Groundwater
Chemistry
9.1
Introduction
If H2 O molecules were the only thing present in groundwater, this chapter could be very
short. Thanks to the many other substances within or in contact with groundwater, there is
a lot more to talk about. The distribution, reactions, and transport of these other substances
make for an interesting and complex topic.
Solutes are other molecules dissolved within the sea of H2 O molecules in the aqueous
state. Many solutes occur naturally, such as inorganic ions like Ca2+ or SO2−
4 . Sometimes
high concentrations of naturally occurring solutes render the water unfit for drinking,
irrigation, or other uses. Other solutes are chemicals introduced by human activities.
Many of these are troublesome contaminants such as heavy metals and organic solvents.
The solid phases that make up the aquifer matrix can react with and dissolve into the
groundwater. At the same time, some solids precipitate out of water, a phenomenon that
can lead to clogged pipes. Some solids may also exist as tiny particles suspended in the
groundwater.
In the unsaturated zone, water is in contact with pore gases and molecules will transfer
between the liquid and gas states. This mechanism can be an important way that subsurface contaminants migrate, particularly for volatile organic compounds (VOCs).
When organic liquids like hydrocarbon fuels and solvents are spilled into the subsurface, they dissolve sparingly into water and can persist for a long time as a separate liquid
phase. The acronym NAPL, for nonaqueous-phase liquid, is often used to describe these
separate liquid phases.
Chemical reactions can involve substances in the aqueous, gas, solid, or NAPL phases,
and some reactions transfer mass from one phase to another. Some reactions occur within
the bodies of microorganisms; they are a vital link in the attenuation of certain contaminants. Chemical processes in the groundwater environment are both complex and
fascinating. Characterizing and predicting these processes are some of the most challenging problems in groundwater science. Groundwater chemistry is relevant to all users of
groundwater resources, whether it be for drinking, irrigation, industrial, or other purposes.
Chemistry is also central to understanding the fate of groundwater contamination and how
to remediate contamination.
9
280
Groundwater Chemistry
This chapter provides an introduction to aqueous geochemistry as it relates to groundwater. Much more detailed treatment of the subject can be found in aqueous chemistry
texts like those of Drever (1988), Pankow (1991), Morel and Hering (1993), Stumm and
Morgan (1996), and Langmuir (1997). Domenico and Schwartz (1998) cover aspects
of chemistry that are relevant to groundwater. The next chapter introduces groundwater
contamination, building on the fundamentals introduced in this chapter.
9.2 Molecular Properties of Water
The geometry of a water molecule is not unlike the face of a famous cartoon mouse
(Figure 9.1). The two hydrogen atoms are bonded to the oxygen atom by sharing outer
electrons, forming covalent bonds. The angle between the two bonds is about 105◦ .
Water is a polar molecule because the distribution of electrical charge associated with
protons and electrons is asymmetric. The oxygen end of the molecule is somewhat negatively charged, while the hydrogen ends are somewhat positively charged.
The polarity of the water molecule causes electrostatic attraction to other polar molecules and to charged molecules. The hydrogen ends of a water molecule are attracted to the
oxygen ends of other water molecules, forming weak bonds known as hydrogen bonds
(Figure 9.1). Hydrogen bonding causes water molecules to bond together in clusters
within which there is an ephemeral fixed arrangement like in a crystalline solid. These
clusters are continually forming and breaking up, existing for only a short slice of time,
on the order of 10−12 sec (Stumm and Morgan, 1996). These clusters grow as large as
100 water molecules (Snoeyink and Jenkins, 1980).
Several different isotopes of both hydrogen and oxygen occur in natural waters, but the
most common isotopes 1 H and 16 O are far more abundant than all others (see Table 9.13,
Section 9.10). The different isotopes of a specific element differ only in the number of
neutrons in the atom’s nucleus and, of course, their total mass. Because various isotopes
of the same element have the same number of electrons, all isotopes behave similarly in
chemical reactions.
The difference in mass between different isotopes can lead to different behavior in
certain physical processes. For example, water molecules containing the heavier isotopes
2
H, 17 O, or 18 O are less prone to evaporate from liquid water than the common water
molecules containing 1 H and 16 O. This discrepancy leads to near-surface ocean or lake
water that is enriched in the heavier isotopes compared to atmospheric water (more about
that in Section 9.10).
Figure 9.1 Geometry
of a water molecule (left)
and hydrogen bonding of
water molecules (right).
Solute Concentration Units
281
Natural waters are water-based (aqueous) solutions, with other elements and compounds are dissolved as solutes within the solvent water. Most solutes in natural groundwaters carry a charge, either as cations (+) or anions (−). Ions dissolved in water are
typically surrounded by water molecules that orient themselves in accordance with the
charge of the ion, as shown in Figure 9.2. The larger the ion, the more oriented water
molecules can surround it. The orientation of water molecules extends beyond the adjacent layer of water molecules, but the degree of orientation decreases with distance from
the ion.
The polar nature of water molecules makes it a good solvent of ionic and polar molecules. The mutual attraction of ions and polar water molecules allows large numbers of
ions to be accommodated in the midst of water molecules, resulting in high solubilities
for ionic substances. Table salt (NaCl) and other salts dissolve readily into their ionic
components.
On the other hand, nonpolar molecules have a relatively symmetric distribution of
charge and little affinity for water molecules. Lacking attraction to water molecules, relatively few of these nonpolar molecules are accommodated within the water. Nonpolar
molecules have low solubilities; they only dissolve to low concentrations in water.
9.3
Solute Concentration Units
The concentration of a solute in an aqueous solution may be expressed in several different
ways. In chemical calculations, it is standard to use molar concentration units, which are
moles of solute per liter of solution (mol/L), denoted M. For example, a 2.5 M Ca2+ solution contains 2.5 moles of Ca2+ per liter. A mole is an amount of a substance consisting
of N atoms or molecules, where N = 6.022 × 1023 is Avogadro’s number, here rounded
to four significant digits. The mass of a mole of atoms is called the atomic mass and
the mass of a mole of molecules is called the formula mass (also called formula weight).
For example, the atomic mass of oxygen is 16.00 g, and the formula mass of CO2 is
12.01 + (2 × 16.00) = 44.01 g. For calculations involving chemical reactions, it is handy
to use moles and molar concentrations, because chemicals react in direct proportion to the
numbers of molecules present (a periodic table of elements is on the back inside cover).
The concentrations measured in a laboratory water analysis are usually reported in
mass/volume units like mg/L or µg/L. Conversion between molar and mass per volume
units may be done using the formula weight of the solute. Converting from mol/L to
mg/L units could be done as follows:
Figure 9.2 Orientation
of water molecules
around dissolved cation
(left) and anion (right).