Acidity And Alkalinity

Background

The alkalinity of water is its acid-neutralizing capacity. The acidity of water is its base-neutralizing capacity. Both parameters are related to the buffering capacity of water (the ability to resist changes in pH when an acid or base is added). Water with high alkalinity can neutralize a large quantity of acid without large changes in pH; on the other hand, water with high acidity can neutralize a large quantity of base without large changes in pH.

Acidity

Acidity is determined by measuring how much standard base must be added to raise the pH to a specified value. Acidity is a net effect of the presence of several constituents, including dissolved carbon dioxide, dissolved multivalent metal ions, strong mineral acids such as sulfuric, nitric, and hydrochloric acids, and weak organic acids such as acetic acid. Dissolved carbon dioxide (CO2) is the main source of acidity in unpolluted waters. Acidity from sources other than dissolved CO2 is not commonly encountered in unpolluted natural waters and is often an indicator of pollution.

Titrating an acidic water sample with base to pH 8.3 measures phenolphthalein* acidity or total acidity. Total acidity measures the neutralizing effects of essentially all the acid species present, both strong and weak.

Titrating with base to pH 3.7 measures methyl orange* acidity Methyl orange acidity primarily measures acidity due to dissolved carbon dioxide and other weak acids that are present.

Alkalinity

In natural waters that are not highly polluted, alkalinity is more commonly found than acidity. Alkalinity is often a good indicator of the total dissolved inorganic carbon (bicarbonate and carbonate anions) present. All unpolluted natural waters are expected to have some degree of alkalinity. Since all natural waters contain dissolved carbon dioxide, they all will have some degree of alkalinity contributed by carbonate species — unless acidic pollutants would have consumed the alkalinity. It is not unusual for alkalinity to range from 0 to 750 mg/L as CaCO3. For surface waters, alkalinity levels less than 30 mg/L are considered low, and levels greater than 250 mg/L are considered high. Average values for rivers are around 100-150 mg/L. Alkalinity in environmental waters is beneficial because it minimizes pH changes, reduces the toxicity of many metals by forming complexes with them, and provides nutrient carbon for aquatic plants.

Alkalinity is determined by measuring how much standard acid must be added to a given amount of water in order to lower the pH to a specified value. Like acidity, alkalinity is a net effect of the presence of several constituents, but the most important are the bicarbonate (HCO3-), carbonate (CO32-), and hydroxyl (OH) anions. Alkalinity is often taken as an indicator for the concentration of these constituents. There are other, usually minor, contributors to alkalinity, such as ammonia, phosphates, borates, silicates, and other basic substances.

Titrating a basic water sample with acid to pH 8.3 measures phenolphthalein alkalinity. Phenolphthalein alkalinity primarily measures the amount of carbonate ion (CO32-) present. Titrating with acid to pH 3.7 measures methyl orange alkalinity or total alkalinity Total alkalinity measures the neutralizing effects of essentially all the bases present.

Because alkalinity is a property caused by several constituents, some convention must be used for reporting it quantitatively as a concentration. The usual convention is to express alkalinity as ppm or mg/L of calcium carbonate (CaCO3). This is done by calculating how much CaCO3 would be neutralized by the same amount of acid as was used in titrating the water sample when measuring

* Phenolphthalein and methyl orange are pH-indicator dyes that change color at pH 8.3 and 3.7, respectively. Copyright © 2000 CRC Press, LLC

either phenolphthalein or methyl orange alkalinity. Whether it is present or not, CaCO3 is used as a proxy for all the base species that are actually present in the water. The alkalinity value is equivalent to the mg/L ofCaCO3 that would neutralize the same amount ofacid as does the actual water sample.

Importance of Alkalinity

Alkalinity is important to fish and other aquatic life because it buffers both natural and human-induced pH changes. The chemical species that cause alkalinity, such as carbonate, bicarbonate, hydroxyl, and phosphate ions, can form chemical complexes with many toxic heavy metal ions, often reducing their toxicity. Water with high alkalinity generally has a high concentration of dissolved inorganic carbon (in the form of HCO3- and CO32-) which can be converted to biomass by photosynthesis. A minimum alkalinity of 20 mg/L as CaCO3 is recommended for environmental waters and levels between 25 and 400 mg/L are generally beneficial for aquatic life. More productive waterfowl habitats correlate with increased alkalinity above 25 mg/L as CaCO3.

Criteria and Standards for Alkalinity

Naturally occurring levels of alkalinity reaching at least 400 mg/L as CaCO3 are not considered a health hazard. EPA guidelines recommend a minimum alkalinity level of 20 mg/L as CaCO3, and that natural background alkalinity is not reduced by more than 25% by any discharge. For waters where the natural level is less than 20 mg/L, alkalinity should not be further reduced. Changes from natural alkalinity levels should be kept to a minimum. The volume of sample required for alkalinity analysis is 100 mL.

Rules of Thumb

1. Alkalinity is the mg/L of CaCO3 that would neutralize the same amount of acid as does the actual water sample.

2. Phenolphthalein alkalinity (titration with acid to pH 8.3) measures the amount of carbonate ion (CO32-) present.

3. Total or methyl orange alkalinity (titration with acid to pH 3.7) measures the neutralizing effects of essentially all the bases present.

4. Surface and groundwaters draining carbonate mineral formations become more alkaline due to dissolved minerals.

5. High alkalinity can partially mitigate the toxic effects of heavy metals to aquatic life.

6. Alkalinity greater than 25 mg/L CaCO3 is beneficial to water quality.

7. Surface waters without carbonate buffering may be more acidic than pH 5.7 (the value established by equilibration of dissolved CO2 with CO2 in the atmosphere) because of water reactions with metals and organic substances, biochemical reactions, and acid rain.

Calculating Alkalinity

Although alkalinity is usually determined by titration, the part due to carbonate species (carbonate alkalinity) is readily calculated from a measurement of pH, bicarbonate and/or carbonate. Carbonate alkalinity is equal to the sum of the concentrations of bicarbonate and carbonate ions, expressed as the equivalent concentration of CaCO3.

Example 3.3

A groundwater sample contains 300 mg/L of bicarbonate at pH = 10.0. Calculate the carbonate alkalinity as CaCO3.

Answer:

1. Use the measured values of bicarbonate and pH, with Figure 3.2, to determine the value of CO32-. At pH = 10.0, total carbonate is about 73% bicarbonate ion and 27% carbonate ion. Although these percentages are related to moles/L rather than mg/L, the molecular weights of bicarbonate and carbonate ions differ by only about 1.7%; therefore, mg/L can be used in the calculation without significant error.

0.73

CO32- = 0.27 x 411 = 111 mg/L, or alternatively, 411 - 300 = 111 mg/L.

2. Determine the equivalent weights of HCO3-, CO32-, and CaCO3.

molecular or atomic weight eq. wt. =

magnitude of ionic charge or oxidation number eq. wt. of HCO3- = 6102 = 61.0.

eq. wt. of CO32- = 6101 = 30.0. eq. wt. of CaCO3 = 1O0°9 = 50.0.

3. Determine the multiplying factors to obtain the equivalent concentration of CaCO3.

Multiplying factor of HCO3- as CaCO3 = eq. wt of CaC° = 500 = 0.820. 6 3 3 eq. wt. of HCO3 - 61.0

Multiplying factor of CO32- as CaCO3 = —-^ =- = 1.667.

4. Use the multiplying factors and concentrations to calculate the carbonate alkalinity, expressed as mg/L of CaCO3.

Carbonate alk. (as CaCO3) = 0.820 [HCO3-, mg/L] + 1.667 [CO32-, mg/L]. (3.13)

Carbonate alk. = 0.820 [300 mg/L] + 1.667 [111 mg/L] = 431 mg/L CaCO3.

Equation 3.13 may be used to calculate carbonate alkalinity whenever pH and either bicarbonate or carbonate concentrations are known.

Calculating Changes in Alkalinity, Carbonate, and pH

A detailed calculation of how pH, total carbonate, and total alkalinity are related to one another is moderately complicated because of the three simultaneous carbonate equilibria reactions, Equations 3.9-3.11. However, the relations can be conveniently plotted on a total alkalinity/pH/total carbonate graph, also called a Deffeyes diagram, or capacity diagram (see Figures 3.3 and 3.4). Details of the construction of the diagrams may be found in Stumm and Morgan (1996) and Deffeyes (1965).

In a total alkalinity/pH/total carbonate graph shown in Figure 3.3, a vertical line represents adding strong base or acid without changing the total carbonate (CT). The added base or acid changes the pH and, therefore, shifts the carbonate equilibrium, but does not add or remove any carbonate. The amount of strong base or acid in meq/L equals the vertical distance on the graph. You can see from Figure 3.3 that if the total carbonate is small, the system is poorly buffered, so a little base or acid makes large changes in pH. If total carbonate is large, the system buffering capacity is similarly large and it takes much more base or acid for the same pH change.

A horizontal line represents changing total carbonate, generally by adding or losing CO2, without changing alkalinity. For alkalinity to remain constant when total carbonate changes, the pH must also change. Changes caused by adding bicarbonate or from simple dilution are indicated in the figure.

Figure 3.4 is a total acidity/pH/total carbonate graph. Note that changes in composition, caused by adding or removing carbon dioxide and carbonate, are indicated by different movement vectors in the acidity and alkalinity graphs. The examples below illustrate the uses of the diagrams.

Example 3.4

Designers of a wastewater treatment facility for a meat rendering plant planned to control ammonia concentrations in the wastewater by raising its pH to 11, in order to convert about 90% of the ammonia to the volatile form. The wastewater would then be passed through an air-stripping tower to transfer the ammonia to the atmosphere. Average initial conditions for alkalinity and pH in the wastewater were expected to be about 0.5 meq/L and 6.0, respectively.

In the preliminary design plan, four options for increasing the pH were considered:

1. Raise the pH by adding NaOH, a strong base.

2. Raise the pH by adding calcium carbonate, CaCO3, in the form of limestone.

3. Raise the pH by adding sodium bicarbonate, NaHCO3.

4. Raise the pH by removing CO2, perhaps by aeration.

Addition of NaOH

In Figure 3.3, we find that the intersection of pH = 6.0 and alkalinity = 0.5 meq/L occurs at total carbonate = 0.0015 mol/L, point A. Assuming that no CO2 is lost to the atmosphere, addition of the strong base NaOH represents a vertical displacement upward from point A. Enough NaOH must be added to intersect with the pH = 11.0 contour at point B. In Figure 3.3, the vertical line between points A and B has a length of about 3.3 meq/L. Thus, The quantity of NaOH needed to change the pH from 6.0 to 11.0 is 3.3 meq/L (132 mg/L).

Addition of CaCO3

Addition of CaCO3 is represented by a line of slope +2 from point A. The carbonate addition line rises by 2 meq/L of alkalinity for each increase of 1 mol/L of total carbonate (because one mole of carbonate = 2 equivalents). Notice in Figure 3.3 that the slope of the pH = 11.0 contour is very nearly 2. The CaCO3 addition vector and the pH = 11.0 contour are nearly parallel. Therefore, a very large quantity of CaCO3 would be needed, making this method impractical.

Diagram Alkalinity

FIGURE 3.3 Total alkalinity-pH-total carbonate diagram (Deffeyes diagram): In this figure, the relationships among total alkalinity, pH, and total carbonate are shown. If any two of these quantities are known, the third may be determined from the plot. The composition changes indicated in the figure refer to Example 3.4.

FIGURE 3.3 Total alkalinity-pH-total carbonate diagram (Deffeyes diagram): In this figure, the relationships among total alkalinity, pH, and total carbonate are shown. If any two of these quantities are known, the third may be determined from the plot. The composition changes indicated in the figure refer to Example 3.4.

Addition of NaHCO3

Addition of NaHCO3 is represented by a line of slope +1 from point A. Although this vector is not shown in Figure 3.3, it is evident it cannot cross the pH = 11 contour. Therefore, this method will not work.

Removing CO2

Removal of CO2 is represented by a horizontal displacement to the left. Loss or gain of CO2 does not affect the alkalinity. Note that if CO2 is removed, total carbonate is decreased correspondingly. However, pH and [OH] also increase correspondingly, resulting in no net change in alkalinity. We see from Figure 3.3 that removal of CO2 to the point of zero total carbonate cannot achieve pH = 11.0. Therefore, this method also will not work.

Of the four potential methods considered for raising the wastewater pH to 11.0, only addition of NaOH is useful.

Example 3.5

A large excavation at an abandoned mine site has filled with water. Because pyrite minerals were exposed in the pit, the water is acidic with pH = 3.2. The acidity was measured at 3.5 meq/L. Because the pit overflows into a stream during heavy rains, managers of the site must meet the conditions of a discharge permit, which include a requirement that pH of the overflow water be between 6.0 and 9.0. The site managers decide to treat the water to pH = 7.0 to provide a safety margin. Use Figure 3.4 to evaluate the same options for raising the pH as were considered in Example 3.4.

Addition of NaOH

In Figure 3.4, we find that the intersection of pH = 3.2 and acidity = 3.5 meq/L occurs at about total carbonate = 0.0014 mol/L, point A. Assuming that no CO2 is lost to the atmosphere, addition of the strong base NaOH represents a vertical displacement downward from point A to point C. Enough NaOH must be added to intersect with the pH = 7.0 contour. The vertical line between points A and C has a length of about 1.8 meq/L. Thus, The quantity of NaOH needed to change the pH from 3.0 to 7.0 is 1.8 meq/L (72 mg/L).

Addition of CaCO3

In the acidity diagram, addition of CaCO3 is represented by a horizontal line to the right. In Figure 3.4, the CaCO3 addition line intersects the pH = 7.0 contour at point B, where total carbonate = 0.0030 mol/L. Therefore, the quantity of CaCO3 required to reach pH = 7.0 is 0.0030 - 0.0014 = 0.0016 mol/L (160 mg/L).

Addition of NaHCO3

The addition of NaHCO3 is represented by a line of slope +1 (the vector upward to the right from point A in Figure 3.4). Notice that the slope of the pH = 7.0 contour is just a little greater than +1. The NaHCO3 addition vector and the pH = 7.0 contour are nearly parallel. Therefore, a very large quantity of NaHCO3 would be needed, making this method impractical.

Removing CO2

In the acidity diagram, the removal of CO2 is represented by a line downward to the left with slope 2. We see from Figure 3.4 that removal of CO2 to the point of zero total carbonate cannot achieve pH = 7.0. Therefore, this method will not work.

Of the four potential methods considered for raising the wastewater pH to 7.0, addition of either NaOH or CaCO3 will work. The choice will be based on other considerations, such as costs or availability.

+2 -1

Responses

  • jay
    How much total acidity and total alkalinity can drinking water have?
    9 months ago
  • Prospero Goodbody
    How inorganic pollutants,acids and alkalinity affect the water quality?
    8 months ago
  • hessu
    What is the alkalinity in 300 mg/l CaCO3?
    3 months ago

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