11/24/04

Introduction to groundwater chemistry

I. Motivations for studying groundwater chemistry

II. Properties of water

The bonds between hydrogen and oxygen cause a separation of charge and the tetrahedral structure of water molecule clusters.

The polar nature of water makes it a good solvent for ions (e.g Na+ and Cl-) and for other polar molecules (e.g. ethyl alcohol). It is not a good solvent for non-polar molecules (e.g. vegetable oil). Many organic contaminants are non-polar and only sparingly soluble in water.

III. Units of concentration

Solute concentrations are generally reported either per unit mass of solution (or solvent) or per unit volume of solution.

Common mass solute/mass solution units include mg/kg (identical to ppm or to mg/g) and mg/kg (equal to ppb).

Common mass solute/volume solution units include mg/L and mg/L. For dilute aqueous solutions, a liter of solution has a mass of approximately 1 kg, so these units are often used interchangably with the mass/mass units listed above. This equivalence cannot be assumed if you are dealing with seawater or brine.

A more useful concentration unit is one that represents the number of molecules or ions dissolved in either a unit mass or unit volume of water. Units of "moles" (6 x 10E23) are a convenient way to express the number of molecules or ions dissolved. The molal concentration is moles solute/kg solvent (i.e. water). This is a convenient scale to use when mixing a solution, since you can start with exactly one kg of distilled water. This is also the scale that is assumed for thermodynamic calculations. The molar concentration is moles solute/L solution. This concentration varies with temperature, pressure and other solutes in solution that can affect the density of the solution. In dilute solutions at room temperature, the molal and molar concentrations are essentially identical and so they tend to be used interchangeably. As in the case of the mass units described above, this equivalence breaks down when you are considering solutions with high concentrations of solutes.

In addition to converting mass/volume or mass/mass concentration units for ionic solutes to units of mol/L (molar, M) or mol/kg (molal, m) concentration, it is also useful to determine the moles of charge corresponding to a given molal or molar concentration. A mole of charge is one "equivalent". Since we are often working in concentrations that are much less than one mole per liter or kilogram, it is often convenient to work in units of milliequivalents (0.001 equivalents). An ion that has a single charge (monovalent) will have one equivalent for each mole of ions. An ion that has a charge of two (divalent) will have two equivalents for each mole of ions. An example of conversion from mg/L to meq/L is provided in the text. Calculating moles of charge is a first step to determining if there is a significant charge balance error.

IV. Common solutes in groundwater

Common solutes in groundwater are listed in Table 16.1 of S&Z. Some of these are common because they are common components of the atmosphere (nitrogen, oxygen, dissolved carbon dioxide and ions that come from dissociation of dissolved CO2). Others are common in groundwater because they are abundant in common rock forming minerals that dissolve as groundwater moves through sediment and rock (silicic acid, calcium, magnesium, sodium). Still others are extremely soluble and are abundant in the ocean, having accumulated during early degassing of the earth's crust and continued volcanic emissions (chloride, sulfate). The relative abundance of constituents in groundwater is a function of both abundance in the crust, ocean or atmosphere and the mobility of the element in a geochemical cycle.

V. Common water quality paramaters and water quality standards

Table 16.2 of S&Z lists commonly measured inorganic parameters for a routine water analysis. This generally includes the major cations (Ca, Mg, Na, K) and major anions (Cl, SO4, HCO3) as well as several other ions that are important to domestic use of water (Fe, nitrate, flouride). In addition, a routine analysis usually includes pH (negative log10 of the hydrogen ion "activity"), alkalinity (a measure of acid neutralizing capacity), total dissolved solids (either calculated or measured from residue left after evaporation), electrical conductivity (the measurement we made during the field exercise earlier in the semester) and hardness (a function of Ca and Mg in solution - hard water can precipitate "scale" in pipes, leading to clogging, and inhibits formation of soap suds, making it difficult to wash things). Water quality standards for microbial, organic and inorganic constituents in water are summarized by the EPA at

http://www.epa.gov/safewater/mcl.html

VI. Equilibrium models for chemical reactions

Equilibrium is a special type of dynamic steady state for reversible reactions. At equilibrium the reaction does not stop, but the rates of transformation of reactants to products are exactly balanced by rates of transformation of the "products" back to "reactants".

Equations 17.1 and 17.2 of S&Z are generic representations of equilibrium relations that are often referred to as the Law of Mass Action. While S&Z suggest that the appropriate "concentrations" for use in 17.2 are molal or molar concentrations, these should actually be "activities" (or effective concentrations") as defined in equation 17.3 on the following page. Activities will be discussed in more detail in the next lecture.

The equilibrium constant K is related to the Gibbs (or standard) free energy of the reaction, which can be calculated from standard free energies of reactants and products in the reaction. There are many tabulations of free energies, which makes it possible to determine equilibrium constants for a large number of reactions. The equilibrium relation summarized in 17.2 (using activities) holds both for "elementary" reactions that could occur at the molecular level or for "net" reactions that are the sum of a number of elementary reactions in series or parallel. As an example, consider the two reactions

and

with equilibrium relations (shown with parentheses "( )" but should really be "[ ]" to represent activities)

and

These can be combined to obtain the net reaction

for which the equilibrium relation is

Note that the equilibrium constant for the net reaction equals the product of the equilibrium constants for the elementary reactions.