The alkalinity of water is a measure of how much acid it can neutralize. If any changes are made to the water that could raise or lower the pH value, alkalinity acts as a buffer, protecting the water and its life forms from sudden shifts in pH. This ability to neutralize acid, or H+ ions, is particularly important in regions affected by acid rain.


In the diagram below, for example, the lake on the right has low alkalinity. When acid rain falls, it is not neutralized, so the pH of the water decreases. This drop in the pH level can harm or even kill some of the aquatic organisms in the lake. The lake on the left, however, has high alkalinity. When acid rain falls in this lake, the acid is partially neutralized and the pH of the water remains fairly constant. In this way, a high alkalinity level helps maintain the health of the water and the organisms that live there.



Alkalinity should not be confused with pH. The pH of a solution is a measure of the concentration of acid, or H+ ions, in the water. Alkalinity is a measure of the water's capacity to neutralize an acid, or H+ ions, thereby keeping the pH at a fairly constant level.


The alkalinity of surface water is primarily due to the presence of hydroxide (OH-), carbonate (CO32-), and bicarbonate (HCO3-) ions. These ions (bases) react with H+ ions (acid) by means of the following acid-base reactions:

OH- + H+ H2O
CO32- + H+   HCO3-
HCO3- + H+ [H2CO3]   CO2 + H2O

Most alkalinity in surface water comes from calcium carbonate, CaCO3, being leached from rocks and soil. This process is enhanced if the rocks and soil have been broken up for any reason, such as mining or urban development. Limestone contains especially high levels of calcium carbonate.


Alkalinity is significant in the treatment of wastewater and drinking water, because it will influence treatment processes such as anaerobic digestion. Water may also be unsuitable for use in irrigation if the alkalinity level in the water is higher than the natural level of alkalinity in the soil.


Expected Levels

Alkalinity is reported in units of mg/L CaCO3, because the carbonate ion, CO32-, is its primary constituent. Alkalinity levels will vary across the country. Some sample data are shown in Table 1. In general, water in the eastern half of the United States will have a higher alkalinity than water in the west because of a higher occurrence of limestone. Areas in the extreme northeast that have had the limestone scoured away by glacial action will often have a lower alkalinity.


Table 1:  Alkalinity of Selected Rivers



(mg/L CaCO3)


Missouri River, St. Joseph, MO



Missouri River, Garrison Dam, ND



Cataloochee Creek, Cataloochee, NC



Columbia River, Northport, WA



Merrimack River, Lowell, MA





Alkalinity is measured by means of an acid-base reaction. The basic ions naturally present in the sample water - hydroxide (OH-), carbonate (CO32-), and bicarbonate (HCO3-) ions - are made to react with added acid. Acid-base reactions are often carried out as titrations. A titration is any reaction in which the amount of one or more of the reactants is monitored volumetrically. The object of an acid-base titration is usually to determine the point at which exact neutralization of the acid by the base occurs. The neutralization point is known as the equivalence point, since at this point an equivalency between the acid and the base has been achieved. Titrations involving strong acids and strong bases have equivalence points at or near pH 7. Titrations between a strong acid and a weak base (such as CO32- or HCO3-, the predominant ions in this titration), on the other hand, will have equivalence points at pH values significantly below 7. You will find that the equivalence point in alkalinity titrations will be at a pH of approximately 4.5, but will vary slightly depending on the chemical composition of the water. The volume of sulfuric acid added at the equivalence point of the titration is then used to calculate the alkalinity of the water.

In this experiment you will titrate water from BCC's wetland with sulfuric acid. The acid will be added in incremental steps, and you will note how the pH changes with each addition of acid. The pH is monitored using a computer interfaced pH electrode. The glass bulb at the tip of the electrode measures the H+ ion activity and produces a corresponding voltage. This voltage varies linearly with the log of the H+ ion activity. Because pH = -log[H+], the voltage produced varies linearly with pH.

At first, when you have added a small amount of acid and are far from neutralization, the pH will decrease slowly. This snail's pace will be maintained until you are quite close to the neutralization (equivalence) point at which time the pH will change dramatically. After neutralization has occurred, the slow pace of decreasing pH will be resumed and you can stop the titration.

If you plot the pH versus the volume of acid that was added, and draw a smooth curve, a graph similar to the one shown in Figure 1 will result. The inflexion point in the graph (the steepest drop in pH) gives the equivalence point of the titration. If, from the inflexion point, you drop a perpendicular down to the volume of acid axis, you can determine the volume of acid (mL) needed to neutralize the basic ions in your sample.






Figure 1. Alkalinity titration curve





The main reaction occurring in this titration is

H2SO4 + CaCO3 H2O + CO2 + CaSO4

Since the reaction of sulfuric acid with calcium carbonate has a stoichiometry of one to one, the following equality is valid at the equivalence point:

# moles of H2SO4 = # moles of CaCO3

Remembering that molarity (M) equals moles per liter, we can rewrite the above equality as follows:

MH2SO4 x VH2SO4 = MCaCO3 x VCaCO3


From the above equation we can calculate MCaCO3

MCaCO3 (mol CaCO3/L) = (MH2SO4 x VH2SO4)/VCaCO3


(where VH2SO4 = value read from burette, MH2SO4=0.0100 mol/L, VCaCO3=100.0 mL)


Alkalinity (in mg/L CaCO3) can then be calculated from MCaCO3 (mol CaCO3/L) and the molar mass of CaCO3 (100 g/mol) as follows:


Alkalinity = (mol CaCO3/L) x (100 g / 1 mol CaCO3) x (1,000 mg / 1 g) = mg/L CaCO3