Small molecules carrying an acid and/or a basic group tend to be soluble in pure water and exhibit a certain degree of ionization (Section 1.1.6) or dissociation (Section 1.1.5). The ionization or dissociation of the acid HA in aqueous solutions (H2O + HA ⇌ A− + H3O+) at a given temperature is characterized by the acid ionization constants and . They are defined as:
where is the activity of the species . The activity of water is approximately constant at a given temperature and with the low solute concentrations normally used in the ESTs. Therefore, it is assumed that . Equivalently, the conjugated acid BH+ of base B, produced by the equilibrium reaction H2O + BH+⇌ B + H3O+, has the following acid ionization constants:
These acid ionization constants are better represented by , where .
The ranges defined for very strong (), strong (), medium, weak, and very weak acids are poorly defined. Nevertheless, it is safe to say that acids with a of between 4 and 10 are weak acids and that acids with can be considered as very weak acids. The opposite occurs for bases, as for a base to be very strong the of the conjugated acid BH+ (equation 1.15) must be above 14 (), with strong bases exhibiting . Similar to weak acids, weak bases also exhibit of between 4 and 10; however, in this case the very weak bases exhibit .
It is important to note that significantly changes with temperature for some functional groups: , where is expected to be a smooth and slowly varying function of temperature. The of the great majority of amines decreases with temperature, while carboxylic acids exhibit a much smaller change, usually negative, but there are some exceptions and it depends on the temperature range. These temperature sensitivities have important practical implications for method development within the field of ESTs, as they affect the mobility of the analytes and the pH of the buffers.[9–11] Moreover, they are used to promote cyclic band compression in the toroidal layouts, which is an interesting way to get some control of band spreading along the separation mediums (see Appendix G).
1.1.10 Concentration–pH and pa–pH Diagrams
It is important to know the concentration of all chemical species present in a given buffer solution at a given pH, because some species may interfere with the migrating analytes under study. Additionally, it is also important to know the concentrations of all existing species of an analyte present in a separation medium at a given buffer pH. This is important because the analyte species present define the average electrophoretic mobility of the analyte, the system peaks, and the interactions of the species with the BGE components. –pH diagrams are one of the most commonly used tools to visualize the concentration of chemical species, showing the concentration of each species at every pH in the range (which is the maximum range used in ESTs). Figure 1.4 shows the –pH diagram of a triprotic acid (0.1 M citric acid in an aqueous solution), which is a complex acid because its successive ionizations produce many species. In Figure 1.4 only the most concentrated species can be seen, as the curves of the most diluted species (with M) run too close to the line at M to be observed. Nevertheless, many detectors (fluorescence and potentiometric) are able to detect species down to M and even lower concentrations. Therefore, sometimes it is desirable to visualize the concentration profiles of species within lower concentration ranges. p–pH diagrams are ideal for this as they show as a function of pH. Figure 1.5 shows the same case as studied in Figure 1.4 (–pH diagram of citric acid at the initial concentration of M), but it is now represented as a p–pH diagram.
Figure 1.4 c–pH diagram of citric acid species in an aqueous solution in the 0 to 14 pH range. 0.1 mol of citric acid was used to prepare 1 L of solution, i.e.
