What is Alkalinity?
"The Protector of the Stream"
Alkalinity is the water's capacity to resist changes in pH that would make the water more acidic. It is also the protector of your health and piping when it comes to drinking water. This capacity is commonly known as "buffering capacity."
For example, if you add the same weak acid solution to two vials of water - both with a pH of 7, but one with no buffering power (e.g. zero alkalinity) and the other with buffering power (e.g. an alkalinity of 50 mg/l), - the pH of the zero alkalinity water will immediately drop while the pH of the buffered water will change very little or not at all. The pH of the buffered solution would change when the buffering capacity of the solution is overloaded or used up. Pure water (no alkalinity) has a pH of 7 but let that pure water be exposed to the atmosphere and the pH will drop to ~ 5.5 because carbon dioxide in the air will dissolve in the water and create carbonic acid.
This is why I like to call Alkalinity the Protector of the Stream.
Alkalinity refers to the capability of water to neutralize acid. This is really an expression
of buffering capacity. A buffer is a solution to which an acid can be added (within limits)
without changing the concentration of available H+ ions (without changing the pH) appreciably.
It essentially absorbs the excess H+ ions and protects the water body from fluctuations in pH.
In most natural water bodies in Pennsylvania, the buffering system is carbonate-bicarbonate
(H2CO3, HCO3, and CO3).
The alkalinity of natural water is determined by the soil and bedrock through which it passes.
The main sources for natural alkalinity are rocks which contain carbonate, bicarbonate, and
hydroxide compounds. Borates, silicates, and phosphates may also contribute to alkalinity.
Limestone is rich in carbonates, so waters flowing through limestone regions or bedrock containing
carbonates generally have high alkalinity, hence, good buffering capacity. Conversely, areas
rich in granite and some conglomerates and sandstones may have low alkalinity and, therefore,
poor buffering capacity.
The presence of calcium carbonate or other compounds such as magnesium carbonate contribute
carbonate ions to the buffering system. Alkalinity is often related to hardness because
the main source of alkalinity is usually from carbonate rocks (limestone) which are mostly
CaCO3. If CaCO3 actually accounts for most of the alkalinity, hardness in
CaCO3 is equal to alkalinity. Since hard water contains metal carbonates
(mostly CaCO3) it is high in alkalinity. Conversely, unless carbonate is associated
with sodium or potassium, which don't contribute to hardness, soft water usually has low
alkalinity and little buffering capacity. So, generally, soft water is much more susceptible
to fluctuations in pH from acid rains or acid contamination.
How Does Alkalinity Affect Aquatic Life?
Alkalinity is important for fish and aquatic life because it protects or buffers against rapid pH changes. Living organisms, especially aquatic life, function best in a pH range of 6.0 to 9.0. Alkalinity is a measure of how much acid can be added to a liquid without causing a large change in pH. Higher alkalinity levels in surface waters will buffer acid rain and other acid wastes and prevent pH changes that are harmful to aquatic life.
Acid shock may occur in spring when acid snows melt, there are thunderstorms, natural discharges of tannic waters, "acid rain", acidic dryfall, and when other acidic discharges enter the stream. If increasing amounts of acids are added to a body of water, the water's buffering capacity is consumed. If additional buffering material can be obtained from surrounding soils and rocks, the alkalinity level may eventually be restored. However, a temporary loss of buffering capacity can permit pH levels to drop to those harmful to life in the water.
The pH of water does not fall evenly as acid contamination proceeds. The natural buffering materials in water slow the decline of pH to around 6.0. This gradual decline is followed by a rapid pH drop as the bicarbonate buffering capacity is used up. At a pH of 5.5, only very weak buffering materials remain and pH drops rapidly with additional acid. Sensitive species and immature animals are affected first. As food species disappear, even larger, resistant animals are affected.
For the protection of aquatic life, the buffering capacity should be at least 20 mg/L. If alkalinity is naturally low (less than 20 mg/L), there can be no greater than a 25% reduction in alkalinity.
Testing Methodology
1. Measure out 50 mL of sample (recorded as V1, in milliliters) into a clean test container.
2. Add a few drops of phenolphthalein indicator, if the solution turns pink the pH is above 8.3.
3. Add 0.02 N sulfuric acid (H2SO4) one drop at a time and mix. Continue adding until the pink color is discharged or disappears. Record the volume of acid used as A1 (milliliters).
4. Add a powder pillow or a few drops of methyl orange indicator; the solution will turn yellow.
5. Add 0.02 N sulfuric acid until the solution appears more orange and record
the volume of the acid used as A2 (milliliters).
Calculations
Phenolphthalein Alkalinity (P) = [(A1 * 0.02 * 50 * 1000)/ V1] = mg CaCO3/L
Methyl Orange Alkalinity (T) = (A2 * 0.02 * 50 *1000)/ V1) = mg CaCO3/L
V1- Volume of sample (ml) = 50 ml
A1 - Volume of 0.02 N sulfuric acid (20 drops); each drop equals 0.05 ml, therefore, 20 drops
equals 1 ml
A2 - Volume of 0.02 N sulfuric acid (45 drops); each drop equals 0.05 ml, therefore, 20 drops
equals 2.25 ml
Phenolphthalein Alkalinity (P) = (1 * 0.02 * 50 * 1000)/ 50) =20 mg CaCO3/L
Methyl Orange Alkalinity (T) = (2.25 * 0.02 * 50 *1000)/ 50) = 45 mg CaCO3/L
Total Alkalinity = P + T = 65 mg/L
A potentiometric titration is typically taken to an end-point pH of 4.5. The amount of acid
required to reach a pH of 4.5 is expressed in milliliters. The carbonate ions
(CO3=) neutralize the acid in this reaction and show the buffering capacity
of the sample. The amount of acid used will indicate the amount of carbonate
(CO3=) involved in the reaction. This then is expressed as mg of
CaCO3/L even though part of the alkalinity may come from MgCO3,
Na2CO3, or K2CO3.
