Hydroxide Carbonate and Bicarbonate Alkalinity UBC Civil by jnyjhtw


ALKALINITY – A measure of the capacity to neutralize acids (also called acid neutralization capacity,
     The alkalinity of natural waters is due primarily to the salts of weak acids, and weak or strong bases,
and substances act as buffers (a substance in solution which resists pH change when acids or bases are
added). Bicarbonate represents the major form of alkalinity (formed from the action of CO2 on basic
materials in soil). Salts of weak acids (borate, silicate, and phosphate, etc.) and organic acids (e.g., humic
acid) may also contribute to the alkalinity of natural waters. In polluted (anaerobic) waters, salts of weak
acids (acetic, propionic, and H2S) also contribute to alkalinity.
     [Carbonate and hydroxide alkalinity – surface water with flourishing algae growth; water through
softening process using lime or lime-soda ash, Boiler water.]
     Three major classes of materials contributing to alkalinity ranked in order of their association with pH
values: 1) hydroxide, 2) carbonate, and 3) bicarbonate. For practical purposes, alkalinity due to other
materials in natural waters is insignificant and is ignored.
     Alkalinity is measured volumetrically by titration using N/50 or 0.020 N H2SO4 and is reported in
terms of equivalent CaCO3. For samples with pH > 8.3, the titration is made in two steps. In the first step
– titration is conducted until the pH is lowered to 8.3 (the point at which phenolphthalein indicator turns
from pink to colorless). In the second step – the titration is conducted until the pH is lowered to about 4.5
(the bromcresol green end point). For samples with pH < 8.3, a single titration is made to a pH of 4.5.
     The choice of pH 8.3 as the end point corresponds to the equivalence point for the conversion of
carbonate ion to bicarbonate ion:
                                      CO32¯ + H+  HCO3¯                                     (18.1)
     The use of pH 4.5 as the end point in the second step of titration corresponds to the equivalence point
for the conversion of bicarbonate ion to carbonic acid:
                                    HCO3¯ + H+  H2CO3                                       (18.2)
     The exact point for this titration depends upon the initial bicarbonate-ion concentration in the sample.
From Table 4.2 and Eq. 4.66, we have:
              pH (bicarbonate equivalence point) = 3.19 - ½ log [HCO3¯]                      (18.3)
     A [HCO3¯] of 0.01 M corresponds to an alkalinity of 500 mg/L as CaCO3, for which the equivalence
point would be 4.10. [A requirement – the carbonic acid or CO2 formed from bicarbonate during titration
not be lost from solution,]
     The actual pH of the stoichiometric end point in alkalinity determination can best be determined by
potentiometric titration [See Fig. 18.1]
     Due to the wide variety of materials (pure water, polluted water, domestic, industrial and agricultural
wastewaters, etc.), it is necessary to explain the methods in some detail to indicate the areas where the
various methods are employed.
Phenolphthalein and Total Alkalinity
     The alkalinity measured to the phenolphthalein end point is called phenolphthalein alkalinity. If the
titration of a sample (originally contains both carbonate and hydroxide) is continued beyond pH 8.3, the
bicarbonate reacts with the acid and is converted to carbonic acid. The reaction is essentially complete
when the pH is lowered to 4.5. The amount of acid required to react with hydroxide, carbonate, and
bicarbonate represents the total alkalinity.

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Hydroxide, Carbonate, and Bicarbonate Alkalinity.
     In water analysis (e.g., softening process), it is desirable to know the kinds and amounts of the various
forms of alkalinity present. Hydroxide, carbonate, and bicarbonate alkalinities may be calculated from the
information given by the titration curves for strong bases and sodium carbonate (Figs. 4.7 and 4.9.) There
are three procedures for these calculations: 1) from alkalinity measurements alone (empirical &
approximate, esp. for samples with pH > 9.), 2) from alkalinity plus pH measurements (sufficiently
accurate, requires an accurate initial pH measurement), and 3) from equilibrium equations. Either the
second or the third procedure is used to calculate hydroxide-, carbonate-, and bicarbonate-ion
concentrations at all pH levels. [Please read Text p.553-557]
Calculation from Alkalinity Measurements Alone
     In this procedure, phenolphthalein and total alkalinities are determined, and from these measurements
the calculation of three types of alkalinity, hydroxide, carbonate, and bicarbonate, are made. This can be
done by assuming (incorrectly) that hydroxide and bicarbonate alkalinity cannot exist together in the same
sample. This permits only five possible situations to be present, which are as follows: (1) hydroxide only,
(2) carbonate only, (3) hydroxide plus carbonate, (4) carbonate plus bicarbonate, and (5) bicarbonate only.
[Reference to Figs. 4.7 and 4.9 will demonstrate that neutralization of hydroxides is complete by the time
enough acid has been added to decrease the pH to 8.3, and that carbonate is exactly one-half neutralized
when the pH has been decreased to the same degree. Upon continuation of the titration to reach a pH of
about 4.5, a negligible amount of acid is needed in the case of the hydroxide, and an amount exactly equal
to that needed to reach pH 8.3 is required for the carbonate. This is the fundamental information needed to
determine which forms of alkalinity are present and the amounts of each.] A graphical representation of
typical titrations obtained with the various combinations of alkalinity is shown in Fig. 18.2.
     Hydroxide Only Samples containing only hydroxide alkalinity have a high pH, usually well above
10. Titration is essentially complete at the phenolphthalein end point. In this case hydroxide alkalinity is
equal to the phenolphthalein alkalinity.
     Carbonate Only Samples containing only carbonate alkalinity have a high pH of 8.5 or higher. The
titration to the phenolphthalein end point is exactly equal to one-half of the total titration. In this case
carbonate alkalinity is equal to the total alkalinity.
     Hydroxide-Carbonate Samples containing hydroxide and carbonate alkalinity have a high pH,
usually well above 10. The titration from the phenolphthalein to the bromcresol green end point represents
one-half of the carbonate alkalinity. Therefore, carbonate alkalinity may be calculated as follows:
              Carbonate alk. = 2 (titration from pH 8.3 to pH 4.5)  1000 / mL sample
and                      Hydroxide alk. = total alk. – carbonate alk.
     Carbonate-Bicarbonate Samples containing carbonate and bicarbonate alkalinity have a pH > 8.3
and usually less than 11. The titration to the phenolphthalein end point represents one-half of the
carbonate. Carbonate alkalinity may be calculated as follows:
                         Carbonate alk. = 2 (titration to pH 8.3)  1000 / mL sample
And                   Bicarbonate alk. = total alk. – carbonate alk.
     Bicarbonate Only Samples containing only bicarbonate alkalinity have a pH of 8.3 or less, usually
less. In this case bicarbonate alkalinity is equal to the total alkalinity.
     The foregoing methods of approximate calculation generally have been superseded by the more
precise methods that will now be described.
Calculation from Alkalinity plus pH Measurements
In this procedure, measurements are made for pH, phenolphthalein, and total alkalinity. This will allow
calculation of hydroxide, carbonate, and bicarbonate alkalinity.
     Hydroxide First, the hydroxide alkalinity is calculated from the pH measurement, using the
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dissociation constant for water:
                                        [OH-] = KW / [H+]                            (18.6)
This calculation requires a precise pH measurement for the determination of [H+]. Since a hydroxide
concentration of 1 mol/L is equivalent to 50,000 mg/L of alkalinity as CaCO3, the relationship in Eq.
(18.6) can be expressed more conveniently as
                                Hydroxide alk. = 50,000  10(pH-pKw)                  (18.7)
At 25°C, pKw = 14.00. However, it varies from 14.94 at 0°C to 13.53 at 40°C. Therefore, it is important
that a temperature measurement be made and the correct pKw be used. The relationship between pH,
temperature, and hydroxide alkalinity is shown graphically in Fig. 18.3. A nomograph is available in
"Standard Methods."
     Carbonate Once the hydroxide alkalinity is determined, use can be made of the principles from the
first procedure to calculate the carbonate and bicarbonate alkalinity. The phenolphthalein alkalinity
represents all the hydroxide alkalinity plus one-half the carbonate alkalinity. Therefore, carbonate
alkalinity may be calculated as follows:
                 Carbonate alk. = 2 (phenol. alk. – hydroxide alk.                       (18.8)
     Bicarbonate The titration from pH 8.3 to pH 4.5 measures the remaining one-half of the carbonate
alkalinity plus all the bicarbonate alkalinity. It is also apparent that the bicarbonate alkalinity represents
the remaining alkalinity after the hydroxide plus carbonate alkalinities are subtracted. From either
standpoint, the bicarbonate alkalinity becomes
              Bicarbonate alk. = total alk. – (carbonate alk. + hydroxide alk.)           (18.9)
Calculation from Equilibrium Equations
The distribution of the various forms of alkalinity can be calculated from equilibrium equations plus a
consideration of electro-neutrality (charge balance) in solution. In order to preserve electro-neutrality, the
sum of the equivalent concentrations of the cations must equal that of the anions. Total alkalinity is a
measure of the equivalent concentration of all cations associated with the alkalinity-producing anions,
except the hydrogen ion. Thus, as shown in Example 4.17, the balance of equivalent concentrations of
alkalinity-associated cations and anions is given by
              [H+] + alkalinity / 50 000 = [HCO3¯] + 2[CO3 2¯] + [OH¯]                (18.10)
The equilibrium equations that must be considered are those for water [Eq. (18.6)] and for the second
ionization of carbonic acid (ignoring activity corrections),
                                    [H+][CO32¯] / [HCO3¯]           (18.11)
From a pH measurement, [H     +] and [OH¯] can be determined, using Eq. (18.6). The only other unknowns

are [HCO3¯] and [CO32¯], and these can be determined from a simultaneous solution of Eqs. (18.10) and
(18.11). The following equations result:
  Carbonate alkalinity (mg/L as CaCO3) =
          {50,000 [(alkalinity/50,000) + [H+] – (Kw / [H+])]} / {1 + ([H+] / 2KA2) } (18.12)

 Bicarbonate alkalinity (mg/L as CaCO3) =
            {50,000[(alkalinity/50,000) + [H+] - (Kw / [H+])] } / {1 + (2KA2 / [H+])} (18 13)

At 25°C, Kw is 10-14 and KA2 is 4.7  10-11. However, these values vary quite radically with
temperature. Also, the activities of the ions vary considerably with, ionic concentration, as indicated in
Sec. 4.3. These corrections are rather tedious; consequently, "Standard Methods" presents nomographs for
the evaluation of carbonate and bicarbonate based on these considerations. The nomographs, as well as
Eqs. (18.12) and (18.13), yield results in terms of alkalinity expressed as CaCO3. At times the actual
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concentrations of carbonate or bicarbonate ion may be desired. Conversions to milligrams per litre of
CO32¯ or HCO3¯ are as follows:
                      mg/L CO32¯ = mg/L carbonate alk.  0.6                          (18.14)
                     mg/L HCO3    ¯ = mg/L bicarbonate alk.  1.22                    (18.15)
   Molar concentrations may be obtained by dividing milligrams per litre by the mole ionic weight in
                       mg/L CO32¯                       mg/L HCO3¯
        [CO32¯ ] = ————— and [HCO3¯] = ——————                                      (18.16)
                        60,000                             61,000
Carbon Dioxide, Alkalinity, and pH Relationships in Natural Waters
    From the equations
                         CO2 + H2O  H2CO3  HCO3¯ + H+                                 (18.17)
                               M(HCO3)2  M      2+ + 2 HCO ¯
                                                              3                          (18.18)
                                  HCO3   ¯  CO 2¯ + H+                                 (18.19)
                           CO3  2¯ + H O  HCO ¯ + OH ¯                                  (18.20)
                                        2           3
The carbon dioxide and the three forms of alkalinity are all part of one system that exists in equilibrium,
since all equations involve HCO3¯. A change in conc. of any member of the system will cause a shift in
the equilibrium; alter the conc. of the other ions, resulting in a change in pH. Conversely, a change in pH
will shift the relationships. [Fig. 18.4]
Application of Alkalinity Data
    Chemical coagulation Chemicals used for coagulation of water and wastewater react with water to
form insoluble hydroxide precipitates. The hydrogen ions released react with the alkalinity of the water.
Thus, the alkalinity acts to buffer the water in a pH range where the coagulant can be effective. Alkalinity
must be present in excess of that destroyed by the acid released by the coagulant for effective and
complete coagulation to occur. (insoluble hydroxide precipitates – jar test dosage determination);
    Water softening (Alkalinity must be considered in calculating the lime and soda-ash requirements in
softening of water by precipitation methods); the alkalinity of softened water is a consideration in terms of
whether such waters meet drinking water standards.
    Corrosion Control Alkalinity is an important parameter involved in corrosion control. It must be
known in order to calculate the Langelier saturation index.
    Buffer Capacity Alkalinity measurements are made as a means of evaluating the buffering capacity
of wastewaters and sludges. They can also be used to assess natural waters’ ability to resist the effects of
acid rain.
    Industrial Wastes Many regulatory agencies prohibit the discharge of wastes containing caustic
(hydroxide) alkalinity to receiving waters. Municipal authorities usually prohibit the discharge of wastes
containing caustic alkalinity to sewers. Alkalinity as well as pH is an important factor in determining the
amenability of wastewaters to biological treatment.
A number of situations are encountered in practice that involve carbon dioxide-alkalinity-pH
relationships and deserve some explanation.
pH Changes during Aeration of Water
   It is common practice to aerate water to remove carbon dioxide, ammonia, and volatile organic
chemicals. Since carbon dioxide is an acidic gas, its removal tends to decrease [H+] and thus raise the pH

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of the water in accordance with Eq. (18.17). Normal air contains about 0.035 percent by volume of carbon
dioxide, which is equivalent to 10-3.46 atm at sea level. The Henry's law constant [see Eq. (2.15)] for
carbon dioxide at 25°C is about 31.6 atm/M; therefore, the equilibrium concentration of carbon dioxide
with air is 10-3.46 (44,000)/31.6 or about 0.48 mg/L. From Eq. (17.7) it can then be calculated that a water
with an alkalinity of 100 mg/L, aerated until in equilibrium with the carbon dioxide in air, would have a
pH of about 8.6. A water with higher alkalinity would tend to have a higher pH upon aeration, and one
with lower alkalinity would tend to have a lower pH.
pH Changes in the Presence of Algal Blooms
    Many surface waters support extensive algal blooms. pH values as high as 10 have been observed in
areas where algae are growing rapidly, particularly in shallow water. Algae use carbon dioxide in their
photosynthetic activity, and this removal is responsible for such high pH conditions. We have seen that
aeration for removal of carbon dioxide tends to increase the pH to between 8 and 9 in waters with
moderate alkalinity. Algae, however, can reduce the free carbon dioxide concentration below its
equilibrium concentration with air and consequently can cause an even greater increase in pH. As the pH
increases, the alkalinity forms change, with the result that carbon dioxide can also be extracted for algal
growth both from bicarbonates and from carbonates in accordance with the following equilibrium
                              2HCO3 ¯  CO32¯ + H2O + CO2             (18.21)
                              CO32¯ + H2O  2OH¯ + CO2                 (18.22)
Thus, the removal of carbon dioxide by algae tends to cause a shift in the forms of alkalinity present from
bicarbonate to carbonate, and from carbonate to hydroxide. It should be noted that during these changes
the total alkalinity remains constant unless removal results through precipitation of carbonate salts such as
CaCO3(s). Algae can continue to extract carbon dioxide from water until an inhibitory pH is reached,
which is usually in the range of pH 10 to 11.
    During the dark hours of the day, algae produce rather than consume carbon dioxide. This is because
their respiratory processes in darkness exceed their photosynthetic processes. This carbon dioxide
production has the opposite effect and tends to reduce the pH. Diurnal variations in pH due to algal
photosynthesis and respiration are common in surface waters..
    In natural waters containing appreciable amounts of Ca2+, calcium carbonate precipitates when the
carbonate-ion concentration, according to Eq. (18.21), becomes great enough so that the CaCO3 solubility
product is exceeded:
                                 Ca2+ + CO32¯  CaCO3(s)            (18.23)
This precipitation usually happens before pH levels have exceeded 10, and it places a ceiling over the pH
values obtainable. The calcium carbonate precipitated as a result of removal of carbon dioxide through
algal action produces the marl deposits in lakes. Marl deposits are the precursors of limestone.
Alkalinity of Boiler Waters

    Boiler waters contain both carbonate and hydroxide alkalinity. Both are derived from bicarbonate in
the feed water. Carbon dioxide is insoluble in boiling water and so is removed with the steam. This causes
an increase in pH and a shift in alkalinity forms from bicarbonate to carbonate and from carbonate to
hydroxide, as indicated in Eqs. (18.21) and (18.22). Under these extreme conditions, pH levels in excess
of 11.0 are often obtained. If Ca2+ levels are high, precipitation of CaCO3(s) may occur.

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