[This is a reprint from Sharon’s blog “Art of Science Teaching” from June 2016.]
Do you use the hydronium ion when writing a weak acid equilibrium system? This is a question I want to ask the crowd at the next ChemEd conference to find out where high school chemistry teachers fall in this debate. As
I head into my twentieth — yes, that’s a really long time — year of teaching chemistry, I can honestly say that I am a complete hydronium ion convert.
I started off, back in the early years, using the H+ ion exclusively because it was easier and quicker to write.
HA(aq) --> H+(aq) + A-(aq)
HA(aq) <-- H+(aq) + A-(aq)
The acid ionization equation without the hydronium ion notation.
I can imagine my younger chemistry teacher self-justifying this choice by saying “Why make things more complicated? Isn’t a weak acid system challenging enough for my students?” And let’s face it, writing out the equilibrium system with the water and hydronium ion requires not only more effort every time, but also an explanation — the first time, and then at least once a day — about why the water in this case is actually part of the reaction and not an “innocent bystander” like in other aqueous solutions. Streamlining the acid ionization equation with the H+ ion makes the weak acid the star of the show, with the emphasis on the loss of the proton from the acid. Referring to H+ as a proton also requires a bit of explanation and a few moments of digesting by most high school chemistry students, but they usually get the hang of it after a few examples.
And now, many years later, I have joined the hydronium ion camp.
HA(aq) + H2O --> H3O+(aq) + A-(aq)
The acid ionization equation with the hydronium ion notation
The transition happened when I returned to AP Chemistry after a ten-year gap in teaching the course. In my new school, I used Chemistry & Chemical Reactivity by Kotz, Treichle and Townsend for AP Chemistry. This excellent book uses the hydronium ion exclusively throughout the book when describing acid/base equilibria. In an attempt to be consistent, I adopted the practice of writing the hydronium ion, along with the many side conversations about what it is and why it’s there. To my delight, I found that my students engaged in more discussion about the chemistry of the system because of this simple — and some teachers would say insignificant — change in notation. Rather than making the weak acid the star of the show, the hydronium ion helps to make the process of proton transfer the important feature in the reaction. When students can see water as a weak base in the reaction, they have crossed an important hurdle to understanding the equilibrium system.
I recently had the pleasure of discussing the hydronium ion with Dr. Binyomin Abrams of Boston University. Dr. Abrams makes a good case for including hydronium ion in chemical reactions. First of all, the proton doesn’t just fall off the acid, it is transferred to the water molecule in a chemical process. Dr. Abrams uses the analogy of dropping a pencil. If you are holding a pencil in your hand, usually you wouldn’t just drop your pencil on the ground for no reason. However, if a classmate asks to borrow your pencil, you would purposely hand it over to the person. We can imagine something similar happening on the molecular level between a weak acid and a water molecule. The proton from the acid is transferred from the acid to the water in the solution. Additionally, including the chemical process of proton transfer in the weak acid system lays an important foundation for more complex reaction mechanisms. Esterification is a good example of how the proton transfer process in a weak acid can help lay a foundation for a more complex reaction mechanism with a similar action of lone pairs on the oxygen atom of an alcohol. After my brief conversation about the hydronium ion with Dr. Abrams, I held my head a little bit higher as a hydronium ion enthusiast.
But there are also good reasons not to use the hydronium ion. In high school chemistry, there is a fine balance between telling the students enough information to understand the chemistry and giving them too much information about what is “really happening”. We have all had to retract a statement because we dove into an explanation that goes beyond the scope of a high school class. The glazed over look that comes from the “long” explanation of what’s really happening in the system is a red flag for me. But in the case of the hydronium ion, can we say with confidence that it actually exists and is the best explanation of the proton in a weak acid system? In a letter to Journal of Chemical Education (The Solvated Proton is NOT H3O+!, 2011, 88 (7), pp 875–875), Dr. Todd Silverstein from Willamette University argues that the hydronium ion is not an accurate representation of the species; it actually exists as a proton that is surrounded by many water molecules. And, why give a proton special treatment when we exclude the solvating water molecules around other ions such as Cu2+ and Co2+? His letter effectively wiped that self-righteous grin off my face. I can’t look in the mirror and say “at least I’m telling my students the truth about the weak acid system” because the real truth is much more complex than a one-to-one proton transfer. Just like the famous courtroom scene from the movie A Few Good Men, Dr. Silverstein is saying to me “You can’t handle the truth!”
I can hear the arguments from a pair of tiny chemistry professors on each of my shoulders. On one side I hear “The hydronium ion is always the best choice, your students deserve it.”; while the other side is saying “Why do you make things more complicated than necessary? Stick to what they need to know.” For now, I’m listening to the hydronium ion argument because I believe that it best achieves my ultimate goal which is to teach my students how to think like a chemist.
[We asked Carey Bissonnette, a first year University of Waterloo chemistry lecturer and below is his comment on Sharon’s dilemma.
In our first year courses, we describe the self-ionization of water and the ionization of an acid as a proton being
transferred to a water molecule to form an H3O+ ion. We also emphasize that H3O+ is further hydrated to form more complex species. If you do not include a water molecule as a reactant, and H3O+ as one of the products, then students can easily be misled into thinking that H2O molecules dissociate to form H+ and OH- and that an acid molecule (e.g., CH3COOH) dissociates to form H+ and an ion (e.g., CH3COO-).
Dissociation and ionization are not synonymous. Dissociation refers to the separation of entities that already exist whereas ionization refers to the generation of ions.]