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This is the second in a series of articles about rethinking common practices in high school chemistry.

Classification of Chemical Reactions

The most prevalent reaction classification scheme in high school chemistry is the “Five Reaction Types,” which are synthesis, decomposition, single replacement, double replacement, and combustion. Students are generally taught to classify a reaction and predict the products of a given set of reactants. This article describes the shortcomings with the “Five Reaction Types” classification and provides an alternative organization to the study of chemical reactions based on three fundamental physical events: electrostatic attraction between ions, electron transfer, and proton transfer.

Classification of reactions and predicting products of reactions are necessary components of any first-year chemistry course. The shortcoming of the categories listed above is that they obscure the actual mechanism underlying the changes that are occurring. The “replacement” terminology is particularly unfortunate, as no “replacing” occurs per se. The terms “single replacement” and “double replacement” describe a visual pattern in the text of the molecular equation of the reaction, but not a pattern in the mechanism of the reaction itself.1

Consider the production of PbI2 (s) from aqueous solutions of Pb(NO3)2 and KI:

Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

It is common for teachers to say something akin to “lead and potassium switch places” or “lead replaces potassium in the compound.” Neither statement is accurate, however, because none of the substances’ aqueous solutions contain discrete molecules, but rather dissociated ions. Many students struggle with understanding that KNO3(aq) actually represents K+(aq) + NO3-(aq), and the writing of the molecular formula does not aid in their understanding. Furthermore, the vocabulary of “replacement” or “ion exchange” does not allay this misunderstanding. Since all strong acids, strong bases, and soluble ionic salts dissociate into ions, then reactions that represent those ions are more accurate representations of the species in the solution. When researching this topic, I found only one scholarly article that addressed it directly, and it dealt only with double-replacement vocabulary.2 The reaction is more accurately represented with an ionic equation:

Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) → PbI2 (s) + 2K+(aq) + 2NO3-(aq)

Customarily, we do not include non-reacting species (i.e. spectator ions) which, when removed, clearly show that the reaction is a precipitation reaction:

Pb2+(aq) + 2I-(aq) → PbI2 (s)

As teachers, we use simple models in place of more complex ones for beginning students all the time. Additionally, there is nothing wrong with teaching students to recognize patterns, as it is a useful way for anyone to organize information. However, we should be critical of the patterns and models we choose and not unintentionally overlook or omit important chemistry content. The “Five Reaction Types” are an inadequate scheme that can be reorganized to better emphasize the different types of change that occur.

I struggled with the use of the “Five Reaction Types” classification scheme for several years before reading two excellent books by Peter Atkins called Chemistry: A Very Short Introduction and Reactions: The Private Life of Atoms.3 Last year, using these two books, my own understanding of reactions, and how my students struggled with the concept, I reorganized my unit on chemical reactions by dividing it into three sections: precipitation, oxidation-reduction, and acid-base.4 Each section includes a laboratory experience and the opportunity for students to practice identifying reactions and predicting products. Additionally, there is an emphasis on the use of the net-ionic equation. In my experience, using precipitation, oxidation-reduction, and acid-base as the classifying categories is superior to using the “Five Reaction Types” scheme. Unlike that more traditional approach, students must identify reactions based on the mechanism that underlies each type of reaction: ion attraction leading to precipitate formation, electron transfer, and proton transfer.

Precipitation

In Reactions: The Private Life of Atoms, Atkins describes precipitation as “very simple and, I have to admit, not very interesting,” but he nonetheless goes on to describe the concept as a very useful introduction to understanding and visualizing chemical reactions. It was his treatment of precipitation in this book that inspired me to begin my unit with precipitation.

I begin with a demonstration of the precipitation of PbI2 (s) reaction mentioned above. The reaction is always a crowd-pleaser, because of the bright yellow precipitate produced. Performing the reaction as a demonstration, I can minimize the amount of the lead-containing product to be disposed of. Before the solutions are mixed, I use a small conductivity tester to demonstrate that both of them conduct electricity. In groups, students draw particle diagrams of each of the solutions on large white boards to illustrate that each solution contains dissociated ions (as evidenced by their conductive abilities). After I mix the solutions, I challenge the student groups to draw a particle diagram of the mixture, identify the yellow solid, and write a chemical reaction that shows the production of the solid. Students were introduced to basic solubility rules for ionic compounds in the previous unit, though I do not remind them of this at first.

After allowing them time to work on their own, we have a class discussion about the reaction and conclude that the yellow precipitate is PbI2 (s). I proffer the explanation that the solid forms because the ions are more attracted to one another than to the water molecules. Some discussion questions I may ask include:

  1. Are all ionic compounds soluble in water? How have we previously determined which are soluble and which are not?
  2. What possible ionic compounds could be produced from the species present in the solution?
  3. How do we know the formula of the solid is PbI2 instead of PbI or some other ratio of ions?
  4. What ions or molecules are present in the solution after the formation of the solid?

After the discussion, I teach my students the general rules of writing net-ionic, complete ionic, and molecular reaction equations, and we discuss the benefits and shortcomings of each. Given our previous emphasis on the difference in the dissociation of ions and dissolving of molecular compounds, students have a foundation on which to critique the different types of reaction equations.

My students then do a normal practice worksheet where they practice the skills in writing equations and drawing particle diagrams that they learned. They also complete a “dropper lab” that allows them to perform many reactions on a grid, write reactions, and draw particle diagrams.

Oxidation-Reduction

Chemistry is a dance of electrons.” – Richard Dawkins, The Greatest Show on Earth

In my experience, most high school chemistry courses do not introduce oxidation-reduction until later in the year, if at all. Many high school texts have a chapter late in the book, long after chemical reactions have been discussed. I was among those who avoided teaching to topic of redox, assuming it was too complex to discuss in a first-year course.

Then I got to thinking: the vast majority of chemical reactions are oxidation-reduction reactions. In fact, synthesis reactions, decomposition reactions, and combustion reactions all fall under the category of redox. We all already teach our students what happens to the charge of an atom when it gains/loses electrons, so they are primed to understand redox. Additionally, if the introduction to redox is limited to solid metals and neutral solutions, then the concept becomes much easier to illustrate, and the complex and laborious process of balancing equations in acidic and/or basic solutions can be dealt with at a later time. For me, that time is in second-year chemistry.

There are several possible labs that can be used to introduce redox. Some favorites are putting an iron nail into a copper(II) chloride solution, aluminum foil in a copper(II) chloride solution, or a copper wire immersed in a silver nitrate solution. I had my students perform the first reaction over a couple of class periods and collect mass data to uncover the mole ratio between the reacting copper and iron.

The reduction of copper on the surface of the iron nail is quick and recognizable. We let the solution sit overnight to collect sufficient copper and the next day in class, the nail is almost completely gone, a sufficient quantity of copper has been produced, and the solution is noticeably less blue in color than it was originally. I give my students the same challenge as with precipitation: represent the reaction with a particle diagram and a net-ionic equation. To give them a starting point, I suggest they consider the identity of “the red stuff” and where it came from.

At this point of the year, students have already seen the production of copper metal from solution, but some still want to call the substance “rust,” so that must be addressed. They are also familiar with the facts that neutral metal atoms do not dissolve in water but that ions do, and that copper(II) is responsible for the blue color in copper(II) chloride. If your students are unfamiliar with these points, they are easily illustrated by dropping a piece of metal into deionized water (nothing happens) and showing them a series of solutions of CuSO 4 (blue), Na2SO4 (colorless), CuCl2 (blue/green), and NaCl (colorless) — and then asking them to determine which ion is responsible for the color.

Following a discussion, we arrive at the explanation that the iron atoms must have lost electrons to be converted into dissolvable ions, and copper(II) ions must have gained those electrons to become neutral metal atoms. We decide on the following equation for the reaction:

Fe (s) + Cu2+(aq) → Fe2+(aq) + Cu (s)

I have been surprised when some students try to demonstrate electron transfer in the reaction, though their representations generally used subtraction signs to show an electron loss. I teach them the accepted way of showing this using half-reactions:

Oxidation: Fe (s) → Fe2+(aq) + 2e-

Reduction: Cu2+(aq) + 2e- → Cu (s)

Earlier this year, I used half-reaction notation to introduce students to writing electrons in equations when we covered ion formation, so this was not their first exposure to the practice.

Following the lab, I introduce the vocabulary of oxidation and reduction and teach students to assign oxidation numbers. Next, we address predicting the products of reactions. (In prior years, one area of struggle for my students was when to use the activity series to predict products of reactions. I constantly had to remind them to “only use it for single replacement reactions,” but could not provide any real explanation because they did not have an understanding of redox.) This time around, I was able to introduce the activity series as a reference table that showed relative ease of oxidation. It was self-evident that the table should only be used to predict redox reactions: if an element is more easily oxidized than another, then that oxidation will occur.5 The activity series I use has a big arrow at the side pointing upward with the words, “Ease of Oxidation Increases.” My students now are able to articulate which reactions will and will not occur in terms of ease of oxidation, rather than simply stating, “because it is more reactive,” or “because it is higher on the chart.” This improved discourse has been a welcome change to my class.

We then discuss a few different patterns of redox reactions: oxidation of metal by a salt, oxidation of a metal by an acid, oxidation of a metal by water, and combustion. I also mention that the first three of these reactions are referred to as replacement or displacement reactions because they follow the pattern: A + BX → AX + B. I feel that introducing the pattern at this point is appropriate, because students are using it with the understanding that the reaction occurs due to electron transfer. At this point, I also introduce the vocabulary of synthesis and decomposition reactions, many of which are redox reactions.

Acid-Base Reactions (Neutralization)

Lastly, we discuss acid-base reactions. Just as oxidation-reduction was the transfer of one subatomic particle, the electron, neutralization is the transfer of another, the proton (or H+ ion, as we chemists prefer to refer to it). This of course necessitates defining acids and bases, which I do in a historical way: first discussing the Arrhenius definition, and then the broader Bronsted-Lowry definition.6 I demonstrate the conductivity of weak and strong acid solutions, and have students propose a reason for the noticeably dimmer lightbulb in the weak acid solution. We decide that there must be fewer ions in the weak acid solution, meaning that not all the weak acid molecules ionized. I introduce the terminology behind strong and weak acids and bases, and then it is time to address reactions between them.

A simple neutralization demonstration follows, where I mix equimolar amounts of dilute NaOH(aq) and HCl(aq), both containing universal indicator solution. The blue and red solutions mix to form a green solution, showing that neutralization has occurred. It is important to practice this demonstration beforehand to make sure you obtain a green color, as universal indicator is very sensitive. Adding one solution to another using a buret would help in more precise volume additions. Alternatively, you could measure pH with a pH probe, but the same caution applies.

I then walk the students through the writing of the net ionic equation for the reaction. Since we do not have extensive prior work in acids and bases, pH, or using indicators, there are too many variables for me to feel confident in the students coming to the proper equation on their own. They are able to write the reactants on their own, however, and I tell them that a neutralization reaction produces water. Then we write the complete-ionic equation, net-ionic equation, and accompanying particle diagrams.

H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → H2O (l) + Na+(aq) + Cl-(aq)

H+(aq) + OH-(aq) → H2O (l)

I then give the students examples of writing equations for reactions that include weak acids and bases, using our knowledge that a solution of a weak acid contains mostly molecular acid species.

To conclude the unit, I have my students prepare a one-page summary of what we have learned for homework. Then, in groups, they prepare a short review presentation for the class. I also do a stations lab highlighting different types of reactions that I used with my old curriculum, modified with the changes I made.

Conclusion

After trying this organization of my chemical reactions unit last year, my students better understood the differences between different types of reactions, and they were also much more adept at writing and understanding equations, especially net-ionic equations. It is also worth noting that, except for “double replacement,” all vocabulary from the original “Five Reaction Types” is included, so there is not the concern they will be missing out on commonly used terms.

Following this unit, I move into stoichiometry. Alternatively, one could easily move into investigate simple electrochemistry examples, deeper treatment of acid-base chemistry, or gravimetric analysis of precipitation reactions. The best choice for your classroom will depend on where you are in your sequence and what you choose to cover in your first- or second-year course.

I look forward to further conversation on how we organize classifying chemical reactions in general chemistry courses so that we can continue to improve student understanding and discourse.

References

  1. It is worth noting that almost all reactions addressed using this classification scheme occur in aqueous solution. For example, the solid phase reaction between KI(s) and Pb(NO3)2(s) could not be appropriately classified as a precipitation reaction.
  2. Martin, B.R. Replace Double Replacement. Journal of Chemical Education 1999, 76 (1), 133.
  3. I recommend Lowell Thomson’s review of the first Atkins book, found at: https://www.chemedx.org/pick/book-review-chemistry-very-short-introduction-peter-atkins.
  4. Atkins addresses two types of reactions that I chose to omit from my first-year course: Radical Reactions and Lewis acid-base reactions.
  5. My preferred activity series is Table 4.5 in Chapter 4 of Chemistry the Central Science: 14th Edition by Brown, Lemay, Bursten, et al (2018).
  6. Although I currently omit Lewis acids and bases, I am reconsidering that approach for future classes.


Photo credit:
(top) PRHaney, via Wikimedia Commons