Evaluation of Alkalinity in Lake Monroe
and Griffy Lake in Indiana, and Arboretum Pond
at Indiana University
Young Jun Moon+ and Linda Wozniewski++
+
School of Public and Environmental Affairs, Indiana University, Bloomington, Indiana 47405
Department of Chemistry, Indiana University, Northwest, Indiana 46408
++
ABSTRACT
In March, 2003, the alkalinity in Lake Monroe and Griffy Lake in southern Indiana, the
Arboretum Pond at Indiana University was determined. The three water samples were
titrated with strong acid (H2SO4) and a Gran-plot analysis was used to obtain an exact
value for the endpoint. The alkalinity in Lake Monroe, Griffy Lake, and the Arboretum
Pond was 25.89, 113.39, and 70.38 mg CaCO3/L, respectively. The alkalinity was
caused by calcium carbonate (CaCO3) and related to the morphology of water systems.
Therefore the alkalinity in Lake Monroe, which has more watersheds, was lower than
that in Griffy Lake, which has fewer watersheds. The alkalinity of the Arboretum Pond,
which is located at university, was between two lakes.
INTRODUCTION
The alkalinity of water is a measure of its capacity to neutralize acids. The ability to
resist changes in pH by neutralizing acids or bases is called buffering. This measure of
acid-neutralizing capacity is important in figuring out how buffered the water is against
sudden changes in pH. Therefore sufficiently alkaline water is more resistant to acid
precipitation and other acidic inputs than poorly buffered systems.
Although many species contribute to alkalinity of water, the most important compounds
in water that determine alkalinity are carbonate (CO32-) and bicarbonate (HCO3-) ions.
One source of alkalinity in southern Indiana is calcium carbonate (CaCO3), which is
dissolved in water flowing through geological formations that have limestone and/or
marble. Limestone is a sedimentary rock formed by the compaction of fossilized coral,
shells and bones. Limestone is composed of the minerals calcium carbonate (CaCO3)
and/or dolomite (CaMg(CO3)2), along with small amounts of other minerals. Limestone
is converted to marble from the heat and pressure of metamorphic events.
Alkalinity can increase the pH, when the alkalinity comes from a mineral source such as
calcium carbonate (CaCO3). When CaCO3 dissolves in water, the carbonate (CO32-) can
react with water to form bicarbonate (HCO3-), which produces hydroxide (OH-):
CaCO3 (s) ↔ Ca2+ + CO32-
(1)
CO32- + H2O ↔ HCO3- + OH-
(2)
The hydroxide ion (OH-) is a strong base. An increase in OH- concentration will cause
the pH to increase.
In watersheds where calcium carbonate (CaCO3) isn’t available, on the other hand,
carbonic acid (H2CO3) is an important source for carbonate (CO32-) and bicarbonate
(HCO3-). Carbon dioxide and water are converted to carbonic acid through the
following reaction:
CO2 + H2O ↔ H2CO3
(3)
Carbonic acid provides bicarbonate and carbonate ions for buffering:
H2CO3 ↔ HCO3- + H+
(4)
HCO3- ↔ CO32- + H+
(5)
While conversion of carbon dioxide to carbonic acid produces ions capable of buffering
pH, it also causes a decrease in pH because 2 hydrogen ions (H+) are released into the
water during the reaction.
In this study the alkalinity of three local water systems, Lake Monroe and Griffy Lake
in southern Indiana, and the Arboretum Pond at Indiana University in Bloomington,
were determined and analyzed.
EXPERIMENTAL SECTIONS
Principles
When titrating for total alkalinity, there are two distinct inflection points (equivalence
points): pH 8.3 and pH 4.6, where pH changes rapidly with small additions of acid.
Water that has a pH > 8.3 is said to have phenolphthalein alkalinity, which is alkalinity
due to the presence of carbonate or hydroxide ions.1) Bicarbonate is the predominant
carbonate species at approximately pH 8.3:
CO32- + H+ → HCO3-
(6)
Total alkalinity is the final endpoint for the alkalinity titration. At pH 4.5, all carbonate
and bicarbonate ions have been converted to carbonic acid (H2CO3):
HCO3- + H+ → H2CO3
(7)
Method
Three local water systems, Lake Monroe and Griffy Lake, and the Arboretum Pond
were selected to determine and analyze the alkalinity. The water samples from two lakes
were collected in March 2003 and the water sample from the Arboretum Pond, whose
area was glossy, was collected on Mar. 27, 2003. The water samples were allowed to
warm to room temperature and pH meter (Orion 720A) was also calibrated before
beginning the experiment. After the sample was equilibrated, color indicator
(bromocresol) was added. The three water samples were titrated with strong acid
(H2SO4). The titration volume and corresponding pH were recorded. Color indicators
have been used to detect endpoint conditions. Since they are not as accurate as the pHcurve construction method, a Gran-plot analysis was used to obtain an exact value for
the endpoint.
Calculations (Gran-plot method)
Only the pH 4.6 region was used for total alkalinity. Two curves were plotted on the
same graph with the titrant volume against F1 and F3:
F1 = (V0 + V)(10-pH)
(8)
F3 = (V)(10+pH)
(9)
Where V0 = sample volume (mL) and V = titrant volume (mL)
The linear portions of the curves were extrapolated to obtain V2, which is the titrant
value at the inflection point. Alkalinity was determined through the following equation:
Alkalinity (mg CaCO3/L) = 5 ×104(V2)(N)/(mL sample)
Where N = normality of titrant and V2 = endpoint value (mL)
(10)
RESULTS AND ANALYSIS
Alkalinity values varied among three different local water systems, Lake Monroe,
Griffy Lake, and the Arboretum Pond, ranging from 25.89 to 113.39 mg CaCO3/L in
March as shown in Table 1. This experiment indicated that the alkalinity in Griffy Lake
was higher than that in Lake Monroe and the alkalinity in the Arboretum Pond was
between the alkalinity values in two lakes.
Table 1 The Alkalinity of three local water systems
alkalinity
Lake Monroe
Griffy Lake
Arboretum pond
(mg CaCO3/L)
(mg CaCO3/L)
(mg CaCO3/L)
25.89
113.39
70.38
The alkalinity in Lake Monroe in March was similar to the data in literature that
indicated that alkalinity values for Lake Monroe showed little spatial variation among
basins and range from 25 to 30 mg CaCO3/L in the spring to 45 mg CaCO3/L in the
fall.2),3) In addition, other data are shown in Table 2. Since the major source for
alkalinity includes calcium carbonate (CaCO3), it could be derived that Lake Monroe
did not have much limestone. It is also consistent with the general lack of limestone in
most of the watershed and in the chemical nature of the Bordon rocks.2) In addition to
calcium carbonate (CaCO3), alkalinity is often inversely proportional to discharge. At
high discharges, the water has little contact time with carbonate-bearing rocks and does
not gain much alkalinity. In spring time snow-melting also influences to alkalinity.
Usually it affects low alkalinity because the water does not have much time to interact
with the geology such as the influence of discharges of water. On the other hand, the
alkalinity in Griffy Lake was higher than that in Lake Monroe. It was also similar to the
literature data as shown in Table 2. Alkalinity can also be influenced by salts, plant
activity, and wastewater. Wastewater can have higher alkalinity because it has higher
concentrations of nutrients and irons. The alkalinity in Arboretum Pond at Indiana
University was 70.38 mg CaCO3/L, which was between the alkalinity values in the two
lakes. It could result from morphology and location of the pond. The Arboretum Pond at
Indiana University is almost stable and did not have much discharge and snow-melting
influences. However it could have nutrients and irons since the pond is located at the
university. In addition, stormwater runoff may carry limestone (either calcite or
dolomite), which is applied to lawns and agricultural fields.
Hardness in water is caused by dissolved calcium and magnesium. It is usually
expressed as the equivalent quantity of calcium carbonate (CaCO3). Hardness also
measures the buffering capacity or the ability to absorb and neutralize added acid
without major changes to pH. This buffering is referred to as alkalinity.4) There are two
types of hardness, carbonate and noncarbonate. Carbonate hardness is caused by the
carbonates and bicarbonates of calcium and magnesium. Noncarbonate hardness is
caused by the nitrates, chlorides, and sulfates of calcium and magnesium. The sum of
both types of hardness is called total hardness.5),6)
Hardnesses in Lake Monroe in 2002 and in Griffy Lake in 1986 were shown in Table 2.
Since the major sources for alkalinity is calcium carbonate (CaCO3), the concentration
of calcium in Griffy Lake in 1986 was similar to alkalinity in 1986 as shown in Table 2
and was not much different from the data in this study, 113.39 mg CaCO3/L. It was
assumed that time difference caused the difference of alkalinity and concentration of
calcium. Total hardness must be higher than the alkalinity or the concentration of
calcium since it includes other species such as magnesium and nitrate.
This alkalinity and hardness in three local water systems seems to be related in southern
Indiana since southern Indiana has considerable limestone, which is composed of over
97% calcium carbonate (CaCO3), deposits underground.
Table 2 The literature data in local water systems 5)
years
lakes
2002
Lake Monroe
1986
Griffy Lake
classes
alkalinity
total hardness
alkalinity
calcium
total hardness
average
(mg CaCO3/L)
27.3
46.4
95
88
248
range
(mg CaCO3/L)
20 - 38
24 - 62
92 - 100
77 - 108
222 - 298
CONCLUSIONS
The alkalinity in Lake Monroe and Griffy Lake in southern Indiana in March 2003 and
Arboretum Pond in Indiana University in March 27, 2003 were 25.89, 113.39, and
70.38 mg CaCO3/L, respectively. The alkalinity in Lake Monroe, which has more
watersheds, was lower than that in Griffy Lake, which has fewer watersheds. The
alkalinity of the Arboretum Pond, which is located in university, was between two lakes.
ACKNOWLEDGMENTS
We thank Dr. Deb. Backhus in the School of Public and Environmental Affairs, Indiana
University at Bloomington for instructing the experiment and Aaron Coby in the School
of Public and Environmental Affairs, Indiana University at Bloomington for helpful
preparation of the experiment.
LITERATURE CITED
1) Clair N. Sawyer and Perry L. McCarly, “Chemistry for Environmental
Engineering” 3rd edition, McGrow Hill, St. Louis, p.364-376 (1978)
2) William W. Jones, Michael Jenson, Eric Jourdain, Sherry Mitchell-Bruker, Lara
Strong, Laura Bieberich, Jeffrey Helmuth, and Tim Kroeker, “Lake Monroe
Diagnostic and Feasibility Study”, School of Public and Environmental Affairs
in Indiana University (1997)
3) Marianne E. Blom and Daniel E. Willard, “A Study of Water quality Data for
Monroe Reservoir, IN, U.S.A.”, School of Public and Environmental Affairs in
Indiana University (1990)
4) “Guidelines for drinking-water quality” 2nd edition, Vol.1. Recommendations,
World Health Organization p.48 (1993)
5) Black & Veatch Corporation, “City of Bloomington Utilities Water System
Master Plan”, chapter7 evaluation of water system improvements (2003)
6) R.G. Wetzel and Gene E. Likens, “Limnological Analyses” 2nd edition,
Springer-verlag, NY (1990)