Geochemistry

The importance of understanding the geochemical evolution of groundwaters as a basis for isotope hydrogeology cannot be overstated. Interpretations of stable and radioactive isotopes in hydrogeological systems must be complimented by an evaluation of the inorganic and organic geochemistry of groundwater. A variety of special sampling and preservation techniques are required for the analysis of major and certain trace species in natural waters. The following discussion outlines some methods and precautions.

A general practice in sampling water for chemical analysis is to filter the sample through a 0.45-micron pore diameter filter. This will remove most bacteria, almost all suspended clays and a proportion of iron and manganese oxyhydroxides, but will not retain viruses or some organic molecules such as fulvic and humic acids. Filtering is required to assure that the laboratory analysis represents dissolved species (which take part in most geochemical reactions and are used in chemical equilibrium equations) and not suspended constituents, which may be contributed from the wells.

Field measurements

Geochemical studies are based on the measurement of the inorganic constituents or species in a groundwater and a series of parameters that control the interactions of these species. Certain parameters such as temperature and pH are difficult to preserve during storage and should be measured in the field.

Temperature: The temperature of groundwater is a fundamental measurement and is required in all thermodynamic calculations and models. It is also needed to correct EC measurements and for calibrating the pH meter. It must be measured as close to in situ conditions as possible. Use a standard glass thermometer, or the temperature probe that is a component of most pH meters.

Electrical conductivity (EC): Electrical conductivity is proportional to the quantity of dissolved ions present in solution and can provide a rough idea of the total dissolved solids (TDS). For most groundwaters, the EC value, in mS/cm corrected to 25^oC, is about 50% greater than the TDS expressed as mg/L, and can be estimated according to:

pH: A reliable pH measurement is one of the most important field parameters to be measured, and must be made with care and patience. pH is an expression of the negative log of H⁺ activity (pH = –log [H⁺]). It is fundamental to thermodynamic calculations, and to the interpretation of carbonate and d¹³C data. In natural waters, pH is generally between about 6.5 and 8. Measurement at a sampling site requires the calibration of the meter with reference buffer solutions of known pH according to the following recommended protocol.

REDOX potential: Redox measurements provide a measure of the electromotive force of a water, or the relative dominance of oxidized vs. reduced species in solution. Electromotive potential (E) is expressed in volts, relative to the standard hydrogen electrode (Eh), which by this convention has zero potential (Eh = 0 volts). It is generally measured with a silver-silver chloride electrode, which has its own standard potential, E_R.

As electrodes are simply measuring an electromotive force, there is no calibration. The electrode can, however, be checked against a solution of fixed Eh to determine whether it is functioning. Two solutions with fixed Eh are Zobell’s (ferric-ferrous cyanide solution) and quinhydrone [pH 4 and 7 buffers saturated with quinhydrone powder (~0.2 g/100 mL]:

Alkalinity: HCO₃^– and CO₃^2– concentrations are measured by an alkalinity titration. In most natural waters, total alkalinity can be expressed as bicarbonate concentration. CO₃^2– only becomes an important component of DIC above pH 8.5. Other bases can act as proton acceptors, and so alkalinity is defined as:

A variety of organic compounds and colloids can also contribute to alkalinity, although HCO₃^– remains the major contributor in most natural waters. It is determined by titrating with an acid (generally H₂SO₄) to an end point near pH 4.3. Carbonic acid is a weak acid, and so the titration end point increases to about pH 5 in low alkalinity waters. The volume of acid (and normality N) required is proportional to the alkalinity, which is most often expressed as equivalent concentration of HCO₃^–:

Alkalinity (mg-HCO₃^–/L) =

Field measurements are greatly aided by using field kits (e.g. Hach, Merck or others) that replace pipettes, burettes and bottles of acid with a microtitrator and syringes of acid. The titration end-point can be identified with a colour indicator (usually Bromcresol), but it is better to use a pH meter and record about 5 to 10 readings below pH 6 (Fig. Chapter 10 -1).

It is recommended to perform an alkalinity titration in the field, although many will argue that alkalinity is conservative, even if degassing or calcite precipitation takes place. Nonetheless, reactions such as oxidation of Fe²⁺, degassing of H₂S or equilibration with atmospheric CO₂ can have an effect. Field titrations with commercially available kits are simple to do and rival laboratory bench auto-titrators in accuracy.

SYNOPSIS OF FIELD MEASUREMENTS:

SAMPLE: Ideally, field measurements are made in a flow-through cell with a low, constant flow of groundwater. Measurements can also be made in aliquots taken from a well or spring, although one must consider the effects of atmospheric contact (CO₂ degassing, oxidation with O_2, etc.).

EQUIPMENT: ·   Flow-through cell, beaker or sample container.

pH: ·   Select two buffers that bracket the anticipated pH of the groundwater, i.e. pH 4 and 7, or 7 and 10.

·   Equilibrate electrodes and buffers to the temperature of the water to be measured. Failure to do this will result in slowly drifting pH readings as the electrode thermally equilibrates with the sample.

·   Follow manufacturer instructions for specific instrument calibration procedures, with corrections for temperature of buffers:

Eh: ·   Test electrode and meter for functioning against standard solution.

·   Stir sample gently in measuring if there is no flow in the measurement container, to allow electrode to "see" all of solution.

·   Wait for a stable reading (1 to 5 minutes). If readings are dropping and begin to rise, record lowest reading as atmospheric O_2; contamination may be occurring.

·   Correct the measured E to Eh by adding the standard potential, E_R, to the measured value, according to the equation: Eh = E_measured + E_R. The following are the standard potentials for a Ag-AgCl electrode with 3.8 N KCl filling solution (according to the equation E_R = –1.39T(°C) + 305):

·   Clean electrode if there are problems with reference readings or time for stabilization. Use liquid hand soap and a soft brush. For scale, dip in 10% HCl for a couple of minutes and rinse with deionized water.

·   Do not use the Eh aliquot as a sample due to K⁺ and Cl^– contamination.

·   Store electrode in a soaker bottle (bottle with O-ring seal), filled with 2 N KCl.

ALKALINITY:

·   Measurement should be made as close to point of sampling as possible due to alteration of the sample alkalinity from CO₂ degassing, Fe²⁺ oxidation or other reactions. If this is not possible, take filtered sample in a HDLP or glass bottle (fill completely), and analyze as soon as reasonably possible (store cool).

Sulphate should be measured if samples are collected for either sulphate or carbonate isotope analyses, to ensure that sufficient sample is collected, and to calculate the amount of barium chloride needed to fully precipitate the sulphate and carbonate. For this purpose a field kit can be used. Alternatively, a simple test for sulphate is to add barium chloride to an aliquot of sample water (adjust pH to below 6) which will turn cloudy if SO₄^2–- is present. Pour slowly into a graduated cylinder (100 mL will do) with a black X on the bottom until the X is no longer visible. The sulphate concentration is roughly determined from the height of water in the cylinder (in cm), using the equation:

Major anions (Cl^–, F^–, SO₄^–, NO₃^–, Br^–)

Analyses of most anions is normally are done in a laboratory by liquid chromatography, although field kits and ion-specific electrodes exist for most anions of interest. HCO₃^– and CO₃^2– are determined by titration. Less than 5 mL of sample can be analysed for most major anions, although using 25- to 50-mL samples allows repeat measurements, and dilutions for improving precision of some species. Use thoroughly rinsed HDPE or PP bottles. Samples should be filtered (0.45 mm), stored cool, and — obviously — unacidified.

Sulphate analyses on water that contains reduced sulphur species (H₂S, HS^– or S^2–) can be very erroneous since they oxidize very fast and sulphate readings are then too high. It is mandatory to remove dissolved sulphide, preferably as CdS since its yellow colour is very indicative of the presence of reduced S-species. Filter before analysing for sulphate. A second 25-mL sample is required for total S (H₂S and SO₄^2–), filtered out of contact with atmosphere. Preserve with 0.25 mL of 30% H₂O₂ to oxidize all sulphide to sulphate. This measurement for SO₄^2– , less that of the first sample, gives total sulphide.

Major, minor and trace metals

All geochemical analyses require measurement of major metals or cations (Ca²⁺, Mg²⁺, Na⁺, K⁺ and SiO₂). Minor (Fe, Mn, Sr and Ba) and trace metals may be required as well. Cations are best analysed in the laboratory by routine methods, including ICP-AES (inductively coupled plasma-atomic emission spectrometry), ICP-MS (ICP-mass spectrometry), AA (atomic absorption) or FE (flame emission), although several field methods are also available. Only a few millilitres of sample are required for ICP instruments that measure multiple wavelengths or masses simultaneously. A normal sample size is 25 to 50 mL.

Carbonates and oxides can precipitate after sampling, and so samples are acidified to keep metals in solution. To ensure that only dissolved species are analysed, the solution must be filtered first. Ultrapure nitric acid is used to acidify to below pH 2. Reagent grade is adequate for major and minor cations only. HCl should be used for brines as it releases hydrogen ions faster. The water must be filtered and should be stored in fresh HDPE or PP bottles. Used bottles are fine too, but should be acid rinsed.

Dissolved organic carbon (DOC)

Sampling for DOC is recommended for any study where redox processes and the carbonate system are important. It is measured in the laboratory by oxidation of an acidified sample (initially degassed of CO₂) and measuring the evolved CO₂. The sample must be filtered to 0.45 mm at the point of sampling, and acidified to stop bacterial activity. Acidify with HCl, because HNO₃ can oxidize organics in the sample prior to analysis. Collect a 10- to 20-mL sample in a glass vial with a non-organic cap seal (aluminum foil, for example). Scintillation vials are ideal and cheap.

SYNOPSIS OF GEOCHEMICAL SAMPLING:

SAMPLE VOLUME: ·   25 to 50 mL for anions.

EQUIPMENT: ·   HDPE or PP bottles.

METHOD: ·   Take the required field measurements of T, pH etc.