Research: Science and Education
Nontypical Brønsted Acids and Bases ⁄ Tadeusz Micha lowski,* Maciej Rymanowski, and Andrzej Pietrzyk Faculty of Chemical Engineering and Technology, Cracow University of Technology, ul. Warszawska 24, 31-155 Kraków, Poland; *
[email protected] According to Brønsted’s theory, the term “acid” is attributed to forms (ions, molecules) that are capable of splitting off protons, whereas the term “base” refers to forms able to accept protons; the “origin” of H+ or OH− ions is of minor importance here. This means that forms defined as acids or bases can provide the corresponding ions in ready form or can act as factors that split water molecules (in hydrolysis processes) or in other protonophoric components of the system in question. In this article some nontypical Brønsted acids and bases will be discussed. Oxides and Oxide Species Oxide forms, O2−, are seldomly mentioned when considering bases. However, when dissolved in acids, for example, CuO + 2H+ → Cu2+ + H2O the oxide readily accepts protons and is justifiably classified as a base. Other “oxide” forms are rarely (if at all) mentioned as bases, despite the fact that they act as polyfunctional bases (when compared for example with NaOH). As an example, we can consider the MnO4− ion, which “swallows up” protons like an octopus (pardon—an eight-hydroxyl base molecule), as shown in eq 1: MnO4− + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O (1) As a result of progressive addition of KMnO4 solution into an acidified solution of iron(II), the pH of the solution increases monotonically (1). Insufficient concentration of acid (H2SO4) in the solution can cause a change in reaction stoichiometry, with MnO2 precipitating as the product of reaction: MnO4− + 3Fe2+ + 4H+ → MnO2 + 3Fe3+ + 2H2O (2) In other words, the presence of a sufficient excess of sulfuric acid prevents an appreciable increase in pH resulting from the course of reaction 1. Cations and Anions Hydrocyanic acid (HCN, pK1 = 9.2) dissociates only slightly in water. Addition of AgNO3 into this solution causes a substantial increase in hydrogen ion concentration, resulting from the binding of CN− ions by Ag+, mainly in the form of the soluble complex Ag(CN)2− in the first stage of AgNO3 addition. In the second stage, the precipitate of Ag[Ag(CN)2] (i.e., AgCN) is formed:
Figure 1. The pH versus volume relationship for simulated titration of 100 mL of 0.01 mol兾L HCN with V mL of 0.1 mol兾L AgNO3.
ated HCN particles as a result of formation of strong complexes Ag(CN)i1−i (i = 2, 3, 4), (Figure 1). Similarly, Fe3+ ions release protons from water particles in a hydrolysis process, for example, Fe3+ + H2O → FeOH2+ + H+
(5)
Acidic behavior is also attributed to FeOH2+ ions, FeOH2+ + H2O → Fe(OH)2+ + H+
(6)
In effect, Fe(NO3)3 is an acid of moderate strength, similar to acetic acid, although such properties are, in some degree, diminished by the formation of iron(III) hydroxide precipitate from some of the ferric ions. Nevertheless, the pH of the solution does not increase substantially; it is mainly affected by the fact that additional protons are liberated in the reactions shown below: Fe3+ + 3H2O → Fe(OH)3 + 3H+
(7)
FeOH2+ + 2H2O → Fe(OH)3 + 2H+
(8)
Rather unexpectedly, at first sight, Na2S2O3 solution in reaction with an acidified (H2SO4) solution of KIO3 (or KIO3 and KI) acts as a strong base (like NaOH) (1), (Figure 2). This reaction, known also from qualitative analysis, can be written as
2HCN + Ag+ → Ag(CN)2− + 2H+
(3)
Ag(CN)2− + Ag+ → 2AgCN
(4)
IO3− + 6S2O32− + 3H2O → I− + 3S4O62− + 6OH− (9)
From eqs 3 and 4 it can be seen that Ag+ ions act as a strong acid in the system, liberating protons from weakly dissoci-
which can be derived from the related speciation plots (Figure 2C).
470
Journal of Chemical Education
•
Vol. 82 No. 3 March 2005
•
www.JCE.DivCHED.org
Research: Science and Education
Quantitative Description
A
Let As2O3 be dissolved in NaOH and neutralized, successively, with H2SO4 and disodium malonate (Na2Mal). V0 = 50 mL of the resulting solution, containing HAsO2 (C0 = 0.002 mol兾L) + NaOH (0.01 mol兾L) + H2SO4 (0.01 mol兾L) + Na2Mal (CMal mol兾L), is adjusted, with 0.1 mol兾L NaOH, to the pre-assumed pH value (indicated at the corresponding curves in Figure 3) and then titrated with V mL of C = 0.02 mol兾L I2 in 0.2 mol兾L KI. Two values for CMal were chosen: 0.01 in Figures 3A and 0.03 in Figure 3B. The resulting curves indicate, more or less pronounced breakpoint closely related to the fraction titrated Φ = (C兾C0)(V兾V0) = (0.02兾0.002)(5兾50) = 1. At 0.03 mol兾L Na2Mal and other data identical with ones specified above, the system with pre-
B A
B
C
Figure 2. The (A) pH versus volume, (B) E versus volume, (C) log[Xi] versus volume for different iodine species Xi relationships (simulated curves) related to titration of V0 = 100 mL of the mixture composed of KIO3 (C0 = 0.01 mol兾L) + KI (CI mol兾L) + H2SO4 (0.01 mol兾L) with V mL of Na2S2O3 (C = 0.1 mol兾L). In Figures 2A and 2B, the curves labeled a refer to CI = 0.1 mol兾L and the curves labeled b refer to CI = 0; the curves in Figure 2C refer to CI = 0 (2).
www.JCE.DivCHED.org
•
Figure 3. The pH versus volume relationships related to titration with V mL of C = 0.02 mol兾L I2 in 0.2 mol兾L KI (as titrant) of V0 = 50 mL of the solution containing HAsO2 (C0 = 0.002 mol兾L) + NaOH (0.01 mol兾L) + H2SO4 (0.01 mol兾L) + Na2Mal (CMal , mol兾L) and adjusted, to the pre-assumed pH value with 0.1 mol兾L NaOH. Numbers on the curves in (A) indicate initial pH values of the solution titrated; CMal equals 0.01 mol兾L in (A) and 0.03 mol兾L in (B). pH = 6.0 was assumed as the starting pH in (B).
Vol. 82 No. 3 March 2005
•
Journal of Chemical Education
471
Research: Science and Education
Figure 4. The pH versus Φ dependencies (simulated curves) for titration of V0 = 100 mL of the mixture containing KIO3 (C0 = 0.01 mol兾L), H2SeO3 (0.02 mol兾L), HCl (0.02 mol兾L), and HgCl2 (CHg mol兾L) with V mL of 0.1 M ascorbic acid; curve a refers to CHg = 0 mol兾L; curve b refers to CHg = 0.07 mol兾L (1, 2).
Figure 5. The simulated titration curves related to the system specified in caption of Figure 4; (a) CHg = 0; (b) CHg = 0.07 mol兾L (1, 2); Φ = CV兾(C0V0).
adjusted pH = 6.0 value gives the titration curve consisting of nearly rectilinear segments intersecting at Φ = 1 (Figure 3B). In the titration of a mixture containing KIO3 + H2SeO3 + HCl (0.02 mole/L) + HgCl2 (CHg mol兾L) with V mL of ascorbic acid (C = 0.1 mol兾L), some non-monotonic relationships between pH and Φ were observed (1, 2) (Figure 4). This means that the titrant acts alternately as a base within some volume interval and as an acid within another one. Similar phenomena were observed in (i) VSO4 and KMnO4, (ii) VSO4 and K2Cr2O7 systems (both acidified with H2SO4), and in (iii) KI and Cl2 system (1).
were discussed in ref 1 in connection with Figure 5. The calculation procedure suggested enables one to select consciously and control the conditions of analysis. Moreover, it enables one to formulate a generalized equivalence mass concept (2, 5–7) without referring to reaction equation notation. Such an approach contradicts the approach suggested by IUPAC (9–12).
Discussion It is worth noting that the properties of electrolytic systems discussed above were indicated on the basis of simulated titrations made with an iterative computer program MINUIT. The algorithms applied enable one to exploit the attainable physicochemical data on the systems in question. For example, 36 independent equilibrium constants were involved in five concentration balances, one charge balance, and one electron balance related to the system with ascorbic acid and iodate (quoted above) (1). The procedure proposed in refs 1 and 2 does not require any simplifying assumptions and then the complexity of the system has no significance. It enables the analyst to follow, in the course of a simulated titration, the speciation of the system and then formulate equations for the corresponding reactions together with their efficiency (1–7). All analytical procedures can be reconstructed. It is even possible to correct some experimental results or mistakes made; for example the errors made in ref 8
472
Journal of Chemical Education
•
Literature Cited ⁄ 1. Michalowski , T.; Lesiak, A. Chem. Anal. (Warsaw) 1994, 39, 623. ⁄ , T.; Wajda, N.; Janecki, D. Chem. Anal. (War2. Michalowski saw) 1996, 41, 667. ⁄ 3. Michalowski , T. J. Chem. Educ. 1994, 71, 560. ⁄ , T.; Lesiak, A. J. Chem. Educ. 1994, 71, 632. 4. Michalowski ⁄ 5. Michalowski , T. Chem. Anal. (Warsaw) 1981, 26, 799. ⁄ 6. Michalowski , T. Orbital 1999, (3), 174. Michalowski, T. Orbital 1999, (5), 399. ⁄ , T. Calculations in Analytical Chemistry with Ele7. Michalowski ments of Computer Programming (in Polish); PK: Kraków, 2001. 8. Erdey, L.; Bodor, E.; Buzas, H. Z. Anal. Chem. 1951/52, 134, 22. 9. West, T. S. Pure Appl. Chem. 1978, 51, 325. 10. IUPAC Manual of Symbols and Terminology for Physicochemical Quantities and Units. Pure Appl. Chem. 1979, 51, 1. 11. Freiser, H.; Nancollas, G. H. Compendium of Analytical Nomenclature, Definitive Rules; Blackwell Scientific Publications: Oxford, 1988. 12. Mills, I.; Cvitas, T.; Homann, K.; Kallay, N.; Kuchitsu, K. IUPAC Quantities, Units and Symbols in Physical Chemistry; Blackwell, Scientific Publications: Oxford, 1988 and 1993.
Vol. 82 No. 3 March 2005
•
www.JCE.DivCHED.org
