Alkalinity is an important parameter used to describe composition, buffer capacity and pH-behaviour of waters in many different fields of science. Although the alkalinity is widely applied, it seems to be a lot of confusion in connection with the concept, and it is often used incorrectly. One of the reasons for all the confusion is probably due to the first use of this concept to describe fresh water systems containing bicarbonate and carbonate only. In this case, when pH is between 4.5–9, the alkalinity will have the same numerical value as the bicarbonate concentration. Due to this fact the alkalinity concept and the bicarbonate concentration have incorrectly been substituted for one another.

Several experimental methods are available for alkalinity determinations where HCl titrations are most commonly used. With only bicarbonate present, the titration end point is close to pH=4.0–4.5, and in most of the published methods, the endpoint is given to be around this pH. For systems containing other weak acids like acetic acid, this titration method fails. The reason is simply that the presence of these acids changes the endpoint pH. Another method is addition of HCl to pH=2-4, boiling or flushing with nitrogen to evaporate CO2, and then back titration with NaOH.

Our aim is to show that a combination of the two methods, using carefully determined titration endpoints, will give very good measurements of both the alkalinity and the content of organic acids in the water. To clarify the alkalinity concept, we want to present a proper definition. We want to discuss how the alkalinity of a solution is varying with additions of acids and bases, and finally we want to present a simple method to measure this property. The work will focus on alkalinity in typical oil field waters.

Theoretical Considerations

In aqueous systems roughly 4 groups of molecules and ions are related to the alkalinity

  1. Neutral species: H2O, CO2(aq) or H2CO3(aq) CH3COOH(aq) etc.

  2. Dissociation products of water and weak acids: H+ and OH-, HCO3, CH3COO etc.

  3. Dissociation products of strong acids: Cl-, SO42- etc.

  4. Metallic cations: Na+, K+, Ca2+, Mg2+ etc.

For charged species, the equation of electro neutrality has to be fulfilled. As and example we may consider a water solution of NaCl, CaSO4, an acid, HA, and dissolved CO2. By rearranging the electro neutrality equation such that all pH dependent species (group 2) are on the right side, and all other species (group 3 and 4) are on the left, a definition of the total alkalinity is obtained

The concentration of the ions on the left side of Eq. 1, Na+, Cl- etc., are pH independent, and their concentrations are constant if pH changes. At this point, it is important to note that pH in oil field waters vary between 3-8. Species with a pKa value outside this range are completely dissociated or undissociated. That is why [SO24-] is considered pH independent in Eq. 1. The alkalinity can also be treated as the sum of titratable bases. [H+] is then a proton source and enters Eq. 1 with a negative sign.

How is alkalinity changed? One would think that an addition of an acid, HA, to a solution will change the alkalinity.

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