· The alkalinity of a water is a measure of its capacity to neutralize acids. It is the sum of all titratable bases. Conversely the acidity of a water is a measure of its capacity to neutralize bases. It is the sum of all titratable acids.

· Neither alkalinity nor acidity have any known adverse health effects. Although highly acidic or alkaline waters are frequently considered unpalatable. However, a knowledge of these parameters may be important in many situations.

· The alkalinity of a body of water provides information about how sensitive that water body will be to acid inputs such as acid rain.

· Turbidity is frequently removed from drinking water by coagulation and flocculation. This process releases H+ into the water. Alkalinity must be present in excess of that destroyed by the H+ released for effective and complete coagulation to occur.

· Hard waters are frequently softened by precipitation methods. The alkalinity of the water must be known in order to calculate the lime (Ca(OH)₂) and soda ash (Na₂CO₃) requirements for precipitation.

· Alkalinity is important in corrosion control.

· Finally, HCO₃- and CO₃^2- may complex with other elements and compounds altering their transport and fate in the environment.

· The carbonate system is the predominant source of alkalinity and acidity in most natural waters. The reactions involved are

,                                                                        (1)

,    and                                                            (2)

.                                                                          (3)

Here H₂CO₃* represents both dissolved CO₂ and true H₂CO₃, as these two species are very difficult to distinguish analytically.

· The equilibrium constants for these reactions are

     ,  (Note that this is an example of Henry's law)                             (1)

     , and                                                                                             (2)

     , respectively.                                                                               (3)

· Acidity may also include mineral acidity, or small amounts of organic acids such as humic and fulvic acids, or inorganic acids such as sulfuric and nitric acids.

· Carbon dioxide acidity alone will not depress the pH of a water below about 4.0. If the pH of a water falls below 4.0 mineral acidity must be present.

· Small amounts of hydroxide, borates, silicates, and phosphates may also contribute to alkalinity. These are collectively referred to as caustic alkalinity. Carbonate alkalinity alone will not raise the pH of a water above about 10.3. If the pH of a water is above 10.3 caustic alkalinity must be present.

· Total alkalinity may be defined by the equation,

            TA = 2[CO₃^2-] + [HCO₃-] + [OH-] - [H+],  if we assume caustic alkalinity is insignificant.

· Alkalinity and acidity are measured volumetrically by titration. An alkalinity/acidity titration is illustrated in the following figure in conjunction with a pC/pH diagram.

· Notice that at pH 10.7 the [HCO₃-] equals the [OH-]. This is called an equivalence point. This is the endpoint for caustic alkalinity and total acidity.

· Notice that at pH 8.3 the [H₂CO₃] equals the [CO₃^2-]. This is the endpoint for carbonate alkalinity and CO₂ acidity. In the alkalinity titration virtually all CO₃^2- has reacted (thus the term carbonate alkalinity) and half of the HCO₃- has reacted. Carbonate alkalinity is also known as phenolphthalein alkalinity as this is the color indicator usually used for this endpoint. In the acidity titration virtually all of the H₂CO₃ has reacted (thus the term CO₂ acidity) and half of the HCO₃- has reacted.

· At pH 4.5 the [H+] equals the [HCO₃-]. This is the endpoint for mineral acidity and total alkalinity.

· Alkalinity is by convention reported as g/L or mg/mL CaCO₃. This is an arbitrary convention. However, the dissolution of CaCO₃ is the source of much alkalinity.

American Public Health Association, Standard Methods for the Examination of Water and Waste Water, Washington D.C., 1992.

Sawyer, C.N., P.L. McCarty, G.F. Parkin, Chemistry for Environmental Engineering, MCGraw-Hill, New York, 1994.

Snoeyink, V.L. and D. Jenkins, Water Chemistry, John Wiley & Sons, New York, 1980.

Stumm, W. and J.J. Morgan, Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters, John Wiley & Sons, New York, 1996.

1. Accurately measure 200 mL of your sample and quantitatively transfer the sample to a 250 mL beaker.

2. Calibrate a pH meter using pH 4, 7, and 10 buffers with stirring.

3. Measure the pH of your sample to the nearest 0.01 pH unit with stirring.

4. Fill a 50 mL burette with standardized 0.02 N H₂SO4 or HCl.

5. If the pH of your sample is greater than 8.3 add approx. 1.5 mL of phenolphthalein indicator to the sample. The solution should turn pink.

6. Position your sample on the stir plate beneath the buret.

7. Slowly titrate the sample. Record both the volume of acid added and the pH. The first endpoint should be reached around pH 8.4. The solution should change from pink to clear, and the pH should be changing rapidly with small additions of acid.

8. After reaching the phenolphthalein endpoint (or if the pH of your sample was less than 8.3), add approx. 1.5 mL of bromcresol green indicator. The solution should turn blue.

9. Slowly titrate, again recording both volume of acid added and pH until the bormcresol green endpoint is reached. This should occur at about pH 4.5. The solution should turn from blue to green. The pH should again be changing rapidly with small additions of acid.

10. Plot the pH versus the mL of acid added. Indicate the endpoints of the titration on the graph.

11. Calculate both the carbonate (or phenolphthalein) alkalinity and the total alkalinity for the sample.

For example,

Alkalinity, mg CaCO₃/mL

= 

where A = mL standard acid used, N = normality of standard acid, MWCaCO₃= the molecular weight of CaCO₃ and EquCaCO₃ = the equivalent weight of CaCO₃ = 2. Acidity and Alkalinity