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GEOL 6660, 3 Credits, Environmental Field and Laboratory Geochemistry,
Prerequisites: Chemistry 2025 and Geology 4658 or consent of department. Field procedures, analytical laboratory procedures, and geochemical data interpretation are covered and used in an environmental geochemical study completed by each student. One hour of lecture and 6 hours of laboratory and/or field.

GEOL 6660 Syllabus - Spring 2003

Class time: To be arranged

Text: Xerox copies of procedures used in chemical analyses and journal articles of field and laboratory studies will be provided by the instructor.

Instructor: Ronald K. Stoessell, Room 1034, 280-6795,

The first half of the course provides a practical background in collecting aqueous field samples and analyzing them with appropriate field procedures (e.g., pH, Eh, dissolved O2, alkalinity, aqueous sulfide) and laboratory procedures for measuring major cations and anions and some trace constituents using the liquid ion chromatograph, UV-VIS spectroscopy, inductively-coupled plasma emmission spectroscopy, and total element analyzers. The computer program MINTEQA2 is used to calculate saturation indices for various solids from the solution analyses. The second half of the course is used to complete a research project such as determining the changes in nutrient concentrations in Mississippi River water flowing across the 10,000 acre Davis Pond Diversion area southwest of New Orleans or the changes in nutrient concentrations in waste water flowing across the artificial wetlands in the Mandeville Waste Water Treatment Facility, the distribution of aqueous lead in monitoring wells at the Delatte Battery site near Pontchatoula, and the ground-water chemistry of the salt-water intrusion in either the Northshore aquifers near Lacombe or in the Baton Rouge aquifers.

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The course outline is given below.

Lecture, Laboratory, and Field Meetings will be held in which the field and laboratory meetings will last longer than the lectures and will generally count more than a single lecture. For example, field meetings on the weekend will count for 2 (1&1/4 hr) class meetings because they will require 8 hours. Laboratory work will be scheduled as a substitute for lecture in which the students will work with the instructor, with their fellow students, or by themselves in the Geochemistry Laboratory.


#1 & #2 Lectures on concentration units, field sampling procedures, preserving samples, and doing field measurements of pH, Eh, conductivity, alkalinity, dissolved oxygen, aqueous sulfide, and hardness.

#3 Practice in the laboratory doing the field measurements of pH, Eh, conductivity, alkalinity, dissolved oxygen, hydrogen sulfide, and hardness.

#4 Meet on a weekend to take aqueous samples from field sites and measure pH, Eh, conductivity, alkalinity, dissolved oxygen, aqueous sulfide, and hardness. These samples will be used for training in the laboratory on the analytical equipment.

#5 & #6 Practice making up standards and an overview of the use of the total element aqueous analyzer, liquid ion chromatograph for measuring major anions, cations, and nutrients; the UV-VIS spectrograph for measuring aqueous silica; and the use of the inductively coupled plasma emission spectrograph for measuring aqueous trace metals.

#7 Practice laboratory use of the total element analyzer for aqueous C and N in field samples.

#8 Practice laboratory use of liquid ion chromatograph to analyze major anions and nutrient anions in field samples.

#9 Pracice laboratory use of liquid ion chromatograph to analyze major cations and ammonium in field samples.

#10 Practice laboratory use of UV-Vis spectrograph to measure aqueous silica in field samples.

#11 Practice laboratory use of the ICP to measure aqueous metals in field samples.

#12 Set-up MINTEQA2 to compute mineral saturation indices for use in interpreting the data for the individual projects.

#13 Meet to discuss student field projects and plan sampling trips.

#14 & #15 Instructor will go with the students individually or in groups to the field to take samples for projects and to measure pH, Eh, dissolved oxygen, alkalinity, and aqueous sulfide (if present)

#16 Laboratory analyses of field project samples - total element analyzer - aqueous C and N

#17 Laboratory analyses of field project samples - liquid ion chromatograph - anions and nutrient anions

#18 Laboratory analyses of field project samples - liquid ion chromatograph - cations and ammonium

#19 Laboratory analyses of field project samples - UV-Vis spectrograph - aqueous silica

#20 Laboratory analyses of field project samples - ICP - metal cations

#21 Laboratory analyses of field project samples - total organic C and total organic N

Students can schedule additional time to finish their analyses.

Last week of class - Give a brief summary to the class of your project and your results, turn in short paper (minimum 5 pages, double-spaced, with an Introduction, Procedure, Results, Summary, and Bibliography) on the project and a copy of your notebook showing field and laboratory procedures. The course grade will be based on the student's work in the class project (paper) and in learning the field and laboratory procedures (notebook).

Field and Laboratory Environmental Geochemistry Topics

  1. Concentration Units
  2. Laboratory Acids
  3. Sample Collections
  4. Field Measurements
    1. pH
    2. Eh
    3. Conductivity
    4. Alkalinity
    5. Dissolved Oxygen
    6. Aqueous Sulfide
    7. Hardness
  5. Laboratory Analyses
    1. Liquid Ion Chromatography in Anion and Cation Analyses
    2. UV Visible Spectroscopy in Silica Analysis
    3. ICP Emission Spectroscopy in Cation Analyses
    4. Total Aqueous Element Analyzer for Carbon and Nitrogen
  6. Computer Programs
    1. MINTEQA2
  7. Examples of Research Projects
    1. Davis Pond Freshwater Diversion
    2. Mandeville Artificial Wetlands
    3. Nutrient Removal with Aluminum Stearate
    4. Salt Water Intrusion in the Big Branche Aquifer in Lacombe
  8. Example References for Research Projects

Concentration Units

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In solutions, the fundamental concentration unit is the mole fraction Xi; in which for j components, the ith mole fraction is

Xi = ni/(n1 + n2 + ...nj),

where the number of moles n of a component is equal to the mass of the component divided by its molecular weight. Note that a substance composed of only one substance will have a mole fraction of unity for that substance. In an aqueous solution, the mole fraction of water, the solvent, is always near unity. In solids that are nearly pure phases, e.g., limestone, the mole fraction of the dominant component, e.g., calcite, will be near unity. In general, only the solutes (minor components) in a liquid solution and gas components in a gas phase will have mole fractions that are significantly different from unity.

Other more commonly used concentration units for the solutes are given below:

per cent, parts per hundred
ppt, parts per thousand
ppm, parts per million
ppb, parts per billion
ppt, parts per trillion

These units are often assumed to be in terms of mass (weight) unless otherwise stated. Note that one mg/kg is one ppm where a mg is 1/1000th of a gram and kg is 1000 grams. Similarly, one ug/g is also one ppm where a ug is 1/1,000,000th of a gram. One ug/kg is a ppb.

Because one liter of water weighs approximately one kg, mg/liter units of solution are nearly equal to ppm units. The precise equivalence is obtained by dividing by the density p:

ppmi = [mgi/liter]/p

where the solution density is in grams/cm3. Older texts will substitute specific gravity for density in the above equation. The specific gravity is the ratio of the solution density to that of the density of pure water at 4oC. Since the density of pure water at 4oC is 1 gram/cm3, the specific gravity is equal to the solution density when expressed in metric units of g/cm3.

Note that a ml (1/1,000th of a liter) is equal to a cm3. The units of ppt and ppm are commonly used for the concentrations of solutes in aqueous solutions such as sea water. Trace components are represented in the ppb and ppt range.

The units of g/cm3 are used in diffusion. This unit is nearly equivalent to weight fraction for an aqueous solution since one cm3 of water approximates one gram.

Other concentrations units are described in terms of the number of moles of a components.

The molarity of i, Mi, is the moles of component i per liter (1000 cm3 or 1000 ml) of solution.

The formality of i, fi, is the moles of component i per kg of solution.

The molality of i, mi, is the moles of component i per kg of solvent (e.g., water). Molality is used only in aqueous solutions. In making thermodynamic calculations, the units used are molality for all aqueous components.

The equivalents of i, ei, is generally the moles of charge of i in molarity, formality, or molality units. If the equivalents refer to a particular reaction, than it refers to the equivalents of charge needed for the reaction to be completed in terms of molarity, formality, or molality units./P>

Multiplication of the above mole concentration units by one thousand will convert them to milli units, e.g., mMi, mfi, mmi, and mei.

Note that molarity is defined in terms of volume, a liter of solution. Any concentration unit involving volume will change with temperature and pressure, because volume is a function of those parameters. For this reason, geochemists prefer not to use molarity units in aqueous solutions; however, chemists prefer molarity units.

To convert from ppm to formality units

fi = ppmi/(1000 Mwi) where Mwi is the molecular weight of i.

To convert from ppm to molality units

mi = [ppmi/(1000 Mwi)] [1/(1 - tds/1,000,000)]

where tds is the total dissolved solids in ppm in the solution.

To convert from ppm to molarity units

Mi = [ppmi/(1000 Mwi)] p

The concentration units are not absolute amounts of a component in the solution but are relative amounts, having been normalized to the total mass or volume of the solution or solvent. Hence, changing the amount of a solution does not change the concentration of any of the components. Concentration units, as well as temperature T and pressure P, are intensive parameters that do not depend upon the total mass of a solution. Volume is an example of an extensive parameter, one that depends upon the total mass of the solution.

Laboratory Acids

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Laboratory reagent-grade acids and bases are commonly diluted for use in chemical analyses. Concentrated strengths (per liter of solution) and 25oC liquid densities (g/cm3), are given below to help in doing dilutions.

Hydrochloric Acid, HCl, 1.18-1.19 g/cm3, 12N (12M)
Nitric Acid, HNO3, 1.41 g/cm3, 15.5N (15.5M)
Sulfuric Acid, H2SO4, 1.84 g/cm3, 36N (18M)
Hydrofluoric Acid, HF, 1.15 g/cm3, 26.5N (26.5M)
Acetic Acid, CH3COOH, 1.05 g/cm3, 17.4N (17.4M)
Ammonium Hydroxide, NH4(OH), 0.90 g/cm3, 14.3N (14.3M)

Sample Collections

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Unfiltered samples should be taken to measure the temperature, conductivity, pH, Eh, and dissolved oxygen of the sample. Four clean plastic beakers with 50 ml sample each can be collected on site and used to complete these measurements within a few minutes of collection. Temperature is usually measured simultanously with the conductivity, pH, and dissolved oxygen probes. If the Winkler titration is being done for dissolved oxygen, replace one of the beakers with a clean 100 ml glass bottle and fill it to the brim with unfiltered sample before inserting the glass stopper. The Winkler titration should be done as soon as possible. If aqueous sulfide is preesnt, the sample will have a rotten egg smell. Collect a 125 ml sample in a plastic bottle for subsequent analyses of aqueous sulfide within a few hours of collection.

Approximately 120 ml of sample needs to be collected and filtered through a 0.45 micron filter ot Whatman #42 filter. Twenty ml of the filtered sample will be needed for an alkalinity titration. If water hardness is to be measured, use 20 ml of the filtered sample for the titration. One 35 ml aliquot of the filtered sample should be stored in a clean plastic bottle and acidified with HCl to a pH of about 2 to keep metals soluble and to analyze later in the laboratory. Another 35 ml aliquot of filtered non-acidified sample should be stored in a clean plastic bottle for laboratory anion analyses and for aqueous silica. The two filtered aliquots should be refrigerated to inhibit biological activity until analyzed in the laboratory.

Field Measurements


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The pH is defined as the -log aH+ in an aqueous solution. In field samples, the pH must be measured immediately, before gas exchange takes place with the atmosphere. The field temperature of the sample at the time of pH measurement should be recorded to subsequently compute the the total hydrogen concentration in the sample. This allows the pH to be back-computed for the sample at the in situ temperature (assuming it is known) and then the computation of gas saturations and mineral saturation indices.

The definition of pH generally relates to our concept of acids and bases in aqueous solutions, the Bronsted concept in which the acid donates a proton and the base accepts a proton. An acidic solution has a higher concentration of hydrogen ions H+ than of OH- ions. These two ions are related through the disassociation reaction with water H2O. A broader definition of an acid or a base comes from Lewis in which the acid can accept an electron pair (has an empty outer shell orbital) to share in a covalent bond from a base which has an unshared electron pair. In this other ions besides H+ count as acids and other ions besides OH-.

H2O = H+ + OH- where KH2O = aH+aOH-/aH2O is the mass-action law.

The a is called an activity and is the product of the concentration (molality which is moles per kg of H2O) times a fudge factor called the activity coefficient.

At 25oC and 1 bar (near atmospheric pressure), KH2O is 10-14.

The activity of water in most aqueous solutions (except brines) can be assumed to be unity, leading to

10-14 = aH+aOH- for the mass-action law.

From the mass-action law, in a neutral solution the H+ and OH- activities must each be 10-7. This neutral solution is neither acid nor basic and has a pH of 7 (-log aH+ = 7). Acidic solutions have a pH lower than 7 and basic solutions have a pH greater than 7.

The pH can be measured with a glass electrode in combination with a reference electrode. In practice, the electrode should be calibrated within a few minutes of the field measurement by using two pH buffer solutions, usually pH 4 and pH 7. The pH is measured in the unfilterd solution because filtering causes gas exchange which will change the pH. Excessive stirring should be avoided because that too will cause gas exchange. Since gas exchange with the atmosphere cannot be completely avoided, it is important to measure the pH within a few minutes of taking the sample. Even if the pH is not stable, it will usually rapidly change as the electrode equilibrates and then change more slowly with gas exchange. The trick is to pick the pH corresponding to the electrode equilibration and record that value.


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Eh is the voltage produced by connecting the aqueous solution through a platinum electrode to the standard hydrogen electrode (SHE). The voltage is defined relative to oxidation of 0.5H2 to H+ in SHE. The voltage is the sum of two half cell reactions: the reaction in SHE and the aqueous solution reaction. A positive voltage means that electrons flow from SHE to the aqueous solution. A negative voltage means that electrons are flowing to SHE for a reversal of the half cell reactions. If electrons flow towards the aqueous solution than the solution is deficient in electrons (relative to SHE) so it is called an oxidizing solution. If electrons flow away from the aqueous solution than the solution has excess electrons (relative to SHE) so it is called a reducing solution. Hence, a positive Eh corresponds to oxidizing conditions and a negative value corresponds to reducing conditions in the aqueous solution.

SHE consists of hydrogen gas being bubbled under a pressure of 1 bar through an HCl aqueous solution in which H+ has unit activity, i.e., a pH of 0. By convention, the voltage of SHE is set to zero, so the measured cell voltage (at zero current flow) is assigned to the aqueous solution connected through it by the salt bridge. The platinum electrode contributes no voltage but is simply a conductor for the flow of electrons.

In practice, the SHE half cell is too hard to maintain for making measurements. Instead, the cell voltage is measured across an external circuit between the platinum electrode in the sample solution and a reference electrode, connected by a salt bridge to the solution. The Ehref of the reference electrode relative to oxidation in SHE has to be added to the measured voltage EMF to compute Eh for the solution. Ehref is 0.2415 volts at 25oC and 1 bar for the saturated (4.16 M KCl) calomel electrode, corresponding to the reduction of 0.5Hg2Cl2 to Hgo + Cl-, relative to oxidation in the standard hydrogen electrode. Ehref is 0.2802 volts at 25oC and 1 bar for the 1 M KCl calomel electrode. The saturated Ag-AgCl reference electrode has a voltage of 0.1986 volts for the reduction of AgCl to Ag(s) + Cl- at 25oC and 1 bar.

Eh = EMFmeasured + Ehref

The Eh can be shown to be related to the pe by the following relationship.

pe = FEh/(2.303RT)>

where F is the Faraday constant (96,485 J/V/mole), R is the gas constant (8.314 J/mole) and T is the temperature in degrees Kelvin (298.15oK = 25oC or oK = 273.15 + oC).

The pe is defined as the -log ae- where ae- represents an aqueous electron activity in solution. In reality, free aqueous electrons do not exist in a solution; however, we can still use the concept. The pe is a measure of how oxidizing or reducing a solution can be. An oxidizing solution can be considered to have a deficiency in aqueous electrons, A reducing solution can be considered to have an excess of aqueous electrons, promoting the acceptance of electrons by ions being reduced. Hence an oxidizing solution has a low ae-(positive pe) and a reducing solution has a high ae- (negative pe).

Measuring the Eh with a platinum electrode is not easy, because the surface of the electrode is easily poisoned upon contact with the solution. Theoretically, the measured Eh should represent the availability of electrons for all oxidation-reduction reactions in aqueous solution. But, because the oxidation-reduction reactions are often not in equilibrium, the reaction kinetics produce a different voltage for each reaction. If they were all all at equilibrium, the voltage would be the same for each reaction because they are in the same solution. One particular (unknown) oxidation-reduction reaction may be responsible for the measured voltage, making the voltage hard to interpret for geochemical purposes.


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The electrical conductivity of an aqueous sample is the reciprocal of the resistivity and must not be confused with hydraulic conductivity. Conductivity increases with increasing salt concentration, with increasing temperature, and differs for different salts and varies with the geometry of the conductivity cell. In order to rule out the effect of temperature and cell geometry, conductivity is often reported as specific conductivity. This is the reciprocal of the resistance in ohms measured between opposite faces of a centimeter cube of an aqueous solution at a specified temperature, usually 25oC. The units are in reciprocal ohms, Siemens or mhos, per cm at a specific temperature. Seawater has a specific electrical conductivity of about 0.050 mhos/cm (or 0.050 Siemens/cm) at 25oC and distilled water approaches zero mhos/cm.

In general, for a given dissolved salt, the electrical conductivity increases linearly with an increase in salt concentration in dilute waters. At seawater concentrations, the increase becomes less than linear. Nevertheless, electrical specific conductivity is often used to predict total dissolved salts of a sample or its salinity by assuming the salts are in the same relative concentrations as in sea water and using an algorithm that fits conductivity to a seawater-type solution.

The measurement is done with a probe that gives the specific conductivity and temperature of the sample and also the estimated specific conductivity at 25oC. The probe does not usually need to be calibrated in the field. The probe units are usually micro mhos/cm, giving seawater a value of 50,000 to 55,000 on this scale, depending on temperature.


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Alkalinity is the capacity of water to neutralize H+ ions. The alkalinity in most natural waters is due to bicarbonate and carbonate ions. High pH waters usually had a high alkalinity and low pH waters have a low alkalinity. Alkalinity is the opposite of acidity which is the capacity of water to neutralize OH- ions. Most ground waters are basic, possessing alkalinity but no acidity. Acidity is rarely measured on water samples, except for low-pH, e.g., acid mine drainage, polluted samples.

Hardness is related to alkalinity. Hardness is the sum of the concentrations of calcium and magnesium and is usually low when alkalinity is high because these cations precipitate out as carbonates. Usually when the alkalinity is high the dominant cations are sodium and potassium, i.e., the Na and K soda lakes in East Africa. Hence, hard water are high in calcium and magnesium and soft waters are high in sodium and potassium.

In practice, alkalinity is usually measured by recording the amount and strength of acid added until the solution pH is lowered to a particular value, e.g., 4.5 for inorganic carbon alkalinity. Below a pH of 4.5 all inorganic carbon is in the form of H2CO3 cannot accept more H+ ions. Acidity is measured by recording the amount and strength of base added until the solution pH is raised to a particular value, usually a pH of 8.3. Alkalinity and acidity are usually not affected by exchange of gases which change the solution pH, because gas exchange usually contains self-cancelling effects in the calculation of alkalinity and acidity. However, precipitation and dissolution reactions will change alkalinity and acidity. Alkalinities on field samples are measured in the field on filtered samples.

Important Alkalinity Reactions in Natural Waters:

Inorganic Carbon: CO32- + 2H+ = H2CO3 and HCO3- + H+ = H2CO3

Below a pH of 8.3, carbonate (CO32-) is nearly absent. This pH is near the phenolphthalein end point which is pink in color above 8.2 pH. Below a pH of 4.5, bicarbonate (HCO3-) is nearly absent which is near the pH 4.3 end point of methyl orange. Below a pH of 4.5, all the inorganic carbon is in the forme of dissolved CO2 and carbonic acid H2CO3, in which they are both called H2CO3*. Without a pH meter, inorganic carbon alkalinity is often determined by titrating acid to the methyl orange end point. However, other H+ neutralizing species may contribute to the alkalinity such as HS.

Water: OH- + H+ = H2O

Hydrogen Sulfide: HS- + H+ = H2S

Acetate: CH3COO- + H+ = CH3COOH

Alkalinity in meq/l is computed from the following formula in which the acid normality is in meq/l. To convert alkalinity to mg HCO3-/liter, multiply by the gram molecular weight of HCO3-. To convert alkalinity to mg CaCO3/liter, multiply by 1/2 the gram molecular weight of CaCO3.

Alkalinity = (ml of acid added)(normality of acid)/(ml of sample)

The pH used as the final end-point in an alkalinity titration is the pH corresponding to the point where all species capable of accepting H+ ion have been neutralized. In practice, this pH is easily found by recording the change in pH with a change in addition of acid. These end-points have rapid changes in pH with addition of acid and plot as an inflection point on a pH (y axis) versus mls of acid added (x axis). There may be more than one pH end-point for an alkalinity titration, e.g., inorganic carbon has two end points; 8.3 and 4.5. Organic acids generally have a lower pH end point than 4.5; whereas, hydrogen sulfide has a higher pH endpoint. In a mixture of different H+ accepting groups, the end points are difficult to interpret. Below the final end point, a plot of pH versus ml of acid added will show by the continuous slope that the pH is changing only by the addition of H+ ions without any neutralization of basic species.

Effects of Gas Exchange on Alkalinity

Gas exchange of the sample with the atmosphere that results in the loss of species such as HCO3- or HS- can be written as:

HCO3- + H+ => H2O + CO2 and HS- + H+ => H2S

The hydrogen ions used-up in forming the gases are from the solution. The solution is losing equivalent numbers of both hydrogen ions and species capable of accepting hydrogen ions. Hence, the solution pH will change but not the alkalinity.

Dissolved Oxygen

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Dissolved oxygen can be measured by the Winkler titration or by using an oxygen probe. We will use both methods in the field. Dissolved oxygen in saturated water is only about 10 mg/l at 20oC. The saturation concentration decreases with rising temperature and increasing salinity.

The oxygen probe contains a gold cathode electrode (where electrons are removed) and a silver anion electrode (where electrons are generated). These electrodes are immersed in a chloride containing electrolyte solution that is separated from the sample solution by a membrane that is permeable to dissolved oxygen. The dissolved oxygen is reduced at the cathode and silver is oxidized at the anode and precipitated as AgCl, producing an electric current used to measure the concentation of dissolved oxygen. The two electrodes are covered by the electrolyte solution (without any air bubbles) and separated from the solution by the membrane. Calibration is usually done with 100% dissolved oxygen saturation by placing the probe tip just above the water in a beaker or bottle. Some probes also need to be calibrated for zero dissolved oxygen using a solution supersaturated to solid sodium sulfite and a trace of cobalt chloride.

The Winkler analysis consists of making the solution basic, adding excess manganeous cations (Mn+2 to precipitate as Mn(OH)2 which absorbs dissolved oxygen as it forms a flocculant mass. The oxygen is reduced by oxidizing Mn(OH)2 to MnO(OH)2. After acidification, iodide is added which is oxidized to iodine by reducing the Mn+4 back to Mn+2. The iodine turns the solution yellow. This iodine is then titrated with thiosulfate using starch as an indicator in which the end point is from blue to colorless.

The manganese is added as a MnSO4 solution. The hydroxide solution contains the iodide and also NaN3 (an azide) to eliminate nitrite interference. Organic matter will not interfere if the thiosulfate titration is done immediately after iodine formation. Aqueous sulfide can reduce iodine so its presence will cause underestimation of dissolved oxygen. However, if aqueous sulfide is present, the rotten-egg order will be noticable. Fluoride can be added to eliminate ferrous iron interference.

To reduce one mole of O2 requires oxidation of 2 moles of Mn+2 to Mn+4 which subsequently generates 2 moles of I2 when it is reduced. These 2 moles of I2 require oxidation of 4 moles of S2O3-2 when it is reduced.

Mn+2 + 2OH- => Mn(OH)2 floc, followed by 2Mn(OH)2 + O2 => 2MnO(OH)2

Upon acidifying and adding iodide, 2MnO(OH)2 + 8H+ + 4I- => 2I2 + 2Mn+2 + 6H2O

I2 + 2S2O3-2 => S4O6-2 + 2I-1

Hence, (M of O2)(ml of sample) = [M of S2O3-2/4](ml of titrant),

or mg/l O2 = (8,000)(M of S2O3-2) (ml of titrant / ml of sample)

Aqueous Sulfide

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The analysis consists of acidifying the solution (sulfide is overestimated if not acidified), adding excess iodine, and titrating the unreacted iodine with thiosulfate. Starch is used as an indicator which turns the yellow solution to blue. The end point is from blue to colorless. (The end point is hard to see under a blue sky.)

Add HCl to acidify the solution, and then add a known (excess) amount of I2 to oxidize S-2 in the sample to S while reducing the iodine to I-1. Oxidizing one mole of S-2 requires reducing one mole of I2. The remaining I2 is then reduced by titration with S2O3-2 which is oxidized to S4O6-2. One mole of I2 is reduced by oxidizing 2 moles of S2O3-2.

S-2 + I2 => S + 2I-1, followed by the thiosulfate titration:

I2 + 2S2O3-2 => S4O6-2 + 2I-1

Hence, M of S-2 = (M of I2)(ml of I2 used to oxidize S-2)/(ml of sample).

In terms of thiosulfate,

M of S-2 = [M of S2O3-2/2] (ml of S2O3-2 titrate blank - ml of S2O3-2 titrate solution) / ml sample.

Or, mg/l S = [16,032 / ml sample] [M of S2O3-2] (ml of S2O3-2 titrate blank - ml of S2O3-2 titrate solution)


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Hardness is the sum of the Ca + Mg concentrations. It can be easily measured in the field using an EDTA titration on a filtered sample. EDTA is ethylenediaminetetraacetic acid which forms strong complexes with divalent, trivalent, and tetravalent metal ions. With these cations the complex forms as one metal cation per molecule of EDTA. It does not complex strongly with monovalent alkali cations. The general procedure is to raise the sample pH using ammonium hydroxide and ammonium chloride buffer to a pH of about 10 where metals other than Ca and Mg immediately precipitate out and then titrate the solution using Erichrome Black T as an indicator. The titration needs to be done with 15 minutes because CaCO3 will gradually precipitate out. Mg needs to be in the sample in order for the end-point to be easily determined. Often MgEDTA is added to the sample as part of the pH buffer to insure Mg is present. The end-point color change is from wine red to blue. The concentration of Ca plus Mg is the hardness of the water sample.

On a second sample the pH is raised higher to about 12 using 1N NaOH so that Mg(OH)2 precipitates out, leaving only Ca which is then titrated with EDTA using murexide as the indicator. Again the titration needs to be done quickly because CaCO3 may precipitate out. The end-point color change is from pink to purple. The Ca concentration measured in this titration is subtracted from the Ca + Mg concentration measured in the first titration to obtain the Mg concentration.

The general titration equation is:
Mmetal = MEDTA(ml of EDTA titrant)/(ml of sample)

Laboratory Analyses

Liquid Ion Chromatography in Anion and Cation Analyses

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The liquid ion chromatograph used in this course is the DIONIX 100 in the Geochemistry Laboratory within the Geology and Geophysics Department. Liquid ion chromatography utilizes an exchange column to separate anions or cations and a detector to measure a signal proportional to their concentrations. A conductivity detector is usually used for doing major anions and major cations. The sample is injected into a moving eluent (pumped at a constant rate) which carries the sample through the column and provides ions for filling the exchange sites. The anion eluent is usually a weak sodium carbonate-bicarbonate solution, and the cation eluent is a weak sulfuric acid solution. Prior to reaching the detector, the eluent signal is removed in a suppressor in which ions are exchanged for H+ and OH- ions to form H2O that doesn't give a signal on the detector.

The component arrival time is dependent upon the column length and the sample concentration. The separation of components increases with increasing column length, producing increasing arrival times at the detector. At low component concentrations, the exchange of a component on the column can be described by a Distribution Coefficient in which the concentration ratio to that in solution and that on the exchange sites is constant. Under these conditions the retardation of the component in a coloun of fixed length is independent of the component concentration, leading to a constant arrival time at the detector. However, in practice, the concentration range of the standards is often greater than the range that the Distribution Law is obeyed and the arrival time slowly decreases as the component concentration increases.

For standard concentrations, the standards usually contain a number of different components at concentrations such that their arrival times at the detector do not interfere with each other. For major anions, utilizing a conductivity detector, your instructor usually uses three standard solutions with the concentrations (given below) in which the nitrite begins to significantly decrease after a week. Hence the standards should be diluted weekly from two stock solutions with one of them (the nitrite stock solution) made weekly. The anions are listed in the order in which they arrive at the detector. A complete run takes about 7 minutes.

standard    #1      #2      #3  
solutions  mg/l    mg/l    mg/l 

fluoride   0.25,   1.25,   2.50
chloride   0.80,   4.00    8.00
nitrite    0.25,   1.25    2.50
bromide    1.00,   5.00   10.00
nitrate    1.00,   5.00   10.00
phosphate  1.00,   5.00   10.00
sulfate    1.00,   5.00   10.00

For major cations utilizing a conductivity detector, your instructor usually uses three standard solutions with the concentrations (given below) in which the ammonium begins to significantly decrease after a week. Hence the standards should be dilute weekly from two stock solutions with one of them (the ammonium stock solution) made weekly. The cations are listed in the order in which they reach the detector. A complete run takes about 14 minutes.

standard    #1      #2      #3  
solutions  mg/l    mg/l    mg/l 

lithium    0.25,   1.25,   2.50
sodium     0.80,   4.00    8.00
ammonium   0.25,   1.25    2.50
potassium  1.00,   5.00   10.00
magnesium  1.00,   5.00   10.00
calcium    1.00,   5.00   10.00
strontium  1.00,   5.00   10.00

In practice, the liquid ion chromatograph usually needs to warm up for about 2 hours to become completely stable. The standards are run once, and then the unknowns can be run for up to 6 hours without the standards shifting significantly. Much of the hastle in the analytical process is diluting the unknowns so that they fall within the standard range or so that they do not interfere with each other. Sodium and chloride are often in concentrations that significant dilution is required to prevent interference with other components. In some cases, the dilution may reduce the other components' concentrations to the point that the analysis is no longer accurate, i.e., insignificantly low, less than 0.05 mg/l of the component. Regardless, this is usually the best way to do anions! The older colorimetric methods are extremely time consuming.

UV-Visible Spectroscopy in Silica Analysis

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The UV-VIS Spectrograph used in this course is the Turner 100 in the Geochemistry Laboratory within the Geology and Geophysics Department. UV-Visible spectroscopy is used in colorimetric analyses of chemical components. Through the addition of chemicals, the chemical environment of the sample solution is adjusted such that the component forms a colored complex. The sample concentration is determined from the degree of light absorption by the colored complex at a particular wavelength. The Lambert-Beer law is used to compute the component concentration. The absorption is measured with a UV-Visible Spectrophotometer which has a monochromator to isolate light at a particular wavelength and then measures light absorption passing through a cuvette, serving as an optical absorption cell. Different cuvettes have different optical properties and even the same cuvette may have different optical transmission properties from different directions. These problems can be elimnated by using the same cuvette in the same orientation to measure the absorption of the blank, standards, and the sample unknowns.

The Lambert-Beer Law states that the decrease in transmission dI with the increasing thickness of the absorbing layer dx is proportional to the transmission I and the concentration c of the absorbing molecules. Let Io be the transmission when x = 0 and I be the transmission when x = b and let the conversion from natural logarithms to base 10 logarithms be included in the proportionality constant k. The integrated equation is

log (I/Io) = -kcb or log (Io/I) = kcb

The thickness of the absorbing layer b is held constant by the thickness of the cuvette. The base 10 adjusted proportionality constant is called the absorbancy index and the log (I/Io) is called the absorbance or optical density. The important point is that the absorbance is proportional to the concentration so that a plot of absorbance versus concentration will be linear over the range that the Lambert-Beer law is obeyed. The blank and standards are used to create a linear plot from which the unknown concentrations in the samples can be obtained.

Aqueous silica forms a yellow silicomolybdic acid complex in the presence of ammonium molybdate which can then reduced to molybdenium blue by the addition of a reducing agent. The initial formation of the yellow complex is done at about a pH of 1 to get complete formation of the complex. After 10 minutes, a reducing agent is added. The reduction is done at a very acidic pH (1.5 N sulfuric acid) to prevent reduction of the excess molybdate in solution. Oxalic acid is also added to remove phosphate interference by preventing the reduction of phosphomolybdic acid. The absorbance is measured after 3 hours at 812 mu. Three standards and a blank are usually sufficient at concentrations of 1, 2, and 3 mg/l SiO2. If the same cuvette is used with the same orientation in the spectrophotometer, the absorbance of the standards can be nearly constant as shown from repeated measurements of newly made up standards over time periods of months.

One reducing agent recommended by your instructor is a solution of metol and sodium sulphite. The procedure is described by Mullins and Riley (1955, Analytica Chimica Acta, v. 12, p. 162-176).

ICP Emission Spectroscopy in Cation Analyses

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The ICP used in this course is the Varian VISTA-MXP in the UNO Chemistry Department that is maintained and supervised by Dr. Matt Tarr. ICP refers to inductively-coupled plasma emission spectroscopy used to analyze cations. Earlier emission spectrographs utilized a direct-current plasma or DCP spectroscopy. An ICP produces a higher temperature plasma than a DCP which results in fewer interferences. An aqueous sample is pumped into the plasma and the emission spectrum measured for the sample and compared to that produced from standards at different concentrations. The standards and samples are aspirated in sequence and multiple elements are analyzed simultaneously using one previously selected emission wavelength for each element. The measured light emission from the standards in intensity counts per second at a wavelength for each element is used to make a plot of emission versus concentration. The emission at a particular wavelength has a linear response over a limited concentration range; however, the values are reproducible outside of that range, allowing the curve to be used for unknowns outside of the linear range. The wavelength is selected to be a less sensitive wavelength if the element is in higher concentration. Anions cannot be analyzed, only metals and not individual species of an element, e.g., not ammonium, only total nitrogen. This method is much quicker than the liquid ion chromatograph for cations, requiring only about a minute to run each sample for multiple metals. However, your instructor has found that the liquid ion chromatograph is less prone to interferences and has a higher accuracy.

Emission spectroscopy has largely replaced AA or atomic absorption spectroscopy for analyzing cations. AA utilzed a much lower flame temperature (acetylene and air or nitrous oxide and acetylene) into which the sample was pumped. The lower flame temperature did not eliminate ionization of easily ionized elements or chemcial reactions of refractory elements which produced interferences which required additions of chemicals to the samples to overcome. In general each element analysis required a lamp to produce the wavelength that an element absorbed at. The absorption for an element was measured as a sample was vaporized in the flame through which the light was passed. The lamp had to be changed for each element necessitating a warm-up time and much longer analytical times.

Total Aqueous Element Analyzer for Carbon and Nitrogen

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In October, 2001, the Geology and Geophysics Department received two total element analyzers for N and C in aqueous samples (Shimadzu TOC-V (6001 System) with a TON analyzer) and for N, C, H, S, and O in sediment samples (CHNSO Flash EA 1112 vy CE Elantech, Inc). The class will be taught to use the Shimadzu for aqueous analyses of total organic carbon and total organic nitrogen. The TOC uses an infrared detector and computes total organic carbon by measuring the total carbon and subtracting the inorganic carbon. Total organic carbon is determined from combustion of the sample to convert organic carbon to carbon dioxide and then detecting all the carbon dioxide from both organic and inorganic sources. Inorganic carbon is measured by adding acid to a sample to drive out the inorganic carbon dioxide and detecting that with the detector. The TON uses a chemi-luminescence detector on nitrogen released following combustion of the sample.

Computer Programs

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Computer programs are used to take the measured bulk solution composition, speciate it into free ions and complexes and compute their thermnodynamic activities. This allows saturation indices to be computed for various precipitates and to calculate the changes in amounts of each component that could occur by precipitation, absorption, degassing, and dissolution.


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MINTEQA2 is a widely used EPA program for computing speciation and saturation indices for bulk solution compositions. The DOS Fortran executable program file with the executable input file setup and a number of input data files can be downloaded with the operation manual from a number of web sites. Run the search engine GOOGLE using EPA MINTEQA2 to find a URL to download the files. One such site is Download the files into a directory and run as a DOS executable file within the Windows environment of an IBM-compatible PC. To do this, bring up the MSDOS prompt on the screen. Type in the executable file name Prodefa2 to set up the input file and then Xminteqx to run the program. Follow the directions in the manual to set up the input file.

Input your complete chemical analyses for each aqueous sample in Prodefa2 and run Xminteqx to determine mineral saturation indices. If supersaturated or undersaturated to calcite, rerun the program allowing the solution to either precipitate or dissolve calcite to see what the water-rock interaction would be. Do the same for any other minerals that you might expect to have in the field.

The program is very flexible and allows you to modify the thermodynamic data base. To do this properly, the operator needs a background in equilibrium thermodynamics, e.g., as taught in a physical chemistry or geochemical thermodynamics course.

Research Projects

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Several environmental geochemical projects are listed below as examples for your field and laboratory project.

Davis Pond Freshwater Diversion

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The Davis Pond Freshwater Diversion is a Mississippi River Diversion southwest of New Orleans of river water and suspended sediment into Upper Barataria Bay. See the URL The diversion passes through a shallow 9,200 acre holding pond (Davis Pond) with input rates of up to 10,500 cubic feet per second. The estimated hydraulic retention time is of the order of 1 to 14 days, allowing time for denitrifcation of the aerobic Mississippi River Water in the wetlands within the pond. Samples can be taken entering and exiting the holding pond and the amount of denitrification measured as a function of input rate and water temperature. The inlet samples can be taken at the highway 90 bridge that crosses the inlet into Davis Pond; however, The project necessitates access to a boat to obtain samples at the exit point into Lake Cataouatche

Mandeville Artificial Wetlands and Oxidation Lagoons

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The Mandeville Waste-Water Treatment Facility has three oxidaiton lagoons and an artificical wetlands in which the waste water moves sequentially through the lagoons and then through the wetlands prior to discharge into the Bayou Chinchuba swamp which drains into Lake Pontchartrain. The flow rates through the treatment facility are monitored at the input into the 1st oxidation lagoon and the exit from the artificial wetlands, and the normal flow is from 1 to 2 million gallons per day. The estimated hydraulic residence time based on volume calculations is from 4 to 6 days in the oxidation lagoons which are heavily aerated and mixed. The measured hydraulic retention time in the wetlands, using a bromide tracer, is from 4 to 28 hours. The flow in the wetlands is complicated by the recycling of 50% of the waste water back to the fron of the wetlands where they pass through gravel beds and are aerated by sprayers before flowing again through the wetlands. These waters do not pass through the gravel beds in their initial pass through the wetlands do to their passing under the gravel beds and then discharging into the wetlands.

One project involves the oxidation lagoons. The major processes occurring in the oxidation lagoons are the oxidation of organic carbon and the oxidation of ammonium to nitrate. Samples can be taken over time at the entrance and the exit area of the 3rd artificial lagoon to quantify the oxidation of ammonium to nitate as a function of water temperature and flow rates. Another project involves the artificial wetlands in which more chemical processes are occurring. Samples can be taken over time at the entrance and the exit area of the wetlands to quantify (as a function of water temperature and flow rate) the oxidation of ammonium to nitate, denitrification of nitrate, and removal of phosphate.

Nutrient Removal with Aluminum Stearate

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Aluminum stearate is an insoluble, hydrophobic powder that bacteria prefer as a metabolic source for denitrification. The destruction of the stearate by the bacteria releases aluminum which precipitates some of the phosphate in solution. The process could be used in the removal of nitrate and phosphate from the oxidized effluent of individual home treatment plants and is being tested on a Sea Grant for that purpose. The powder can be placed on volcanic scoria within a container which is filled with waste waters. At selected time intervals, waste-water samples can be taken to measure the amount of denitrification and phosphate removal as a function of reaction time. See Stoessell et al. (2001) for laboratory data.

Salt Water Intrusion in the Big Branche Aquifer in Lacombe

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The Big Branch Aquifer is at a depth of 1500 feet in the vicinity of Lacombe, Louisiana. The aquifer has a salt component which is not present further updip to the north or to the east and west. The TDS is about 10% that of seawater in the vicinity of Lacombe. Salt-water samples can be taken from old wells completed in the aquifer and fresh-water samples can be taken from other wells completed in the aquifer outside of the salt-water area. Conservative mixing relations (e.g., those using Br, Na, Cl) can be used to infer the salt composition of the salt-water end member and its source. See Stoessell (1997) for an example.

Example References for Research Projects

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Examples of earth-surface (environmental) field and laboratory geochemical studies are listed below that have been published in journals by your instructor and his co-authors. Geochemistry is usually combined with mass flow in the projects. Several of the papers are general applications of interpreting low-temperature geochemical data (Stoessell, 1992; 1997; 1999).

Stoessell, R. K. and J. S. Hanor (1975) A nonsteady state method for determining diffusion coefficients in porous media: J. Geophys. Research, 80, 4979-82.

Stoessell, R. K. and R. L. Hay (1978) The geochemical origin of sepiolite and kerolite at Amboseli, Kenya: Contrib. Mineral. Petrol., 65, 255-267.

Stoessell, R. K. and P. A. Byrne (1982) Methane solubilities in clay slurries: Clays and Clay Minerals, 30, 67-72.

Stoessell, R. K. and P. A. Byrne (1982) Salting-out of methane in single-salt solutions at 25oC and below 800 psia: Geochim. Cosmochim. Acta, 46, 1327-1332.

Stoessell, R. K. (1988) 25oC and 1 atm dissolution experiments of sepiolite and kerolite: Geochim. Cosmochim. Acta 52, 365-374.

Stoessell, R. K., W. C. Ward, B. H. Ford, and J. D. Schuffert (1989) Water chemistry and CaCO3 dissolution in the saline portion of an open-flow mixing zone, coastal Yucatan Peninsula, Mexico: Geol. Soc. Amer. Bull. 101, 159-169.

Moore, Y. H., R. K. Stoessell, and D. H. Easley (1992) Ground-water flow along the northeastern coast of the Yucatan Peninsula: Ground Water, 30, 343-350.

Stoessell, R. K. (1992) Effects of sulfate reduction on CaCO3 dissolution and precipitation in mixing-zone fluids: Journal of Sedimentary Petrology, 62, 873-880.

Stoessell, R. K., Y. H. Moore, and J. G. Coke, (1993) The occurrence and effect of sulfate reduction and sulfide oxidation on coastal limestone dissolution in Yucatan cenotes: Ground Water, 31, 566-575.

Stoessell, R. K. (1997) Delineating the chemical composition of the salinity source for saline ground waters: An example from east-central Concordia Parish, Louisiana: Ground Water, 35, 409-417.

Stoessell, R. K. (1999) Predicted retardations of concentration fronts using a mass-balance approach: Ground Water, 37, 701-706.

Stoessell, R. K., D. H. Easley, and G. H. Yamazaki (2001) Denitrification and Phosphate Removal in Experiments Using Al Stearate: Ground Water Monitoring and Remediation, 21, 89-95.