Arrhenius confuses students

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Stephen J. Hawkes
Oregon State University
Corvallis, OR 97331 USA

Published in the Journal of Chemical Education
Vol 69, No. 7; July 1992; pages 542-543

Note to student readers: this is an article published for chemistry teachers. As such, it assumes a fair amount of chemical knowledge that you may not yet have. The ChemTeam's advice is to learn the Brønsted-Lowry theory of acids and bases really really well!

The traditional presentation of the Arrhenius acid-base concept before the Brønsted-Lowry confuses and misleads students. The Brønsted-Lowry approach should be presented first, because it is simpler; it involves only the transfer of a proton; and it does not lead to the misconceptions and complications discussed below. The Arrhenius definitions could serve later as an historical footnote but not as a viable concept.

If a base is defined by its ability to produce OH¯ as suggested by Arrhenius, it requires a convoluted argument to show that NH3 and HCl react in an acid-base reaction in the absence of water. It can be done, invoking the fact that NH4Cl is produced either in the absence or presence of water and is the "salt" produced from NH3(aq) and HCl(aq). The relevance of the OH¯ still remains as a stumbling block to understanding, even then. This problem leads to other misconceptions. I have seen statements in textbooks that acetic acid and ammonia, for example, react in aqueous solution because they produce H+ and OH¯ which then interact. The student is thus led away from the simple understanding that the acid and the base interact directly.

The meaning of acidity also is confused by the Arrhenius definitions. Water at 100 °C has pKw = 12.3. When asked whether it was, therefore, more acidic than at 25 °C when pKw = 14.0, some of my students argued that the increase in [H+] made no difference: it was offset by the increase in [OH¯]! This is clearly an artifact of the Arrhenius concept, because they gave the opposite response when asked about a decrease in pKa of acetic acid, the basicity of the acetate ion notwithstanding.

The need for an OH group led Arrhenius to propose NH4OH as the formula for hydrated ammonia, giving us the "ammonium hydroxide" label on reagent bottles. This misconception is apparently not eliminated even now, for it has appeared on the pages of this Journal (1) as late as 1991.

A student struggling to discard the Arrhenius concept cannot absorb the concept of conjugate acids and bases and has difficulty in believing that negative ions are bases unless they are hydroxide ions. The Arrhenius approach requires that they first appreciate that negative ions react with water producing OH¯ without providing any reason to expect such reaction, while the Brønsted-Lowry concept makes it obvious that any negative ion must be basic simply because of its negative charge.

It is inherent in human nature that we accept what we are told first and relinquish or change it with difficult. Like most students of my generation, I learned the Arrhenius concept in my first year of college. (My high school teacher held pre-Arrhenius concepts.) I learned the Brønsted-Lowry concept as a college senior. Consequently, I found the latter difficult, and it took decades to realize that it is the simpler of the two concepts. We should not confuse another generation of students in this way.

Protons Are Not "Donated"

The reaction

HCl ---> H+ + Cl¯

requires an input of 1.4 x 106 J/mol (2). When that much energy is required, it makes no more sense to speak of HCl "donating" a proton than of "donating" your purse to a mugger. It creates a false concept of the action of an acid, as if it somehow expelled its proton by means of some internal force. A base must tear the H+ from the powerful attractive forces holding it to an acid, breaking its bonding by superior force.

My chemistry teacher, William Gerrard, nearly half a century ago provided me the above argument and the simple definitions:

I recommend these definitions to teachers and publishers and especially to students.

Literature Cited

1. Tuttle. T. R. J. Chem. Educ. 1991, 68, 533.

2. Calculated from data In Chase, M. W.; Davies, C. A.; Downey, J. R.; Frurip, D. J.; McDonald, R. A.; Syverud, A. N. JANAF Theremochemical Tables, 3rd ed., American Chemical Society: Washington, DC, and American Institute of Physics: New York, 1986

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