Universal pH Indicator as a Colorimetric Reagent for Differentiating Inorganic Anions

A simple colorimetric approach using a universal pH indicator to differentiate inorganic anions according to their relative acidity or basicity is presented. Anions caused changes in the pH of the solution, producing various colors of the universal indicator. Among halides, Fwas differentiated from Cl-, Brand I-. The indicator was also used on conjugate acid-base pair anions to distinguish HCO3 from CO3 2-, HSO4 from SO4 2-, HSO3 from SO3 2-, and various phosphate species. Oxyanions SO3 2-, HSO3 -, ClO-, ClO2 and NO2 can be differentiated from oxyanions with more oxygens attached, namely, SO4 2-, HSO4 -, ClO3 -, and NO3 -, respectively. Results can be correlated with the acid ionization constant Ka and/or base hydrolysis constant Kb of the anion.


INTRODUCTION
Acid-base indicators or pH indicators are a class of dyes that are common in most chemistry laboratories. These are weak acids or bases of which its undissociated form exhibits a color different from its ionic form. 1 It has been developed, primarily, for determining the pH of solutions. However, its application has expanded to include sensing various analytes such as gases, 2-7 organic compounds , 8,9 and cations. 10,11 Recently, our group has shown that anions can be differentiated based on its acidic or basic properties. Using flower pigments 12 or common laboratory pH indicators, 13 different anions change the color of the indicator depending on the pH produced in solution. This approach has been applied to differentiate conjugate acids and bases such as carbonates, sulfates and phosphates. While flower pigments can be extracted easily from natural sources, its color was found to be unstable for long periods of time and tend to vary depending on the extraction procedure and its source. Common laboratory pH indicators, on the other hand, are more stable and are readily available. However, the pH range of individual indicators is limited; thus, requiring several indicator solutions to perform the analysis.
In this study, a universal pH indicator serves as a "single" colorimetric reagent for the qualitative profiling of anions. This approach is simple, easy to conduct, and requires small amounts of reagents for analysis. The method can have practical applications in chemical education. It can serve as a teaching demonstration or microscale laboratory experiment in senior high school or undergraduate general chemistry, or analytical chemistry class to illustrate solution properties of anions, understand the nature of amphiprotic anions, and relate the effect of equilibrium constants K a and K b to varying anionic species.

Reagent and Microplate Colorimetry Preparations
1.5 mL of 0.1M salt solutions in distilled or deionized water were prepared. 0.3 mL of each solution were placed into 96-well microplates, with one column composed of three wells serving as the three trials for each salt. A separate column was used for the control group (water only). 15.6 mL of the universal pH indicator was added into each test solution.

RESULTS
One of the most versatile acid-base indicators developed is the universal pH indicator. 14 Also known as Yamada Universal Indicator, this solution consists of thymol blue, bromothymol blue, phenolphthalein and methyl red. 15 It operates at long pH ranges (pH 4-10). In this study, a universal pH indicator was used as a "single" reagent for differentiating anions. The addition of the universal pH indicator to individual solutions of anions led to color changes as shown in Figure 1. Among the halides, F -, Cl -, Br -, and I -, only Fwas generally distinct from the others. Fled to an olive-green solution while other halides have colors similar to the control solution at 0.1M concentration. At a higher concentration (1M), Fis still significantly different from Cl -, Br -, and I -, but the color has turned to blue. Moreover, Icould now be distinguished from the other halides with a color change from yellow to olive-green. This observation may be attributed to an increase in ionic strength of halide solutions which caused a shift in the color of the indicator. This behavior has also been noted by Rodriguez and Mirenda 16 in which an increase in salt concentration led to a shift of the pH indicator towards the basic form. Thus, it is important to compare anions with the same concentration and under dilute conditions. C o n j u g a t e a c i d s / b a s e s we r e a l s o distinguished from each other. It was observed that the color of the SO 4 2solution is yellow while that of HSO 4 is red. Meanwhile, SO 3 2was notably violet while HSO 3 was red-orange. Different phosphate species exhibited different colors: PO 4 3is violet, HPO 4 2is blue, and H 2 PO 4 is orange. It was also observed that CO 3 2is violet while HCO 3 is blue. As expected, the conjugate acid form of the anion tends to shift the color of the indicator towards the acidic color relative to its basic color.
Finally, oxyanion pairs sulfate-sulfite, bisulfate-bisulfite, nitrate-nitrite, and oxychlorides were also differentiated from each other. Lessoxygenated oxyanions SO 3 2and ClO 2 resulted in blue to violet colors. ClO -, at the onset, is violet but eventually turns to light green, then colorless. It should be noted that hypochlorite, ClO -, is a widely-used bleach and a strong oxidant, and is known to degrade dyes. 17,18 Meanwhile, oxyanions with a higher number of oxygen atoms attached, SO 4 2and ClO 3 -, result in a yellow solution. On the other hand, for oxyanions NO 3 and NO 2 -, nitrate exhibited a dark green color, while nitrite ion yielded a yellow-green solution.
The analysis was also performed in degassed distilled and degassed deionized water. Similar results were obtained as those for simple distilled water. This implies that minimal impurities like dissolved gases in the air do not significantly affect the results of the analysis.

DISCUSSION
The colors observed for anion solutions correlate well to the colors of the indicator at a given pH of the solution. This was verified using a pH meter. Conjugate bases and lessoxygenated oxychlorides ClO 2 -, ClO -, CO 3 2and PO 4 3with pH of 10 or greater exhibited a violet color. SO 3 2and HPO 4 2with pH 9 to less than 10 imparted a blue solution. HCO 3 with pH ranging from 8 to less than 9 yielded a blue-green color. On the other hand, NO 3 and Fwith pH 7 to less than 8 presented a green color. Oxyanions SO 4 2-, ClO 3 and NO 2 with pH from 6.4 to less than 7 displayed a yellow-green solution. Meanwhile, halides I -, Cl -, and Brwith pH from 6 to less than 6.4 imparted a yellow color. For protonated anions, H 2 PO 4 -(pH 4-5) gave off orange color, while HSO 3 and HSO 4 -(pH < 4) yielded redorange or red, respectively.
Aside from pH, the color of the solution imparted by anions can be correlated to its acid ionization constant (K a ) or base hydrolysis constant (K b ). 19 This allows grouping of anions based on its acidic or basic properties as shown in Fig. 2. Anions which turn the pH indicator to red or red-orange (HSO 4 -, HSO 3 and H 2 PO 4 -) are protonated species and were found to have K a values greater than its K b (Table 1). Thus, these anions tend to produce more H 3 O + in solution via acid ionization than OHions produced via base hydrolysis, as exemplified by HSO 4 in Figure 3.  *Calculated from K b = K w /K a where K w is the ionization constant of water, and K a is ionization constant of the conjugate acid of the anion. This approach is simple and can be used as a demonstration or laboratory experiment in high school or freshmen college chemistry to help learners understand the acidic or basic nature of anions. It also applies principles of chemical equilibrium, ionization constants, and nature of amphiprotic anions through visual observation.