emissions from other industries and 
from homes and businesses that 
shifted from coal to oil or gas. The 
total amount of sulfur oxides released 
into the atmosphere therefore did not 
increase as dramatically as the amount 
of nitrogen oxides. The increased ni- 
trogen oxide emissions were primarily 
from motor vehicles, factories, and 
power plants, in that order. 
Ironically, the installation during 
this time of extremely tall smoke- 
stacks at some power plants often 
helped to transform a local problem 
into a regional one. The tall stacks 
emit large amounts of pollutants high 
into the air where they are carried 
long distances by the wind, thereby 
gaining more time in which to react 
chemically to form sulfuric and nitric 
acids. 
There is little evidence so far that 
airborne acidity has injured crops and 
forests, but the evidence of environ- 
mental damage in freshwater lakes 
and streams in severely affected re- 
gions is unmistakable. Some lakes are 
more susceptible to acid precipitation 
than others. Among the factors that 
determine the extent to which a lake 
may become acidified are the amounts 
of acid or acid-forming materials that 
are deposited, the sensitivity of the 
soil and bedrock to acidification, and 
the topography of the watershed. 
When scientists speak of a lake be- 
ing acid, they are referring specifically 
to the concentration of ionized hydro- 
gen in the water. In chemical short- 
hand, this form of hydrogen is speci- 
fied as H + . The presence of H + is 
what gives solutions their acidity. 
For the sake of convenience, the 
concentration of H + is often expressed 
in terms of the pH scale. The scale 
is logarithmic, so that a rise or fall 
of one pH unit represents a tenfold 
change. Thus, pH6 is ten times more 
acidic than pH7, and pH5 is a hundred 
times more acidic than pH7. In un- 
polluted, clear, natural waters, the pH 
level is rarely below 6. 
In neutral water (with a pH of 7), 
the H + concentration is very low, 
about one-tenth part of H + to a billion 
parts of water, or 0.1 parts per billion 
(ppb). Pure water, which is in equi- 
librium with the atmosphere, will have 
carbon dioxide dissolved in it, making 
a slightly acidic solution of pH5.6, 
or 2.5 ppb. This is often referred 
to as the pH of “unpolluted” rain. 
Many lakes that have been acidified 
by acid rain now have concentrations 
of 30 parts of H + per billion parts 
of water — 300 times the concentration 
of acid in neutral water. We know 
that the concentration of H + in rain 
in the northeastern United States is 
sometimes as high as 3,000 ppb, which 
is 30,000 times as high as neutral 
water and more than 1,000 times more 
acidic than “unpolluted” rain. 
The ability of water to neutralize 
H + is called its buffer capacity, and 
it is determined, for the most part, 
by the alkalinity of the water. Alka- 
linity, in turn, comes primarily from 
the weathering of minerals in the wa- 
tershed. Limestone and other, similar 
minerals are the most important sub- 
stances that provide alkalinity and, 
hence, buffering capacity. 
As acidification proceeds, the alka- 
linity is removed in the form of carbon 
dioxide and lost from the water. This 
diminishes the water’s acid-neutraliz- 
ing capacity. When acids deposited 
from the atmosphere percolate 
through a watershed, aluminum is also 
dissolved from soil and rocks. Hydro- 
gen ions are exchanged for aluminum 
ions in the minerals of soil and rocks. 
By decreasing the amount of H + , this 
process may decrease acidity, but alu- 
minum can be a threat to fish. Alu- 
minum concentrations, which are usu- 
ally well below 0.1 parts per million 
(ppm) in lakes and streams, are greatly 
elevated through the exchange of H^ 
in regions with very acidic precipi- 
tation. Acidified waters in the Ad- 
irondacks, for example, have concen- 
trations of 0.2 to 1 ppm of aluminum. 
At 0.2 ppm, aluminum is quite toxic 
to fish. 
Most of the regions in which lakes 
and streams are acidified have thin 
soils and relatively large amounts of 
George R Hendrey 
60 
