DECONSTRUCTING SEAWATER

“My soul is full of longing

for the secret of the sea,

and the heart of the great ocean

sends a thrilling pulse through me.”

--Henry Wadsworth Longfellow, The Secret of the Sea

Kellen Hansen

April 28, 2014

Chemistry 112

Section 402

Professor Jake Emmert

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ABSTRACT

This paper is a study of the ion content of seawater. This analysis was conducted via

testing of seawater samples, taken from the Deer Island. This paper will focus on an analysis of

collected seawater samples in order to establish which ions are present and in what quantity

through the use of qualitative, gravimetric, and titration analyses. The ions tested for were

ammonium, barium, calcium, halides, magnesium, potassium, sodium, and sulfate. Although

there were errors that occurred in this experiment, the results gathered were consistent and any

errors of significance are discussed and evaluated in the results section. The results of this lab

showed that in the gravimetric analysis, sulfate was found to be the most abundant, and in the

titration analysis, magnesium was the most abundant. In the qualitative analysis, the question of

which specific ions were present in the seawater was examined.

INTRODUCTION

The destructive and creative potential of the sea has fascinated man since the dawn of

time. For example, ancient Greek mythology revered Poseidon, ruler of the ocean, and attributed

the ocean’s ever-changing state to the god’s temperament; The Rime of the Ancient Mariner by

Samuel Taylor Coleridge fascinated and bemused audiences with its haunting depiction of a sea

voyage, ultimately culminating in a face-off between man and nature; or, finally, the persevering

myth of Atlantis that has permeated present-day popular culture, and only added to the mystery

of the sea. Only in the nascent days of oceanography when positivism finally imbued man with

an ideological framework and materialistic skillset to rationalize and understand the sea did the

veil begin to lift on a large scale in order to begin to satisfy this eternal curiosity.

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In order to fully understand the Earth’s oceans, it is important to understand what exactly

they are made of. The composition of seawater is the main focus of this paper, with an analysis

of the dissolved ions in seawater: which ions are present, and in what quantity? The reason why

most materials in seawater exist in their ionic forms is because of water’s ability to act as a

solvent. The two hydrogen atoms and oxygen atom that form a covalent bond to create water

makes a highly polar molecule, and as water is the primary component of seawater, this polarity

gives seawater its solvency. The most plentiful ions in seawater are chloride, sodium, sulfate,

magnesium, calcium, potassium, and bicarbonate. Along with hydrogen and oxygen, these ions

make up 99.9% of seawater (Szaflarski, n.d.).

When studying the composition of the oceans, it is important to look at the salinity of

seawater. Salinity can be vastly different depending on location. Sea salt cations (ions that have a

net positive charge) are formed mostly from weathering rock present on land; anions (negatively

charged ions), for the most part, come from the interior of the earth. It was once believed that the

ions in seawater came from volcanic eruptions, and the initial rapid release that occurred when

the earth melted; however, it has since been discovered that volcanic rocks do not contain

enough anions for this to be the case. Weathering is a slow process that is performed by water

and carbon dioxide, and while it may have occurred more quickly on younger earth, it would still

take several million years. As such, ions in seawater are thought to come from sedimentary rocks

(“Ocean Health,” n.d.).

The composition and concentration of the oceans has proven to be steady, and does not

change significantly over time (Ocean Health). Evidence and research suggest that, for the most

part, the ionic composition of Earth’s seawater has remained relatively steady for a long period

of time. This state of homeostasis has occurred because the removal rate of salts from the ocean

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is equal to the input rate, kept constant by rivers that carry dissolved ions to the ocean (“Ocean

Health,” n.d.).

For the purposes of this experiment, a near-total analysis of the seawater samples was

utilized. This entailed the use of qualitative experiments, gravimetric experiments, and titration

methods in order to determine the kinds of ions and their abundance in the seawater samples.

During each experiment, a standard solution was run alongside each step to ensure that a positive

test could be seen for the ion that was being tested for.

As qualitative analysis focuses on observations as opposed to measurements, the purpose

of the qualitative portion of this study was to discover which ions were present. The reactions

employed did not measure how many ions were present. For the qualitative analysis portion of

this study, a variety of precipitation and oxidation-reduction reactions were used in order to

determine which anions were in the sample being tested. In a similar fashion, cations were tested

for by using both precipitation reactions and flame tests.

Gravimetric analysis is used to study the measurement of a mass, and can be divided into

two categories: precipitation and volatilization. Precipitation was used in this experiment. In this

type of gravimetric analysis, an ion is isolated in a solution by a precipitation reaction. It should

first be filtered, washed, and then converted by calculation to a product of known composition.

Finally, the precipitate is to be weighed and the mass is to be determined by the difference. The

amount of the original ion can be found by using the mass of the precipitate and its known

composition.

Titration is the determination a solution’s concentration, with respect to water that has a

pH of 7. The first step of titration is to ensure that each sample has an equivalent weight in

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grams, then putting drops of Calcon indicator into the solution and titrating with Na2H2EDTA.

The color of the sample will then change from pink to blue. Specific calculations are also

necessary to complete the titration process successfully.

This paper utilizes all of these different types of analyses in studying samples of seawater

in order to find out which of the major ions were present and the levels at which they were

present. Qualitative analysis was used to verify whether ammonium, barium, calcium, potassium,

and sodium were present in the samples. In addition, the gravimetric analysis was used to

measure sulfate and halide ions. (A few examples for the latter are bromide, chloride, fluoride,

and iodide.) Finally, calcium and magnesium were measured using titration.

METHODOLOGY

To execute this experiment, samples of seawater were first collected by our TA Jake.

Four locations were selected from the areas surrounding the Galveston Island, and from each

location both a ‘shallow’ and ‘deep’ samples were collected. Shallow samples were taken from

just below the surface of the water, and deep samples were taken from just below the water’s

surface. For each sample, the salinity of the water was measured. The sample water my group

received was from Deer Island (locations shown in Figure 1). The weather that day was 69

degrees with overcast. The salinity that was at 21.5 ppt and had south winds of 27mph.

Fig. 1 Collection Site

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Before beginning any of the analyses, a standard solution was prepared by placing seven drops of

0.1 M BaCl2, 0.1 M CaCl2, 0.1 M KNO3, and 0.1 M NaCl into test tube. The explanation of how

solution and the seawater sample were comparatively analyzed follows below.

QUALITATIVE ANALYSIS FOR CATIONS AND ANIONS IN SEAWATER

To establish the precipitation of Ba2+ and Ca2+, one drop of 6 M NH3 was added to the

solution, followed by six drops of 3 M (NH4)2CO3 while stirring constantly. The solution was let

to sit for five minutes so that the precipitation of the BaCO3 and CaCO3 salts could occur, and

was then placed in the centrifuge for another five minutes. 3 M (NH4)2CO3 were then added one

drop at a time to test for total precipitation. If the precipitation was not completed, the solution

was placed in the centrifuge again. Once the precipitation was completed, the supernatant

solution was decanted into test tube and preserved for later testing of K+ and Na+.

To test for Ba2+, the precipitate from the previous step was washed and centrifuged. The

wash water was then discarded into the waste container. Three drops of 6 M acetic acid

(CH3COOH) were added, followed by 0.5 mL of water. The solution was then buffered by the

addition of three drops of 3 M ammonium acetate (NH4CH3COO). Next, two drops of 1 M

potassium chromate (K2CrO4) were added, followed by five minutes of mixing and centrifuging.

If the solution did not appear yellow after five minutes, more 1 M K2CrO4 was added very

gradually until the solution turned yellow, taking care to avoid the persistence of a highly yellow

color. The purpose for this exercise was to detect the presence of BaCrO4, which was indicated

by the yellow color. If a definite yellow color persisted, adding 1 M K2CrO4 in excess of twelve

drops was avoided. Any precipitate was then washed. The supernatant solution was decanted into

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a clean 6 x 50 mm test tube and then saved for future testing (of Ca2+). After the yellow solution

had been decanted, three drops of 6 M HCl were added to dissolve the yellow solid residue of

BaCrO4 in order to confirm the test for Ba2+. This was followed by the addition of 0.5 mL of 0.1

M Na2SO4. If necessary, the test tube was then centrifuged to determine whether the precipitate

was white in the yellow solution; a white precipitate of BaSO4 proves the presence of Ba2+.

To test for Ca2+, five drops of 1 M potassium oxalate (K2C2O4) were added to the yellow

solution from which the BaCrO4 was removed. To make the mixture basic, one drop of 6 M NH3

was added. Up to ten minutes of wait time were then allowed for any precipitate to appear. Any

presence of Ca2+ was shown by the development of a white precipitate of calcium oxalate,

CaC2O4. The presence of Ca2+ was reconfirmed by first decanting the yellow solution and then

adding two drops of 6 M HCl and 0.5 mL of water so that the solid residue would be dissolved.

Next, one drop of 1 M K2C2O4 was added, and the solution was made to be basic with 6 M NH3.

This would cause the precipitate of CaC2O4 to reappear, indicated by the presence of small flakes

in the solution.

After testing for Ba2+ and Ca2+, the flame test was utilized in order to test for Na+ and K+.

A wire loop was submerged in a solution that contained sodium or potassium ions, and then

heated with a Bunsen burner flame. The purpose of this process was to dissociate the salts into

neutral atoms. When heated to high temperatures, the electrons of atoms are put into an agitated

state, and as they return to their normal state, the light of certain colors are produced. Bright

yellow is indicative of the presence of sodium; a dull violet color indicates that potassium is

present. When both colors occur simultaneously, the sodium emission obfuscates the potassium

emission’s color, so in this experiment a cobalt glass filter was used to filter out the yellow light

for the violet light to be seen. Figure 5 shows the flame test being performed.

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Any tests for NH4+ were carried out on the original samples or solutions, as the

ammonium salts produced were used as reagents throughout the procedures that studied other

cations. Twelve drops of 0.1 M NH4Cl and twenty-one drops of water were added to a large well

on the micro-drop tray, followed by seven drops of 6 M NaOH. The reasoning behind the

addition of twenty-one water drops is to achieve the proper dilution ratio. Next, a piece of red

litmus paper was fixed over the well and then covered with a watch glass. The release of NH3

gas would cause an even blue color to appear, showing the presence of the ammonium ion. The

test was repeated using thirty-three drops of seawater—instead of the 0.1 M NH4Cl—as well as

deionized water.

QUANTITATIVE GRAVIMETRIC ANALYSIS FOR ANIONS AND CATIONS IN

SEAWATER

Before beginning the gravimetric analysis, four test tubes were labeled with a black

marking pen. The test tubes were placed in a 150 mL beaker and dried in an oven set to 110

degrees Celsius for thirty to sixty minutes. After drying, the test tubes were cooled in a

desiccator and then each tube was weighed to the nearest milligram. 1 mL of seawater was then

added and the tubes were then re-weighed. Next, the water was acidified using two drops of 6 M

Fig. 5 Flame Test

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nitric acid (HNO3) so that the OH- and HCO ions could be neutralized. The halide ions (Cl, Br,

and I) were precipitated by the addition of 6 mL of 0.1 M AgNO3 to the seawater sample, which

caused the formation of a white precipitate of silver chloride (AgCl). The solution was then

stirred with a stirring rod, ensuring that no precipitate was lost. The test tube was then

centrifuged for between one and two minutes to settle the precipitate. If all of the precipitate had

not collected at the bottom of the test tube, the sample was centrifuged again until all of the

precipitate was at the bottom of the tube, approximately one minute. Figure 6 shows an even

distribution of the sample about to be centrifuged.

An additional one to two drops of 0.1 M AgNO3 was added for any clear solution above

the precipitate. If more white precipitate formed, this indicated that the precipitation was

incomplete, and five more drops of AgNO3 were added, the solution was stirred, and then re-

centrifuged. The lack of white precipitate indicated that the precipitation was complete, so it was

Fig. 6 Centrifuge

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then washed so that any remaining Ag+ ion would be freed. This was done by decanting the

supernatant solution with a polyethylene pipet. When pouring off the supernatant solution, the

precipitate stayed in the bottom of the tube. Had it not, it would have indicated that it had not

been centrifuged for the right amount of time. Next, 6 mL of deionized water were added to the

precipitate, which was then stirred with a glass rod and re-centrifuged. The supernatant solution

was again drawn off, and this step of the experiment was performed once more with a new

amount of deionized water. As much water that could be removed without losing any precipitate

was then taken out of the precipitate. The test tube containing the precipitate of silver halide was

then dried overnight. After the samples had dried, the sample was let to cool in a desiccator until

reaching room temperature. The masses of all four test tubes containing precipitates were then

recorded.

GRAVIMETRIC ANALYSIS OF S AS BARIUM SULFATE

The process of gravimetric analysis was begun by doing the calculation for the ionic

equation for the reaction of barium ions with sulfate ions to form insoluble barium sulfate

(BaSO4). Next, the concentration of sulfate in seawater in units of mg/L and mmol/L was

calculated. Using this estimate, the volume of seawater required to produce 100 mg of barium

sulfate (BaSO4), as well as the volume of 0.3 M BaCl2 solution required to precipitate the entire

sulfate ion (+ 20% extra Ba2+) was calculated.

After verifying the accuracy of the calculations, the process of gravimetric analysis for

BaSO4 was begun. The first steps of the gravimetric analysis were to dry, cool, and weigh the

test tubes, recording the weights to the nearest milligram. An exact volume of seawater was then

put in the test tubes. Each sample was acidified with 0.5 mL of 6 M HNO3 and heated to remove

Hansen 10

any carbonate ions as CO2. After heating, the exact volume of 0.3 M BaCl2 was added to each

acidified seawater sample. Next, the precipitate was centrifuged and washed two times with 5

mL portions of deionized water. For the final wash liquid, it was carefully drawn off and the test

tubes containing the BaSO4 precipitate were placed in a 150 mL beaker and let to dry overnight

in an oven. The next day, after the tubes were cooled a desiccator, they were then re-weighed and

the mass of each precipitate recorded to the nearest milligram. From the mass of barium sulfate,

the moles and mass of sulfur in the sample were calculated. The concentration of S that was in

the original seawater sample in units of mg/Liter of sulfur was then calculated using the mass of

sulfur and volume of the water sample.

EDTA TITRATION OF CALCIUM AND MAGNESIUM IN SEAWATER IN THE

PRESENCRE OF MAGNESIUM USING CALCON INDICATOR

For the EDTA titration analysis, it was imperative that a waste collection container and

safety goggles were used at all times. A 1-mL microburet was filled with 0.0500 M Na2H2EDTA

and weighed to the nearest milligram of 4.1 g of seawater in a 25 mL Erlenmeyer flask.

Approximately 10 mL of 0.15 M NaOH were then added, and a small stirring bar was placed in

the flask. Next, the flask was placed with a sheet of white paper underneath in order to make the

color change that would occur more visible. Five drops of 0.4% Calcon indicator were then

added, followed by an immediate titration with the 0.0500 M Na2H2EDTA from the 1-mL

microburet. At the end, the visible color change was pink and then to blue. Figure 7 displays both

the before- and after-image of the color change. These steps were repeated three times, and this

entire process was done for calcium, magnesium, and calcium plus magnesium, for both the

shallow and deep locations at each sample collection location.

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The Ca2+ for the samples titrated was calculated, followed by a calculation to determine the

average of the concentrations, measured in units mmol/L and mg/L. The totals of Mg2+ and

Ca2+ in units of mmol/L were obtained using the Calmagite indicator in a different titration. The

Mg2+ ion concentration was obtained in units of mmol/L by the difference of (Mg2+ + Ca2+)

mmol/L – (Ca2+) mmol/L .

RESULTS

The following tables below contain the data collected during this lab. Table 1 shows

where the presences of ammonium, barium, calcium, potassium, and sodium were located. A

plus sign (+) indicates that there was a presence and a minus sign (-) represents the ion’s

absence. Table 2 contains the gravimetric analysis data showing the weight change and ppm of

AgCl and BaSO4-. Table 3 contains examples of titration, by working out and showing how the

results gathered were found.

Fig. 7 Titration

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Barium Calcium Potassium Sodium Ammonium Group 1 Shallow

+ + + + +

Group 1 Deep + + + + - Group 2 Shallow

+ + + + +

Group 2 Deep + + + + + Group 3 Shallow

- - - - +

Group 3 Deep + + + + + Group 4 Shallow

+ + + + -

Group 4 Deep + + + + -

Weight Change of AgCl ppm for AgCl

.01 11.6235ppm

.02 23.2799ppm

.03 30.9198ppm

Weight Change of BaSO4- ppm for BaSO4-

.03 19,294ppm

.06 38,588.24ppm

.1 64,471ppm

.005 3,233ppm

.037 5,819ppm

.023 14,872ppm

.018 11,640ppm

.02 ,932ppm

.017 10,992.3ppm12

Table 1 Qualitative Analysis of Barium, Calcium, Potassium, Sodium, and Ammonium

Table 2 Gravimetric analysis data/results

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20 drops = !"  !"!"  !"#$%

= .8 mL

.8 mL X !  !!"""

𝑚𝑙 = .0008 L

.!"  !"#  !"#$!!

X.!!!"  !!

= .00004 mol EDTA

.00004 mol X !"""  !!"#!  !"#

= .04 mmol EDTA

.04 EDTA = .04 mmol Ca2+

.!"  !"#  !"#!$%&.!!"!  !"#$#%"&

= 10 mmol/L Ca2+

!".!!  !"#$%&'!

X .00004 mol = .001604 g Ca2+

.001604 g Ca2+ = !"""  !"  !"#$%&'!  !  

= !.!"#$  !"  !"#$%&'.!!"!

= 401 mg/L – 10= 301 mg/Mg

30 drops = !  !"!"  !"#$%

= 1.2 mL

1.2 mL X !  !!"""

𝑚𝑙 = .0012 L

.!"  !"#  !"#$!!

X.!!"#  !!

= .00006 mol EDTA

.00006 mol X !"""  !!"#!  !"#

= .06 mmol EDTA

.04 EDTA = .04 mmol Ca2+

..!"  !"#  !"#!$%&.!!"!  !"#$#%"&

= 10 mmol/L Ca2+

!".!!  !"#$%&'!

X .00006 mol = .002406 g Ca2+

.002406 g Ca2+ = !"""  !"  !"#$%&'!  !  

= !.!"#!"  !"#$%&'.!!"!

= 601.5 mg/L – 15= 586.5 mg/Mg

Table 3 Titration

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23 drops = !  !"!"  !"#$%

= 0.92 mL

0.92 mL X !  !!"""

𝑚𝑙 = .00092 L

.!"  !"#  !"#$!!

X.!!!"#  !!

= .000046mol EDTA

.000046 mol X !"""  !!"#!  !"#

= .046 mmol EDTA

.046 EDTA = .046 mmol Ca2+

..!"#  !"#  !"#!$%&.!!"!  !"#$#%"&

= 11.5 mmol/L Ca2+

!".!!  !"#$%&'!

X .000046 mol = .0018446 g Ca2+

.0018446g Ca2+ = !"""  !"  !"#$%&'!  !  

= !.!""#!"  !"#$%&'.!!"!

= 461.15 mg/L -11.5= 449.65mg/Mg

33 drops = !  !"!"  !"#$%

=1.32 mL

1.32 mL X !  !!"""

𝑚𝑙 = .00132L

.!"  !"#  !"#$!!

X.!!"#$  !!

= .000066mol EDTA

.000066 mol X !"""  !!"#!  !"#

= .066 mmol EDTA

.066 EDTA = .066 mmol Ca2+

..!""  !"#  !"#!$%&.!!"!  !"#$#%"&

= 16.5 mmol/L Ca2+

!".!!  !"#$%&'!

X .000066mol = .0026466 g Ca2+

.0026466g Ca2+ = !"""  !"  !"#$%&'!  !  

= !.!"!!!"  !"#$%&'.!!"!

= 661.65 mg/L - 16.5= 645.15mg/Mg

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31 drops = !  !"!"  !"#$%

=1.24 mL

1.24 mL X !  !!"""

𝑚𝑙 = .00124L

.!"  !"#  !"#$!!

X.!!"#$  !!

= .000062mol EDTA

.000062 mol X !"""  !!"#!  !"#

= .062 mmol EDTA

.062 EDTA = .062 mmol Ca2+

.!"#  !"#  !"#!$%&.!!"!  !"#$#%"&

= 16.5 mmol/L Ca2+

!".!!  !"#$%&'!

X .000062mol = .0024862 g Ca2+

.0024862g Ca2+ = !"""  !"  !"#$%&'!  !  

= !.!"#!"  !"#$%&'.!!"!

= 621.55mg/L - 15.5= 606.05mg/Mg

Discussion

For the qualitative experiment, according to the data from trial 1, ammonium showed up

the least in seawater. But for barium, calcium, potassium, and sodium did not show up at all in

shallow water sample, but it was in all of the other samples. From trial 2 the data collected

calcium and ammonium was not present in shallow water. Trial 3 had most of their data entered,

except their data was almost opposite of class one. Barium showed up the least, mostly in

shallow water. Potassium, sodium, and ammonium all showed up equally in shallow and deep

locations. Factors that could have contributed to these results could have been from a numerous

of factors. The biggest factor could have been weather climate from previous day.

The gravimetric experiment had a huge amount of difference in its data. There were a

very small numbers of sulfates and halides present in trial 1 and a very large amount of sulfate

and halides present in trial 2. Overall, the shallow samples had less amounts of sulfate present.

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For the halides, the deep samples were more present. While the shallow samples had less sulfates

present.

The titration part of the experiment was very time consuming and tedious. It was also the

more complicated out of the other test, due to the fact of trying to understand how to do the

calculations. Magnesium appeared to be more prevalent in deep water. For trial 1, more

magnesium is present in shallow water. Trial 2 has more magnesium present in deep water and

same for class three. Calcium showed up more an equal amount of times in shallow water in all

trials.

There were a few significant errors in this experiment. One is that the oil spill prevented

classes from sampling water on the boats. Instead the TA’s all went and collected samples

around the island. This could significantly change the data in the experiment because samples

were collected in completely different areas of the island. Another error was time management

and productiveness from the groups during lab time.

Research regarding the differences between seawater samples collected at different

depths was inconclusive. The lack of widespread research into this specific area of seawater

analysis means that the kind of research done in this paper is all the more important. The general

scientific consensus does seem to be that chloride and sodium are the most common ions found

in seawater, followed by calcium, magnesium, potassium, and sulfate.

In conclusion, the ocean is comprised of many different elements that change and affect it

every day. From human waste and detritus that is discarded into the oceans every day, among the

long-term effects of other forms of pollution, man’s influence on seawater is only one of the

factors that change the ocean’s makeup. Rivers that run off into oceans, natural animal waste,

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changing global temperatures, and other natural causes also contribute to changes in seawater’s

ionic makeup at any given location at any given time. This experiment sought to quantify and

qualify the ionic makeup of the Earth’s oceans in order to learn more about the composition of

seawater and what affects it, using historical and scientific perspectives to do so.

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http://www.marinebio.net/marinescience/02ocean/swcomposition.htm

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http://www.seafriends.org.nz/oceano/seawater.htm

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Chemistry of Seawater. (n.d.). Seawater Salt Concentrations. Chemistry of Seawater. Retrieved

from http://oceanplasma.org/documents/chemistry.html#6_major_ions

Duedall,  I.  W.,  &  Weyl,  P.  K.  (1967).  The  partial  equivalent  volumes  of  salts  in  seawater.  

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G. P. Hildebrand and C. N. Reilley, “New Indicator for Complexometric Titration of Calcium in

Presence of Magnesium, Anal. Chem. 1957, 29(2), 258-264.

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http://www.chem.tamu.edu/class/majors/tutorialnotefiles/gravimetric.htm

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Retrieved from http://people.chem.duke.edu/~jds/cruise_chem/oceans/seawater1.html