After completing this chapter, you will be able to:
- Describe the most important chemical ingredients of precipitation and explain which ones may cause acidity to develop.
- Outline the spatial patterns of acidic precipitation in North America, and identify factors influencing this distribution.
- Explain the difference between the wet and dry deposition of acidifying substances and how their rates vary.
- Describe how water chemistry is affected as precipitation interacts with vegetation and soil, and explain the implications for surface waters.
- Identify factors that make fresh waters vulnerable to acidification.
- Describe the effects of acidification on freshwater organisms.
- Discuss the roles of liming and fertilization in the reclamation of acidified lakes.
- Explain the importance of reducing emissions of sulphur and nitrogen gases to mitigating the acidification of surface waters.
Acidification is a process that is characterized by increasing concentrations of hydrogen ions (H+) in soil or water. It can cause metals and their compounds to ionize, producing ions (such as Al3+) in concentrations high enough to be toxic to plants, animals, and microorganisms. Consequently, increasing acidification is usually interpreted as a degradation of environmental quality. Acidification is caused by many influences, both natural and anthropogenic, but the most widespread problems are associated with a phenomenon commonly referred to as acid rain.
Acid rain has been an important problem in parts of North America since at least the 1950s, but it did not become a high-profile issue until the early 1970s. This rather sudden attention resulted from the discovery that acid rain was a widespread problem in Western Europe, and the realization that the same conditions likely occurred in North America. This awareness stimulated research in Canada and the United States, which demonstrated that acid rain was causing a widespread acidification of lakes and streams, and possibly of soil. The acidification of aquatic ecosystems was resulting in important ecological damage, including the loss of many fish populations. Buildings and other materials were also being damaged because acidity erodes metals, paint, and some kinds of quarried stone.
Strictly speaking, the term “acid rain” refers only to acidic rainfall, which along with snowfall accounts for wet deposition. However, acidifying chemicals are also deposited from the atmosphere when it is not raining or snowing, through the dry deposition of certain gases and particulates. A suitable phrase to define this complex of processes is “the deposition of acidifying substances from the atmosphere”, or more simply, acidifying deposition. In this chapter we examine natural and anthropogenic causes of the acidification of ecosystems. We focus on the chemical qualities of acidic precipitation and dry deposition, their effects on terrestrial and aquatic ecosystems, and how acidification can be avoided or mitigated.
In Detail 19.1. Acids and Bases
An acid is defined as a substance that donates protons (hydrogen ions, H+) during a chemical reactions. An aqueous solution is acidic if its concentration of H+ is more than 1 × 10-7moles per litre. In contrast, a base (alkali) donates hydroxyl ions (OH–) in chemical reactions. A solution is basic if its concentration of OH– exceeds 1 × 10-7mol/L. (A mole is a fundamental unit that measures the amount of a substance and is equal to 6.02 × 1023 molecules, atoms, or ions. This number is known as Avogadro’s constant and it is derived from the number of carbon atoms contained in 12 g (1 mole) of carbon-12.)
Acids and bases react together to form water and a neutral salt. If equal numbers of moles of each are present, the solution has both zero acidity and zero alkalinity – the concentrations of H+ and OH–are both exactly 1 × 10-7mol/L. Such a solution is said to be neutral.
Because extremely wide ranges of H+ and OH– concentrations are encountered in nature and in laboratories, acidity is measured in logarithmic units, which are referred to as pH (an abbreviation for “potential of hydrogen”). pH is defined as –log10H+], or the negative logarithm to base 10 of the aqueous concentration of hydrogen ion, expressed in units of moles per litre. Acidic solutions have a pH less than 7.0, while alkaline solutions have a pH greater than 7.0. Note that a one-unit difference in pH implies a 10-fold difference in the concentration of hydrogen or hydroxyl ions. The scale illustrated below shows the pH of some commonly encountered substances.
Figure 19.1. The pH scale and the pH of some familiar substances.
Chemistry of Precipitation
Scientists have adopted a functional definition of acidic precipitation as having a pH less than 5.65. This was chosen as the cut-off because at pH 5.65, an aqueous solution of carbonic acid (H2CO3) is in equilibrium with atmospheric CO2, as follows: CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3– ⇌ 2H+ + CO32-
This definition assumes that “non-acidic” precipitation is essentially distilled water, in which the acidity is determined only by the atmospheric concentration of CO2 and the amount of carbonic acid that subsequently develops. This is why the threshold below which precipitation is deemed “acidic” is set at the slightly acidic pH of 5.65, rather than at the strict zero-acidity pH of 7.0 (see In Detail 19.1).
It is, however, too simplistic to consider atmospheric moisture as consisting merely of distilled water in a pH equilibrium with gaseous CO2. Additional chemicals are also present in trace concentrations in precipitation. For example, on windy days, dust containing calcium and magnesium is blown into the atmosphere, and precipitation containing these elements may develop a pH higher than 5.65. This is especially true of agricultural and prairie landscapes, where the ground surface is often bare of plant cover and soil particles may be easily eroded into the atmosphere. In some other regions, a relatively high concentration of naturally occurring sulphate in the atmosphere may result in precipitation having a pH less than 5.65.
The most abundant cations (positively charged ions) in precipitation are hydrogen ion (H+), ammonium (NH4+), calcium (Ca2+), magnesium (Mg2+), and sodium (Na+). The most abundant anions (negatively charged ions) are sulphate (SO42-), chloride (Cl–), and nitrate (NO3-). Other ions are also present, but only in trace amounts that have little influence on the pH (see In Detail 19.2).
In Detail 19.2. Conservation of Electrochemical Neutrality
The principle of conservation of electrochemical neutrality states that in any electrically neutral solution (one that does not carry an electrical charge), the total number of positive charges associated with cations must equal the total number of negative charges of anions. For the purposes of calculating a charge balance, the concentrations of ions are measured in units known as equivalents. These are calculated as the molar concentration multiplied by the number of charges on the ion. (When dealing with precipitation or surface waters, microequivalents, or µeq, are generally the units reported.)
This principle is relevant to the acidification of water. The concentration of H+ can be determined as the difference in concentrations of the sum of all anion equivalents minus the sum of all cations other than H+. Therefore, if the total equivalents of anions exceed the total equivalents of cations other than hydrogen ion, then H+ must go into solution to balance the cation “deficit,” as follows: H+ = (SO42- + NO3– + Cl–) – (Na+ + NH4+ + Ca2+ + Mg2+)
The above equation has proven to be useful in studies of acidic precipitation. Prior to about 1955, the measurement of pH was somewhat inaccurate. There were, however, reliable analyses of other important ions in surface waters and precipitation. In such cases, the equation can be used to calculate pre-1955 pH values, providing important data for studies of the historical pH in waters sensitive to acidification.
One of the longest-running records of precipitation chemistry is from a research site at Hubbard Brook, New Hampshire, in a region exposed to intense acidifying deposition. During 1967-1971, when acid rain was relatively severe, the average pH of precipitation at Hubbard Brook was 4.1. This level of acidity then relaxed somewhat to pH 4.9 in 1991-1995 because of decreased industrial emissions, particularly of the acid-forming gas SO2, and then even more so in 2009-2013 because of further decreases in SO2emissions (Table 19.1; see In Detail 19.2 for an explanation of equivalents). Sulphate and nitrate are the most important anions in precipitation, and from 1967-1971 they occurred accounted for 88% of the anion equivalents. During 2009-2013 these two still contributed 87% of the anion equivalents, although their total amounts were considerable smaller. These data suggest that most of the acidity in the precipitation occurs as dilute solutions of sulphuric and nitric acids. The precipitation events at Hubbard Brook that are most acidic are associated with storms that have passed over the large metropolitan regions of Boston, New York, and New Jersey. These areas have enormous emissions of SO2 and NOx, which are the precursor gases of much of the SO42- and NO3– in acidic precipitation.
Table 19.1. Chemistry of Precipitation at Hubbard Brook. The data represent the average concentration (in microequivalents per litre, or µeq/L) of various ions in precipitation during three 5-year periods. The small difference between the sums of cation and anion equivalents is due to analytical inaccuracies, which are inevitable in even the best chemical data. Source: Data from Buso et al. (2003).