In the Classroom
Teaching Brønsted–Lowry Acid–Base Theory in a Direct Comprehensive Way Jamie L. Adcock Department of Chemistry, University of Tennessee, Knoxville, TN 37996-1600;
[email protected] As a teacher of general chemistry for more than 20 years, I have sought to teach each subject efficiently and effectively. One subject that provokes many questions and takes its toll on student grades is the myriad relationships existing between strong and weak acids, bases, and their conjugates. Adding to student’s confusion is the inexactness of the terms strong, weak, and very weak in regard to these species and how these qualifiers affect the interaction of acids and bases with each other and with water, itself both acid and base. Adding to this the concept of the leveling effect that occurs when strong acids/ bases are placed in water and the non-interaction when very weak acids/bases are involved can create chaos in the mind of a student trying to make predictions and draw insight and solve quantitative problems. The problem is that each relationship, although simple, is linked to every other and we teach them individually in separate sections and expect students to be diligent and thoughtful enough to make the connections that practicing chemists/scientists take for granted. Spurred by the persistence of a particularly challenging class during the summer of 2000, which demanded succinct answers, I devised a figure (Fig. 1) that depicts or implies
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HI I− HBr Br− HCl Cl−
all the relationships simultaneously. The students found it helpful, and compared with students in all previous classes they improved their grades on the examination that was heavily weighted with acid/base equilibria and buffer questions. The Figure The assumptions implicit in Figure 1 are (i) the use of logarithmic pK values; (ii) the relationship that pKa + pKb = pKw; and (iii) ignoring the leveling effect of water in order to suggest it! Logarithmic pKa values for a series of inorganic and organic monobasic acids including the hydrohalic acids and the weak base ammonia are included. In each example the pKa value for the conjugate acid and the pKb value for its complementary conjugate base are shown aligned with the numeric scale above and with the corresponding symbol/formula at the left. Most acids and their conjugates are considered dilute; the one exception is water, which as both conjugate acid and base is by nature 55.33 mol/L at 25 °C. This skews the values for the hydroxide and oxonium ions to 1.74 pH units greater than 14 and less than zero. The pKa and pKb values tabulated are derived from Ka and Kb values commonly quoted in general chemistry texts (1, 2) and a general inorganic text (3). The figure is descriptive, not quantitative, although an effort was made to not misrepresent values for any chemical species. For each species, the sum of the pKa and pKb values is 14.00. In fact, this relationship was used to calculate approximate values for the conjugate base strengths for chloride, bromide, iodide, and nitrate ions. Application of the Figure
HF F− H3O+ H2O HNO3 NO3− HClO2 ClO2− H2CO2 HCO2− HOAc Ac− HCN CN− NH4+ NH3 H2O OH− Strong
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Figure 1. The numerical values refer to pKa values for the conjugate acids (CA, 䊊) and pKb values for the conjugate bases (CB, 䊉). The modifiers strong, weak, and very weak apply to either acid or base. The divisions implied by the shading are somewhat arbitrary, but encompass the 0–14 pH scale where pKa /pKb values apply most importantly to buffers.
The concepts of strong, weak and very weak can now be quantified for both acids and bases, although few strong nonionic, non-hydroxide bases are commonly known. Quantification of these qualifiers allows important and definitive inferences to be made about the reactions possible when acids and bases in each category are dissolved in water. For example, values for the hydrohalic acids H–X (X = I, Br, Cl) are estimated from thermodynamic considerations, since they immediately react with water to form the hydrated oxonium ion. All acids stronger than the oxonium ion are thus leveled to the acid H3O+(aq) when actually dissolved in (reacted with) water. The corresponding conjugate bases of these acids, the anions I ᎑, Br᎑, and Cl ᎑ are very weak and will not hydrolyze. A strong acid is one whose pKa is smaller than ᎑1.74; only a negligible amount of the original acid will remain after mixing, and the reaction with water is effectively 100%. This reaction is the “leveling effect” of water on acids stronger than the aqueous hydrogen ion. A very weak acid is one whose pKa is greater than 15.74 and it will not react with water to effectively alter pH. This point allows one to logically explain why the very weak conjugate bases of strong acids (iodide,
JChemEd.chem.wisc.edu • Vol. 78 No. 11 November 2001 • Journal of Chemical Education
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In the Classroom
bromide, chloride, etc.) do not hydrolyze; their pKb values are larger than 15.74. A weak acid is one whose pKa is between 15.74 and about 2.6 (if one assumes 5% dissociation); it will produce a non-negligible change in pH due to partial reaction (equilibration) with water. Conjugates of weak acids are also weak bases and as a result undergo hydrolysis. Conjugate base pKb values are displayed on the same scale with the conjugate acid’s pKa value and utilize the same qualifiers. A strong base is one whose pKb is smaller than ᎑1.74; only a negligible amount of the base will remain because the reaction with water is effectively 100%. This reaction is the leveling effect of water on bases stronger than the aqueous hydroxide ion. A very weak base is one whose pKb is larger than 15.74. Very weak bases will not react with water to effectively alter pH; that is, they do not hydrolyze if they are ions. The pKb of a weak base is between 15.74 and about 2.6 (if one assumes 5% dissociation) and a weak base will produce a non-negligible change in pH. This layout illustrates the complementary nature of the conjugates of acids and bases and makes the point that there is essentially no difference in nature between the acids HCN and NH4+ except in the way we, as chemists, have previously classified them. The same relationship is seen with the conjugate bases CN ᎑ and NH3. We thus explicitly describe the hydrolysis of salts of weak acids and bases as simply a consequence of having an ionic or charged acid or base rather than a neutral, molecular one.
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Conclusion This figure enables a lecturer to unify all the Brønsted– Lowry acid–base concepts in a quantitative way. It is frustrating when students who are introduced to acid–base behavior memorize numerous specific equations as a problem-solving approach. These students fail to see the exquisite unity presented in the general concept. A few of my better students will realize that two simple equilibria and two simple equations can be used to solve all acid, base, buffer, and hydrolysis problems. Using this diagram in my general chemistry classes as a handout has resulted in more students realizing this unity of concept, and that makes the effort worthwhile. Acknowledgment I would like to acknowledge the helpful suggestions of George K. Schweitzer and the persistence and determination of a gifted summer 2000 class in general chemistry. Literature Cited 1. Hill, J. W.; Petrucci, R. H. General Chemistry, 2nd ed.; Prentice Hall: Upper Saddle River, NJ, 1999. 2. Ebbing, D. D.; Gammon, S. D. General Chemistry, 6th ed.; Houghton Mifflin: Boston, NY, 1999. 3. Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth Heinemann: Boston, 1997.
Journal of Chemical Education • Vol. 78 No. 11 November 2001 • JChemEd.chem.wisc.edu
