Alkalinity and OLI

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What does OLI mean when it discusses Alkalinity? Alkalinity is a frequently measured and reported quality of many waters. Stumm and Morgan define alkalinity as:

“Acidimetric or alkalimetric titrations of carbonate-bearing water to the appropriate end points represent operations procedures for determining alkalinity and acidity, that is, the equivalent sum of the bases that are titratable with strong acid and the equivalent sum of the acids that are titratable with strong base. Alkalinity and acidity are then the capacity factors that represent, respectively, the acid- and base-neutralizing capacities … of an aqueous system. For solutions that contain no protolysis system other than that of aqueous carbonate, alkalinity is a measure of the quantity of strong acid per liter required to attain a pH equal to that of a total concentration (molar) solution of H2CO3. Alternatively, acidity is a measure of the quantity per liter of strong base required to attain a pH equal to that of a total concentration (molar) solution of Na2CO31

The key to this statement involves the fact that many users think that the alkalinity is the concentration of various forms of carbonate ion. This would be true of other acid or base systems were not present in solution. Even simple ions such as sodium and magnesium may affect the free carbonate in solution and have markedly different alkalinities.

OLI considers alkalinity to be the total base capacity of the brine. We will us a titration to determine the alkalinity exactly. We will now show some examples featuring the OLI/LabAnalyzer™ program.

We will consider a simple brine with the following concentration

 1“Aquatic Chemistry. An Introduction Emphasizing Chemical Equilibria in Natural Waters”. Werner Stumm and James J. Morgan. John-Wiley & Sons, New York. 1981 p 185

Table 7 Updated.jpg

Figure 1 Brine composition.

The user would suspect that the alkalinity would be the same as the bicarbonate concentration of 375 mg/L. The electrically neutral pH of this brine is:

Summary 1.jpg

Figure 2 the pH is 7.68

We will use HCl to titrate the brine to the standard end point pH of 4.5.

The amount of HCl added to bring the pH down to 4.5 is:

Table 3.jpg

Figure 3 215.4 mg/L of HCl required.

This amount of HCl needs to be converted to equivalents of bicarbonate ion for reporting purposes.

Formula 1.jpg

So there is actually less alkalinity than the input concentration of bicarbonate ion would indicate. The reason for this is that some of the bicarbonate ion is tied up in the form of a complex, NaHCO3o and is not available to the alkalinity titration.

What would happen if organics acids were present in the brine? The following brine concentration has 150 mg/L of acetic acid added:

Table 8.jpg

Figure 4 Added acetic acid

The alkalinity was determined as before to an end point pH of 4.5.

Table 5.jpg

Figure 5 164.8 mg/L of HCl required.

This corresponds to an alkalinity of 275.8 mg/L as HCO3-1. This is considerably less than the original bicarbonate ion concentration of 375 mg/L. The problem occurs in that the acetic acid is also reacting with the HCl. At a pH of 4.5 (the titration end point) not all of the acetic acid has reacted. The reason is that the dissociation constant for acetic acid is near the end point pH.

CH3COOH = CH3COO-1 + H+1 pKa ~ 4.7

Most procedures state that if organic acids are present the pH end point must be lowered to react both the carbonate and the acid. We used a lower pH end point of 3.0 and obtained this result:

Table 6.jpg

Figure 6 263.6 mg/L of HCl to reach the end point pH of 3.0

This corresponds to an alkalinity of 441 mg/L as HCO3-1. This value is greater than the bicarbonate ion concentration that was entered. This means that the organic acid is also a source of alkalinity.

It is true that any organic or inorganic species that has significant acid/base chemistry near the end point of the alkalinity titration will contribute the alkalinity. Species such as acetic acid (as just seen) and formic acid contribute to alkalinity. Inorganic species such as hydrofluoric acid (pKa = 4.3) also contribute. Some boric acids also contribute.

The OLI code does not make any direct calculation of alkalinity since we do not know apriori what ions may appear in solution. We calculate the alkalinity via a titration. The OLI/ScaleChem program also uses this philosophy. When ions exist in the brine (e.g., acetate, formate, fluoride, borate, etc.) it is left to the user to determine what end point pH is appropriate for the alkalinity titrations.

This was former Tip35.