composition of Ground Natural

Carbonate Sediments

S amp le --# -I- V. SapeV.7Mesgkh SijZe A-ra aoie Calcite Mesh SizeArao ite Calcite

1 53 47 4-120 87 13

2-18 95 5 4- l20 4/ 87 13

2 35 97 3 4-230 84 16

2-60 95 5 5 100 --

3 100 -- 6F 100 -

4-10 74 26 6W 100 -

4-l04-/ 74 26 7 100 --

4-18 70 30 13-1/2" 31 69

4 8/ 71 29 13-3/8" 32 68

4-35 72 28 13-1/4" 34 66

4-5/ 72 28 13-8 31 69

4=60 86 14 13=10 41 59

4-604/ 86 1413=1 43 5

Lowenstam (1954) Assigns a rather large error, -10 per-

cent, to the analysis of calcite :aragonite ratios. The

relationship between the percentages of MgCO3 and aragonite

/Minera!Qgical composition at the onset of the grinding

9

in the sediment samples is shown in Figure i. As the per-

centage of aragonite increases in the samples, less MgCO3

is found. This can be easily explained as the MgCO 3 is

found as a calcite type structure, being present in a solid

solution series between calcite and dolomite, and the per-

centage of MgCO3 present would decrease as the ratio of

Calcite to aragonite decreases.

Solution 9xperiment.

Introduction.

In order to measure the resistance of recent natural

marine carbonate sediments to solution, sediment samples

were suspended in various salt solutions and after a number

of months, the amount of calcium and magnesium given up to

the solution was determined. The degree of solution of each

sample in the various salt solutions was then related to

sample Composition, mineralogy, and the particle size.

Procedure.

Five 0.5 gram samples of each of the sediment fractions

listed in Tables 2 and 3 were weighed out on glazed weighing

paper, folded, and set aside. One-hundred milliliters of

distilled water were pipetted into a 250 ml beaker, placed

on a magnetic stirrer, and after the stirring bar was added,

the pH was recorded. Then a 0.5 gram sample of a sediment

10

The Relationship Between Aragonite and lMgCO3 in Rcn

Marine Carbonate Sediment Samples

3000

400

-50

40

700

00

96000

00

0 2 81

% MgCO 3

IN THE SAMPLE

11

was added and the change in pH' recorded contifuously for

about 10 to 15 minutes. The suspension was then transferred

to a 250 ml erlenmeyer flask and stoppered with cotton and

gauze. This procedure was repeated until one 0.5 gram sample

of each Sediment sample was placed in distilled water and

the pH measurements recorded continuously.

The process was then repeated for the NaCl, MlgC 2,

NaCl + MgC+2, and NaHCO ion pairs of synthetic sea water.

The Solution experiments for each sediment sample were not

run in duplicate; however, the analyses for Ca and Mg

were the average of two determinations each. Duplicate

blanks of each solution, minus the sediment sample, were

also prepared.

The pH of each suspension was recorded after 1, 2, 3,

4, and 10 weeks, and at the end of the experiment. The pH

results are given in Appendix III.

After the final pH was recorded, the solution was

analyzed for dissolved Ca + and Mg + , using the procedure

given in Appendix I, Part F. The results of these analyses

are given in Appendix II.

C HA PT 9f R 11

RESULTS AND DISCUISSION

NaliCO 3Solution Experiment

The amiounits of Cali+ and M-g+ dissolved out Of the marinie

carbonate Sediments in NaHCO- 3are recorded in Appendix II,

Table 1. These analyses were obtained after the sedim-ents

had been in contact with the NaHCO 3solution for eight months.

Figure 2 relates the dissolved amounts of Ca+ aind Mg+ in

siillimoles to the percentage of MgCO 3in the sediments.

From Figure 2, it is evident that little change occurs

in the amount of dissolved Ca++, averaging about 0.15 milli-

moles per liter, from the ground sediment samples regardless

of Mg:CO 3 content. The amount of dissolved Mg ,ont the other

hand, increases with increasing MgOO content only to begin

to level off around 0.6 millimoles per liter at higher MgCO 3

percentages. The amount of dissolved Ca + from unground sedi-

ments is also about 0.15 millimoles per liter; however, no

dissolved Mg ++ is found in solution.4

Figure 4 indicates that the carbonate alkalinity-- Of

the solution of ground sediment samples in NaRCO 3is 2 to 3

1/arontea .al -~y =the HC0- plus C0 contributed to the

solution by the dissolving CC 3 adMC 3

C.. CO~ -2CO= (expressed as equivalents)

12

13

FIGURE 2

Dissolved CA"~ and XSg in 0.0023 H NaHCO 3 with Respect

33

LEGEND

Obissolved Ca++ from ground marinecarbonate sediment samples

EDissolved CA+ from unground Matinecarbonate sediment samples

ODissolved Mg+ fromh ground Marinecarbonate sediment samples

ObDissolved Mg++ from unground marinecarbonate sediment samples

FIGURE 3

Dissolved Ca++ and Mg+ in Distilled Water with Respect

to Percent MSC0 3 in Sediment Samples

LEG-END

*Dissolved Ca ++ from ground marine4carbonate sediment samples

*Dissolved Ca ++from unground marine

carbonate sediment samples++

0 Dissolved Mg from ground marinecarbonate sediment samples

O3 Dissolved Mg + from unground marinecarbonate sediment samples

0

0ais

IS 0 i=

0 2 4 6 8 10 12

% MgC0s

IN THE SAMPLE

00 6 10 1

% 0 0CO

INTESML

14

FIGURE 4

Carbonate Alkalinity Resulting from the Solution

of Sediment Samples in NaHCOI3, Distilled Water,

and NaCi with Respect to Percent MgCO3

in Sediment Samples

LEGEND

C.A. from ground marine carbonatesediment samples in NaHCO 3

A CA. from unground marine carbonate

sediment samples in NaHCO3

0 CiA. from ground marine carbonatesediment samples in distilled water

OC.A. from unground marine carbonate

sediment Samples in distilled water

U C.A. from ground marine carbonatesediment samples in NaCl

C.A. from unground marine carbonatesediment sampics in NaCl

00 00

00

0 2 4 6 8 10 12

% MgCO3

p IN THE SAMPLEI

times that of the unground carbonate Sediments, and approaches

a constant value at the higher percentages of MgCO3.

Figure 5 indicates that the carbonate ion concentration

resulting from the solution of ground carbonate samples in

NaHCO 3 solution is not leveling off with greater percentages

of MgCO3 present in the samples. In NaRCO3 solution, the

ion products of [Ca] [COt]: and [Hg -] (GO change very

little as MgCO3 increases in the sample (Figure 6).+

The indication from Figure 2 is that Mg may have some

influence on the solubility of CaCO 3 in ground carbonate sedi-

ments. The Solubility of Ca++ from aragonite and magnesium

calcite appears the same regardless of the amount of dissolvedM++ -

Mg present.

The situation with regard to the coarse, unground car-

bonate sediments is more complex. Magnesium from the coarse

sediments is not going into solution (Figure 2). The reason

for this is not known with certainty; however, there are 2 pos-

sible explanations.

The first explanation is derived from the fact that the

sediments were washed prior to any subdivision or grinding.

During the removal of the sea salts from the sediment surfaces

by distilled water, it is possible that the limited amount+4

of Mg at the surface was removed. Subsequent solution of

the unground sediment samples would result in lower dissolved

16

FIGURE 5

Carbonate Ion Concentration of the Solution of Sediment

Samples with Respect to Percent MgCO-3

in Sediment Samples*

LEGEND

ICO:] of dissolved marine carbonatesedimeent samples in NaHCO 3

O[CO=] of dissolved marine carbonatesediment samples in distilled water

[ICOn] of dissolved marine carbonatesediment samples in NaCi

X [CO ] of dissolved marine carbonate3sediment samples in MgC1 2

+[CO I of dissolved marine carbonatesediment samples in NaCl + MgCl2

*See Appendix II, Tables 6 and 7

1- Xo

0 A 0 1

% MgCO3IN THE SAMPLE

17

FIGURE 6

Iont Products of the Solution Of Sediment Samples

in NaHCO 3 Distilled Water, and NaCi with -Respect

to Percent MgCO, 'n Sediment Samples*

LEGEND

t(ca'" [CO:] of 'ligsolved marine carbonat,33

A[Mg"I] 'C0 ] of dissolved marine carbonate3se, iment samples in NaHCO0

OCa" ] 'CO]j of dissolved marine carb-nat-3sediment samples in distilled vater

* [mg ] LCO%] of dissolved marine Larbonatesediment samples in distilled water

O [Ca ++] [CO'] of dissolved marine carbonate3 sediment samples in NaCl

0 [Mg++j C031 ol' dissolved marine carbonat3 ediment samples in NaCl

*See Appenuix II, Table 6 and 7

400

3000-00

I- 09-

0 Op

1000 0

/ ... ~A LA

0 2 4 6 8 10 12

% MgQCO3I N THE SAMPLE

i8

+

Mg values than for ground samples. It is further suggested

that the washing would not remove all available surface Ca

as the samples are composed of 88 percent plus CaCO 3 , No

change in dissolved Ca++ would therefore be expected.

The second explanation stems from the possibility that

in the marine environment, the sediments, or actually the

sediment surfaces, are in equilibrium with sea water. Upon

solution in NaHCO, an "equilibrium shift" results in the

observed value of dissolved Ca+ and Mg . The inner parts

of the sediments are not in equilibrium with sea water and

grinding exposed a great number of these surfaces. There-

fore, upon solution of the ground sediments, different values

of dissolved Ca+ + and Mg+ result.

Figure 4 indicates that the carbonate alkalinity for

the ground carbonate samples in NaHCO3 is 2 times greater

than for the unground sediment samples when more than 2 per-

cent MgCO3 is in the sediment sample. This also may be due

to washing the sediments, as CO- and HCO3 from the dissolving3 3

MgCO3 of the unground sediments would be lost to the distilled

water wash.

The carbonate alkalinity resulting from the solution

of the ground samples approaches constancy with increasing

MgCO3 content in the sediment, indicating that additional

HC0 and CO' ions are not going in solution (Figure 4).3 3

19

The (CO ]Z/ of the solution, which takes K2 and pH into con-

sideration, may be approaching a constant value (Figure 5).

This value gives some indication as to the status of equi-

iibrium; however, only the ion product attaining a constant

value will tell whether equilibrium is reached, In the ease

of NaHCO CaCO equilibrium is approached in eight months3' 3

by sediments with high MgCO3 content (Figure 6); however,

MgCO 3 equilibrium is not yet reached as indicated by the

increasing [Mg + ] [COn] ion product.

Distilled Water Solution Experiment

The amounts of Ca + and Mg+ dissolved out of the marine

carbonate sediments in distilled water are recorded in

Appendix II, Table 2. These analyses were obtained after

the sediments were in contact with distilled water for six

months. Figure 3 relates the dissolved amounts of Ca and+ +

Mg in millimoles to the percentage of MgCO 3 in the sedi-

ment sampies.

The dissolved amounts of Ca++ from the ground sediment

samples increase from 0.45 to 0.55 millimoles per liter with

increasing MgCO The dissolved amounts of Mg + , meanwhile,

increase from 0 to almost I millimole, and there is only

_ K2

2 -LCO- (C.A.] K 2-3 H+ +2K -2

20

a Slight suggestion that it is beginning to level off for

the same ground sediment samples.

Comparing Figures 2 and 3, it is observed that for the

ground samples, the dissolved Ca++ is greater in distilled

water (0.6 mM/L) than in NaHCO 3 (0.2 mM/L). Likewise, the

dissolved Mg+ has reached almost 1 millimole per liter in

distilled water while reaching only 0.6 millimoles per liter

in NaHCO3,.

Figure 4 indicates that the carbonate alkalinity for

ground samples in distilled water is 50 percent higher than

the carbonate alkalinity for unground sediment Samples in

distilled water. For ground sediment samples in distilled

water, the carbonate alkalinity (Figure 4) and the carbonate

ion concentration (Figure 5) are both steadily increasing

with greater percentages of MgCO 3 present. The ion products

of [Ca+ I [COm] and [Mg+*] [COQ] also continue to increase

with higher percentages of MgCO 3 in the sediment for the

distilled water solution experiment (Figure 6).

Figure 3 is interpreted as indicating that the Mg +

in the ground samples has influenced the solubility of Ca + +

dissolving from calcite. This has resulted in there being

a slight increase in the solubility of calcite over aragonite.

The fact that a considerable amount of the calcite is dis-

solving is indicated by the rapidly increasing amounts of

21

Mg going into solution. The absolute value of Ca + dis-

solved appears to increase slightly with increasing MgCO 3

in the sediment, The presence of Mg has increased the

overall solubility of calcite in natural marine carbonates.

The interpretation of the dissolved Ca+ + and Mg+ + from

the unground carbonate sediments is more complex. The mndi-

cation from Figure 3 is that the Mg ++ dissolved from unground

sediments is depressed, and the Ca+ dissolved is increased.

Two possible explanations have been outlined in the previous

section.

The first explanationj derived from the washing history

of the coarse sediment samples after removal from the marine

environment, does explain the low values of dissolved Mg+ +

from unground sediment samples. From this explanation the++

value of dissolved Ca is not expected to change from ground

to unground sediment samples; however, Figure 3 indicates

that the dissolved Ca++ from unground sediments is 0.1 to

0.2 millimoles per liter greater than dissolved Ca+ from

ground sediment samples. This is reasonable when the chemistry

of the CO2 system is taken into consideration.

The CO2 system can be represented by the following:

CO +HC- H C C) H HCO --- N+C+0O2 + H 2 O 2 3 " + 3- ++

+ Ca

-- CaCO 3

22

In distilled water, the CO2 system is established upon the

solution of CaC1O 3 and a speCifio carbonate alkalinity is3

maintained. In the case of the ground samples, CO- and3

HCO3 is contributed to the solution through dissolving CaCO 3

and MgCO3 , For the unground samples, however, less MgCO3

is available for solution which results in more CaCO dis-3

Solving toL make up the CO deficiency.

The second explanation suggesting an "equilibrium shift"

could result in decreased Mg solution and increased Ca

Solution.

The differences in Solution between distilled water

and NaHCO 3 can be attributed to the common ion effect in

the NaHCO 3 solution. At pH values of around 8, the major

contributing ion in the CO 2 system is HCO3 and the concen-

tration of CO present Will therefore depend upon the HCO33

concentration. As CaCO 3 is added to NaHCO 3 the system will

establish equilibrium with less carbonate from CaCO 3, as

CO0 is already present from the dissolved NaHCO. For this3 3

reason, Ca+ + and Mg + are dissolved more readily in diatilled

water than NaHCO 3.

The steady increase in carbonate alkalinity with higher

percentages of MgCO3 in the sample for ground sediment samples

indicates that more HC03 and CO- are being added to the dis-3 3-

tilled water, when pH and K2 are taken into consideration,

as shown by Figure 5,

23

The changing [Ca + ] [C-] and (Mg ] jC0*] ion products

in distilled water indicate that equilibrium is not yet

reached, although the (Ca ++ [C0;] ion product appears to

be much closer to attaining it (Figure 6).

NaCl Solution Experiment

The amounts of Ca and Mg+ dissolved out of the marine

carbonate sediments in NaCI are recorded in Appendix II,

Table 3. These analyses were obtained after the sediments

were in contact with NaCl for 3 months. Figure 7 relates

the dissolved amounts of Ca+ + and Mg+ + in millimoles to the

percentage of MgCO3 in the sediment samples.

From Figure 7, it can be seen that the dissolved amounts

of Ca+ + from the ground sediment samples are relatively con-

stant, averaging between 0.9 and 1.0 millimoles per liter

regardless of the MgCO3 content, The dissolved amounts of

Mg + + , however, increase rapidly at first, and begin to level

off around 1.2 millimoles per liter at higher MgCO3 contents.

The dissolved amounts of Mg + for unground sediment samples

average about 0.1 millimoles per liter, and Ca+ + about

1.55 millimoles per liter (Figure 7).

The dissolved Ca + and Mg + for ground samples in NaC!

are greater than the dissolved Ca + and Mg + for ground

samples in distilled water (Figures 3 and 7).

'4

24

FIGURE 7

++ +Dissolved Ca and Mg++ in 0.405 M NaCi with Respect

to Percent MgCO3 in Sediment Samples

LEGEND

0 Dissolved CA++ from ground marinecarbonate sedime nt samples+

EDissolved Ca+ + from unground marinecarbonate Sediment samples

o Dissolved Mg from ground marinecarbonate sediment samples

dDissolved Mg + from unground marinecarbonate sediment samples

to- -

0.5 0

00

0p- -

0 2 4 6 8 10 12

% MqCO3 IN SAMPLE

25

The carbonate alkaliiity for the ground samples in NaCl

is somewhat higher than for the unground carbonate samples

with some indication of leveling off at higher percentages

of MgCO 3 (Figure 4).

Figure 5 shows the carbonate ion concentration for NaCi

to be quite high with the absolute value not yet approaching

constancy.

Figure 6 indicates that the [Ca + + It[o ior product

for the NaCl solution experiment may be beginning to level

off; however, the [Mg+ ] ![COO] ion product is still rapidly

increasing at almost a linear rate.

A considerable amount of calcite is dissolving as indi-

cated by the rapidly increasing Mg+ + in solution (Figure 7).

Much of the dissolved Ca must therefore come from the cal-

cite present in the carbonate samples. The absolute value

of Ca dissolved appears to be little changed by increasing

Mg + , suggesting that there is little difference in the solu-

bility of aragonite and the magnesium calcite present in

NaCl solution. Seidell (1940) reports that aragonite is

somewhat more soluble than calcite in NaCl. The indication

is, then, that Mg+ + does influence the solubility of CaCO 3

as seen in Figure 7.

Figure 7 indicates that between 1.5 and 1.6 millimoles

per liter of Ca + + dissolved from unground natural size

2,6

sediments, whereas only about 1 mIllimole per liter of Ca++

dissolves from the ground sediments. Likewise, only a very

small amount of Mg dissolves from the unground sediments,

about 0.1 millimoles per liter, whereas from 0.5 to 1.2 milli-

moles per liter of Mg+ dissolves from the ground sediments

(Figure 7). The reason for these differences again is not

known with absolute certainty; however, there are 2 possible

explanations, one involving washing history) and the other

involving equilibrium both of which have been previously

outlined under the distilled water discussion.

The increa - in the solubility of CaCO3 by alkali

chlorides have long been recognized (Seidell, 1940). This

has been reaffirmed in the present study; however, it is

difficult to say whether the solution of Kg is ultimately

greater or less in NaCl than in distilled water. A compari

son between Figures 3 and 7 indicates that at 10 to 12 per-

cent MgCO in the sample, 1.2 millimoles per liter of Mg

are dissolved in the NaC1 solution compared to just under

! millimole per liter of Mg+ + present in distilled water.

The carbonate alkalinity plot (Figure 4) for ground

sediment samples in NaCl indicates that the total HCO- and

CO content of the solution is not yet beginning to show3

signs of leveling off at high percentages of MgCO3. The

carbonate ion concentration in the NaCl solution, from

27

Figure 5, is continuing to rise as MgCO 3 increases in the

sediments, The latter value takes into account any pH dif-

ference as it was previously seen.

Figure 6 indicates, however, that the [Mg+] + CO] ion

product is still rapidly increasing at 12 percent MgCO 3.

The [Ca ] [COz] ion product On the other hand, may be

approaching equilibrium with increasing MgCO 3 . It is evi-

dent that NaCl does favor the solution of Mg , and the dis-

solved Mg++ does not approach equilibrium in the NaCI solution

(Figure 6).

MgC1 2 Solution Experiment

The amounts of Ca and Mg + dissolved out of the marine

carbonate sediments in M-gC1 are recorded in Appendix Ii,

Table 4. These analyses were obtained after the sediments

were in contact with MgC!2 for eight months. Figure 8 relates

the dissolved amounts of Ca+ + and Mg++ in millimoles to the

percentage of MgCO3 in the sediments.From Appendix II, Table 4, it is evident that something

unusual is taking place, The Mg + content of the MgCl 2 solu-

tion is not increased by either the dissolving ground or

unground samples.

The dissolved amount of Ca from the ground sediment

samples is relatively constant for those carbonate samples

28

FIGURE 8

Dissolved Ca in MgCl 2 and Nati + MgCl 2 with Respect

to Percent MgCO3 in Sediment Samples

and

Carbonate Alkalinity Resulting from the Solution

of Sediment Samples in MgCl2 and NaCI + gCl 2with Respect to Percent MgCO3

in Sediment Samples*

LEGEND

6 Dissolved Ca from ground marine carbonatesediment samples in MgCI 2

0 Dissolved Ca++ from unground marine carbonatesediment samples in MgCl2

UDissolved Ca++ from ground marine carbonatesediment samples in NaCI + MgCl

++EDissolved Ca from unground marine carbonate

sediment samples in NaCl + MC12

2!I

*Carbonate Alkalinity is equal to dissolved Ca ,as no add-tional !g+ is contributed to the solution from the sediment.

3U

0

0 20 4~ 6 80 1

% MgCO3IN THE SAMPLE

29

containing calcite, with the dissolved amount of Ca from

pure aragonite being much lower (Figure 8). The amount of

Ca dissolved from unground carbonate sediments is greater

than the amount of Ca++ dissolved from the ground sediment

samples for the MgCi 2 solution experiment (Figure 8).

Figure 8 indicates that the carbonate alkalinity from

the dissolving ground sediment samples in MgCl2 is on the

average 0.4 millimoles per liter less than the carbonate

alkalinity of the unground sediment samples. The carbonate

ion concentration of the ground carbonate sediments in MgCl 2

may be beginning to level off according to Figure 5, and the

[Ca + ] [CO] ion product from Figure 9 has about reached

a constant value,M++

The absence of Mg in solution can not be explained

either by washing or equilibrium as in the case of NaHCO3,

distilled water, or NaCl. One very possible explanation,

however, is that the high concentration of MgCl2 in solution

inhibits Mg from dissolving. Common ion effects such as

these are well known in solution chemistry.

Seidell (1940) reported that alkaline earth chlorides

depress the solution of CaCO 3 . A comparison between

Figures 3 and 8 shows quite convincingly that the solubility

of Ca+ + is greater in MgCl2 solution than in distilled water.

The increase in the dissolved amount of Ca , above that

11!30

FIGURE 9

Ion Pro-ducts of the Solution of Sediment Samples

in, MgCi2 and NaCi + MgCL2 with Respect

to Percent MgCO3 in Sediment Samples*

LEGEND

.--(Ca + + ] [CO ] of dissolved marine carbonatesediment samples in NaCl(from Figure 6)

X [Ca + + ] LCO3] of dissolved marine carbonatesediment samples in MgCI2

+[Ca ++ ] (CO-] of dissolved marine carbonatesediment samples in NaC1 + MgC1 2

*See Appendix II, Table 7

700+

600-

%~400x

0

& 200x

i 100/

0 2 4 6 8 10 12% MgCos

IN THE SAMPLE

31

which dissolves from pure aragonite, can be attributed to

the high magnesium, calcite.

in order to justify the difference in the amounts of

Ca+ + dissolved from ground and unground sediments, we must

return to the alternatives presented under the discussion

of NaHCO The most likely explanations, derived from3.

either preliminary washing or differences due to an equilib-

rium shift, seem to apply,

It is evident from Figure 5 that the [CO;], from the

dissolving ground carbonate sediments in MgCi2, may have

begun to level off so that further increase in available

MgCO 3 will result in very little additional carbonate being

added to solution.

The [Ca [CO] ion product has nearly reached equi-

librium for ground sediments in MgCI indicating that the

solution will not dissolve greater quantities of CaCO 3

(Figure 9).

NaCl + MgCl 2 Solution Experiment

The amounts of Ca and Mg + dissolved out of marine car-

bonate sediments in NaC! + MgCl are recorded in Appendix II,2

Table 5. These analyses were obtained after the sediments were

in contact with NaCl + MgC!2 for eight months. Figure 8 relates

32

the dissolved a.iounts Of Ca++ and Mg+* in millimoles to the

percentage of MgCO 3 in the sediiments.

As in the, case of MgC12, Appendix I!, Table 5, indicates

unusual behavior on the part of sediment solution. Again++

Mg is not dissolved from either the unground or ground

sediment samples.

Figure 8 indicates that the amount of Ca dissolved

from the ground sediment samples slowly increases at a rather

constant rate from just under 1.9 millimoles per liter to

more than 3 millimoles per liter for 0 percent MgCO to 12

percent MgCO3 , respectively. A comparison between Figures 7

and 8 indicates that the total Ca + dissolved by NaCl + MgCl

is roughly the Same as the Ca + dissolved from NaCI and MgCI2

added separately.

Figure 5 indicates that the [CO' for NaCI + MgCI2 con-

tinues t-o rise with increase in MgCO in the sediment, and3 - +

Figure 9 shows that the rapidly increasing [Ca -] [COi ion

product for NaCI + Mg'Cl 2 does not level off and reach a

constant value.

Figure 7 indicates that NaC1 dissolves over 1 millimolef++

of Mg per milliliter; however, when NaC1 is combined with

MgC 2, no Mg + is in solution after eight months (Figure 8).

Magnesium chloride exerts greater control over the solution

of carbonates than NaC!. It is again suggested that the MgC12

in solution tends to inhibit the solution of Mg from the

sediments.

33

The differences in amount of Ca*+ dissolved from ground

and unground sediments for the NaCi + MgCI 2 solution experi-

ment are attributed to either preliminary washing, or to the

equilibrium shift.

Figure 5 indicates that the [CO'] for NaCi + MgCl2 may

be leveling off, With increasing MgCO 3 in the sample the3

carbonate ion concentration in NaCI + MgCl 2 solution may

become constant. The [Ca++] ICO3 ion product from Figure 9,

however, shows that equilibrium is not being reached in the

++-case of NaC1 + MgCl2. it is also evident that the [Ca 1

[CO3] ion product in NaCl + MgCl2 is merely the additive

effect of the ion products for NaCi and MgCl2 added individually.

Figures 10 and 11 summarize the dissolved amounts of

Ca + and Mg+ + with respect to the percent MgGO3 in the sedi-

ment samples for each of the solutions used in the experiment.

The solution of Ca++ is inhibited by NaHCGO 3, followed by dis-

tilled water, NaCl, MgCl2 , and NaCi + MgCl2. The solution of

Mg is inhibited by MgCl2 and NaCl + MgCl2, followed by

NaHCO3, distilled water and NaC1. Figures 10 and 11 show the

relationship between dissolved amounts of Ca++ and Mg with

increasing MgCO3 content in the sediments for each solution.

34

FIGURE 10

Dissolved Athounts of Ca+ with Respect to Percent MgCO 3in Sediment Samples in the Ioni Pairs of Synthetic

Sea 'Water

3. -12%

0% W

11 2%w

+

0- fNoHCO3 DISTILLED No CL MgCL2 NoCL

WATER +MgCL2SOLVENT

35

FIG URE 1'.

Dissolved Amounts of Mg++ with Respect to Percent mgCo3

in Sediment Samples in the lon Pairs of Synthetic

Sea Water

wow

0U- 49o l

C/0 -m 2

E 0- u 0 m..- 'T

NaHlCO3 DISTILLED Na:CL MgCL2 NaCl.WATER +MgCLt

SOLVENT

36

The Effects of Particle Size and pH

From the pH values observed during the first 10 or

15 minutes as each sediment went into solution, it was evi-

dent that particle size did affect the rate of solution.

The ground samples, in almost every case, go into solution

very rapidly as indicated by the immediate rise in pH. Por

the unground samples the rate of pH change decreases with

increasing grain size. This indicates that the initial rate

of solution is dependent upon grain size. Figure 12 shows

a typical family of pH curves for various size carbonate

sediments as they dissolve. Appendix III contains the pH

values of each dissolving carbonate with time.

All differences in pH, with respect to particle size,

disappear as the solution approaches equilibrium. No cor-

relation appears to exist between particle size and ultimate

solubility in any of the solutions used.

It is apparent that the final pH values do not reflect

solubility in NaHCO 3, distilled water, and NaCl. For MgCl2and NaCl + MgC!2 , however, the final pH values of th& unground

sediment samples are higher than the final pH values of the

ground sediment samples.

Figure 8 indicates that the dissolved Ca from unground

carbonate sediments is greater than the dissolved Ca+ + from

ground carbonate samples. Inasmuch as no C02 is contributed3

37

FIGURE 12

Typical Variations in pH for Various Grain

Size Carbonate Sediments

LEGEND

ApH variationS for 18 mesh sizeSediments

* pH variations for 35 mesh sizesediments

O pH variations for 60 mesh sizesediments

pH variations for 120 mesh Sizesediments

pH variations for 230 mesh sizeSediments

XpH variations for all sedimentsground for 2 hours

r- ClI

CD

OE 4 0 '

38

to the solution from dissolving MgCO in either MgCl or3 2

+++NaCl + Mg~i2, dissolved Ca and 0O are directly related

As dissolved Ca++ increases in a solution, CO: also increases.3

The relationship between pH and CO can be written as:

CospH pl 2 + log

The higher values of dissolved CA" for the unground samples

are thus accompanied by higher pH values in both MgCl 2 and

NaCl + MgCl 2 solution experiments, indicating that a rela-

tionship between pH and the Solution of Ca + exists. An

understanding of the mechanisms underlying this relationship

will require further work.

The pH values reached during the initial stages of solu-

tion are as high as 9.7, which are considerably greater than

the final equilibrium pH values of 8.2 to 8.4 (Appendix Ill,

Tables 1-5). It is suggested here that supersaturation

occurs during initial solution of marine carbonate sediments

in the ion pairs of sea water. Reprecipitation of the cations

is indicated by the lower pH values after 1 week.

The situation with regard to sea water is somewhat dif-

ferent. The pH of each sediment sample going into solution

was followed for eight months (Appendix III, Table 6). No

significant change in pH occurs during the first 15 minutes

39

except for pure aragonite samples which decrease somewhat.

Little Change in pH occurs durinig the entire experiment for

ground sediment samples. The urigroufd Sediment Samples,

however, raise the pH of the solutioni fro about 8.0 t

1 week to 8,2 at 10 weeks.

CHAPTER IV

CONCLUS IONS

The results of this study indicate:

i The grain size of recent marine carbonate sediments have

no relationship to the amount of Ca or Mg dissolved by

the major ion pairs of synthetic sea water and distilled

Water.

2. The pH of solutions containing MgCl2 reflect, to some2

extent, the amount of Ca+ dissolved.

3. The presence of MgCl2 in solution inhibits Mg++ from

dissolving out of sediments, and promotes the solution of

Ca The indication is that MgC1 2 is the major controlling

ion pair of synthetic Sea water, and the effects of the other

ion pairs are relatively unimportant on the solution of Ca++

+

and Mg++

4. The magnesium calcites present in the marine carbonate

sediments of this study are more soluble than 100 percent

aragonite in MgC12 and NaC1 + MgCl 2 1 but are not noticeably

different than pure aragonite in NaHCO 3, distilled water,

and NaC!.

40

41

5. At the conclusion of the solution experiments, it is

evident that when large amounts of HMgCO 3 are present in the

sediment, the ion, product of -Ca CO3] in the MgC 2 soiu

tion nearly reaches equilibrium, and the ion products of

[Ca + +] [CO] and (Mg++], [CO do not reach equilibrium in

NaHCO3 , distilled water, NaCi, or MgCl2 + NaCli

6. The pH values of the suspensions increase rapidly to a

maximum during the first 15 minutes in most cases and

decrease to a constant lowe value after the first week.

This is considered as evidence that during the first 15 minutes

of solution, more cations are in solution than are present

after a week. it is suggested that the initial solution of

each sediment is followed by reprecipitation until an equi

librium value is reached.

BILIOGRAPHY

Chave, X. P,., 1 954, Aspects of the biogeoche-mistry of mag-ne siurm:- 1 . Calcareous marine organisms: lout.GeoLogy, V. 62,) #5, p. 266-283.

Garrels, R. M., Thompsoni, M. E., and Siever, R., 1960,Stability of some carbonates at 250 C and 1 atm. totalpressure: Am. j. cAV. 2 58, p . 4 02-4 18,

Holland, H. D.1, Segntit, E. R., and Oxburgh, V. M., 1960,Carbonate solubility at high temperatures and pressures:.N.S.F. Annual Report G6069, Princeton, N. J.

Kitanoo Y., and Furut~u, T., 1959, The state of a smallamount of magnesium contained in calcareous shells:Bull1. Chem. Soc. Jap, V. 33, #1, P. 1-4.

Lowenstam, H. A., 1954, Fa-ctors affecting the aragonitecalcite ratios in carbonate-secreting marine organisms:Jour. Geology, V. 62, #3, p. 284-322.

Seidell, A. , 1940, Solubilities of inorganic and mnetalorganic compounds, p. 273-275, Van Nostrand, N. Y.

Weyl, P. K., 1961, The carbonate saturometer: Jour. Geology.V. 69, #,p. 32-44.

42

9?? N D 1 X I

APPENDIX I

A. Preparation of Calcium Ion Standard Solution

Reagents.

1) 6N Hydrochloric Acid. Dilute 485 ml of reagent grade

concentrated HCl (S.G. j1.9) to 1 liter.

Equipmen.

1) Hot plate.

2) Glassware:

1-200 ml tall form beaker

l-I liter volumetric flask

1 short stem funnel (slightly greaterin width than the tall form beaker).

Dry several gm reagent grade CaCO3 overnight in a low

temperature (850C) drying oven. Cool in a desicCator and

weight-out 1.000 gm. Quantitatively transfer the weighed

CaCO to a 200 ml tall form beaker and add 20 ml of distilled

water. Place the short stem funnel in the beaker". Through

the funnel add several ml of 6N RC! to dissolve the CaCO 3

completely. The solution is carefully boiled for 5 minutes

to remove CO2 completely, Cool the solution and quantitatively

transfer to a I liter volumetric flask. Dilute the solution

1 /The funnel prevents any loss of Ca ++ either due to theeffervescence of the carbonate while dissolving, or due tothe bubbling of the solution during boiling.

44

45

to 1 liter with boiled distilled water arnd stopper. The

titer of the calcium ion standard is 0.4008 mg Ca+ /ml.

B. Standardization of EDTA for Ca Analysis

1) Methyl Red Indicator (0.02 percent dissolved in

ethanol),

2) 2N Sodium Hydroxide, Dissolve 80 gm of reagentgrade NaOH pellets in distilled water and make tovolume of 1 liter.

3) Calcium Red Indicatori One gram of 2-hydroxy-l-(2-hydroxy-4-sulfo-i-ntaphthylao) -3-naphthoicacid is ground and mixed with 100 gm of reagentgrade NaCi in a mortar and stored in the dryform in a dark colored bottle.

4) EDTA Standard Solution. Dissolve 80 gm of thedi-sodium salt of ethylenediaminetetraaceticacid in 20 liters and allow to stand at roomtemperature with occasional stirring, for

several days.

Equipment.

1) Glassware:

1-50 ml automatic burette with automaticzero point

1-10 ml pipette

3-250 ml erlenmeyer flasks.

With the volumetric pipette, tranfer 1"10 ml aliquot

of the calcium ion standard into each of the 3250 ml erlen-

meyer flasks. Adjust the total volume of solution to about

50 ml with distilled water, Add 1 drop of methyl red and

then add 2N NaOH dropwise until the solution just becomes

46

basic (methyl red turns to yellow). Add an excess of 15 ml

of 2N NaOH (the titration must be carried Out at a pH of

about 13). With a spatula, add approximately 100 mg of the

solid calcium red indicator and titrate with EDTA until 1

drop turns the indicator from red to blue. The triplicate

titrations should agree within 0.05 ml.

The EDTA titer is calculated As follows:

Titer (mg/ml) (ml EDTA required)

C. Standardization of Mg-EDTA for Total Ca++ + Mg + + Analysis

Reagents,

1) Methyl Red (0.02 percent dissolved in ethanol).

2) Ammonium Hydroxide-Ammonium Chloride Buffer. Mix570 ml of reagent grade NH4 0H and 67.5 gm of reagentgrade NH 4C1 and dilute to 1 liter.

3) Eriochrome Black T Indicator (EBT). Add 0.4 gmof EBT to a mixture of 10 mi of 95 percent ethanoland 30 ml 2, 21, 2" nitrilotriethanol. The indicatorshould be stored under refrigeration when not inuse.

4) Mg-EDTA Standard Solution. Mix 160 gm of the di-sodium salt of ethylenediaminetetraacetic acdd and40 gm of reagent grade MgCl2 f6H 2 0, dilute to20 liters, and allow to stand at room temperature,with occasional stirring, for several days.

1) Glassware:1-50 ml automatic burette with automatic

zero point

1-10 ml pipette

3-250 ml erlenmeyer flasks.

47

With the volumetric pipette, transfer 1-10 ml aliquot

Of the calcium ion standard into each of 3-250 thl erlenmeyer

flasks. Adjust the total volume of solution to about 50 ml

with distilled water. Add 1 drop of methyl red and then add

NH 4OHNH4 Cl buffer solution dropwise until the solution Just

becomes basic (methyl red turns to yellow). Add an excess

of 2 ml of the buffer (the titration must be carried out at

a pH of about 9.5). Add 2 to 4 drops of EBT indicator, and

titrate with Mg-EDTA until 1 drop turns the indicator from

red to blue. The triplicate titrations should agree within

0.05 ml.

The Mg-EDTA titer is calculated as follows:

Titer (mg Ca++/ml) a (10.00) (0.4008 mp Ca/ml)(ml EDTA required).

D. The Determination of Ca+ + and Mg + + in Natural CarbonateSediment Samples

Relgents.

1) 6N Hydrochloric Acid (see Appendix I, Part A)

2) Methyl Red Indicator (see Appendix I, Part B)

3) 2N Sodium Hydroxide (see Appendix I, Part B)

4) NH4 OH-NH 4 Cl Buffer (see Appendix I, Part C)

5) Calcium Red Indicator (see Appendix I, Part B)

6) EDTA Standard Solution (see Appendix I, Part B)

7) Eriochrome Black T Indicator (see Appendix I, Part C)

8) MgmEDTA Standard Solution (see Appendix I, Part C)

48

Eqipment,

1) Hot Plate.

2) Glassware:

1-20 ml pipette

2-50 ml automatic burettes.

Wash free of chlorides approximately 100 mg of the

sample to be analyzed, and dry at 85'C overnight. Carefully

weigh the sample and transfer to a 150 ml beaker. Add 20 ml

of distilled water and several ml of 6N HCl. Boil carefully

for 2 or 3 minutes and cool. Transfer the dissolved sample

quantitatively to a 100 ml volumetric flask and bring to

volume with distilled water. Mix thoroughly and withdraw

420 mi aliquots, transferring them into 4-250 ml erlenmeyer+++

flasks. Two aliquots are for Ca analysis and 2 for Mg+ +

analysis. Reserve the remaining 20 ml in the volumetric

flask in case of loss.

Calcium Determination. Add one drop of methyl red to each

of the two solutions reserved for Ca++ analysis. Add2.N NaOH &ropwise until the solution just becomes basic,

and then add an excess of 15 ml NaOHo Adjust the volume of

the solution with distilled water to about 50 ml. Add approxi-

mately 100 mg of the calcium red indicator with a spatula

and titrate with EDTA until 1 drop turns the solution from

red to blue. Duplicate analyses should agree within 0.05 ml.

The calculations are as follows:

49

Ca + (mg) (ml EDTA required) ,0 -6,- (EDTA titer2 /)

GaCO % a(a .003 Sample weight 40.08

Magnesium Determinat-in. The analysis of Mg in carbonate

sediments is obtained by differcnce. The calcium present

in the sample is subtracted from the total Ca"+ + Mg+ as

determined by Eriochrome Black T and Mg-BDTA.

To the 2 remaining aliquots, add 1 drop of methyl red

to each and NH4 OH-NH4 Cl buffer until the solution just becomes

basic. Add an excess of 2 ml of the buffer, and adjust the

volume of solution in each flask to about 50 ml with distilled

water. Add 2 to 4 drops of EBT indicator and titrate with

Mg-EDTA until 1 drop turns the solution from red to blue.

Duplicate analyses should agree within 0.05 ml.

The calculations are as follows:

Mg (mg) ( ml Mg-ZDTA (2) (Mg24DTA,tite 0.08

Lrequired)

M CO (ag) 84.33Sample weight 24.32

-/See Appendix I, Part B

3/See Appendix I, Part C

50

E, The Determination of the Chloride Ion in the Componentsof Synthetic Sea Water

Rkzje-t .

1) Reagent Grade NaClo

2) 0.05 M K2 CrO 4 Solution, Dissolve 1.0 gm of reagentgrade K CrO in 100 ml of distilled water.

2 4

3) 0.05 N AgNO3 Standard Solution, Dissolve 8.5 gm ofreagent grade AgNO 3 in distilled water and diluteto 1 liter. Store in a dark colored bottle away

from light.

-~W~t.1) Glassware:

1-25 ml automatic burette

1-5 ml micro-burette

1 to 5 ml pipettes

3-250 erlenmeyer flasks.

The AgNO3 solution must be standardized before use.

Weight approximately 30 mg of NaCl accurately, and dissolve

in 20 ml of distilled water. Add 1 ml of 0.05 M K 2 CrO 4 as

the indicator, and titrate with the AgNO3 standard solution.

The end point is reached when 1 drop of the AgNO3 produces

a red precipitate of Ag2 CrO 4 which persists. Repeat until

triplicate analyses agree to within 0.05 ml. The AgNO3 titer

is then calculated as follows:

(gm NaCl x 35 46 x 1000) C1 pe r ml AgNO 3

51

To determine the chloride content of a solution, with-

draw an aliquot of the unknown with a volumetric pipette

calibrated to deliver i to 5 ml, depending on the amount of

solution available, and dilute it to about 20 ml with distilled

water. Add 1 ml of 0.05 M K2 CrO 4 indicator and titrate with

the AgN;O 3 standard solution until a single drop produces a

red precipitate that persists. For the chloride determina-

tion of MgCl2 , use the 5 ml micro-burette; for the others,

use the 25 ml automatic burette. The chloride content of

the solution is calculated as follows:

ml AgNO3

m a iquat x AgNO 3 titer = Cl .1

F. The Determination of Dissolved Amounts of Ca++ and Mgin Various Solutions

Reagents.

!) 6N Hydrochloric Acid (see Appendix I, Part A)

2) Methyl Red Indicator (see Appendix I, Part B)

3) 2N Sodium Hydroxide (see Appendix I, Part B)

4) NH4 OH-NH4 C1 Buffer (see Appendix I, Part C)

5) Calcium Red Indicator (see Appendix I, Part 5)

6) EDTA Standard Solution (see Appendix I, Part B)

7) Eriochrome Black T Indicator (see Appendix I, Part C)

8) Mg-EDTA Standard Solution (see Appendix I, Part C)

9) Reagent Grade NaCl

10) 0.05 M K2 Cr0 4 Solution (see Appendix I, Part E)

11) 0.5N AgNO3 Standard Solution (see Appendix 1, Part 9).

52

1) Hot Plate.

2) Glassware:

2-50 m! automatic burettes

1-25 ml automatic burette

1-5 ml micrb-burette

400 ml beakers

long stem funnels

various sizes of volumetric pipettes.

Filter each natural carbonate suspension through #42

Whatman filter paper into a clean dry flask. Do not wash.

Withdraw 4 equal aliquots from the flask so that at least

a few m! of the solution remain, and transfer to 4 erlenmeyer

flasks. Dilute the 4 aliquots to about 50 ml each with

distilled water, add 2 ml of 6N HC1 to each, and boil for

2 or 3 minutes. Two of the aliquots are reserved for Ca+ +

analysis, and 2 for Mg analysis.

Calcium Determination. To the 2 solutions reserved for Ca+ +

analysis, add I drop of methyl red. Add 2N NaOH dropwise

until the solution just becomes basic, and then add 15 ml

NaOH in excess. Adjust the volume of the solution to about

50 ml with distilled water. Add approximately 100 mg of the

calcium red indicator from red to blue. Duplicate analyses

should agree within 0.05 ml.

53

~Agooesium Determination. The analysis of Mg in the car-

bonate Sediments is done by difference. The calcium present

in the sample is subtracted from the total Ca++ +I Mg ++as

determined by Eriochrome M-ack T and Mg-EDTA.

To the 2 remaining aliquots, add 1 drop of methyl red

to each and a-dd NH OH-NH 4Cl buffer until the Solution just

becomes basic. Then add ant excess of 2 ral of the buffer.

Adjust the volume of solution in each flask to about 50 ml

with distilled water, and add 2 to 4 drops of EBT indicator.

Titrate with mg-EDTA until 1 drop turns the indicator from

red to blue. Duplicate Analyses Should agree within 0.05 mi.

The calculations for the sedimen~t samples in distilled

water, NaCl, and NaHCO -are as follows:

3

Dissolved Ca~() (lMgDT (BDTA titer-' 100L -required)ml a-iquot

Jeurd (Mg (mlt 5 /'D

Dissolved Mg ~ reui)d (Ms EgET DTA tier-/ (10010)

Dissolved Ca++ (v)] Z

/'See Appendix I, Part B

5/See Appendix I. Part C

54

Chloride Determination. Since several of the synthetic solu-

tions used contain Mg ) any evaporation of the liquid m edium

will result in concentration of Mg . By determining the Mg

+

and Cl- contents of the original synthetic solution, and the

Cl content at the end of the experiment, the Mg content

of the solution at the end of the experiment is calculated.

From the remainder of the filtered solution, after ali-

quota for Ca ++ and Mg analysis have been taken, a suitable

aliquot is withdrawn with a pipette and placed in an erlen-

meyer flask. Dilute to about 20 ml with distilled water,

and add 1 mi of 0.05 M K2CrO4 indicator. Titrate with AgNO3

standard solution until a single drop produces a red precipi-

tate that persists. Use the 5 ml micro burette for MgC! 2

solutions, and the 25 ml automatic burette for the others.

The calculations for the dissolved armounts of Ca ' and

Mg+ + in MgCl2 and NaCl + MgCl2 are as follows:

(ml EDTA

Dissolved Ca-+ -+eq)ir (EDTA tie 6 " (1000)L (ml requi (e (

6/See Appendix I, Part B

55

Dissolved (ml -Mg-EDTA ieLIrequI red) (MZDTA tiel)(1000)

Hg ~***~~ (ml aliquot)

Disle 2+~ a)1 48 / Concentratio' Hg concentration ofLC (+400 factori/ original solution 9/

i'See Appendix 1, Part C

8/ Cncenratin fator (ml AgN'O3) (*gNtO3 titer) (1000)

(Cl concentration of original

solution in%)

I/g+concentration of (ml Mg-EDTA) (MEDTA)(0)

(~) (mli-q-m-t} titeroriginal solution L L

mlDI EDTA 24.32

mialquo (EDTA titer) (1000)] 0

APPENDIX I

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