Colors in Chemistry

I recently finished the module on Stoichiometry and Chemical Reactions in my G-Chem class. The last group of reactions we looked at were redox reactions. Of these, one of the classic demos is the metal displacement reaction. I usually put in zinc or magnesium, a grey and shiny metal, into a blue solution of copper sulfate. Moments later, the grey metal begins to dissolve, the solution turns colorless, and out pops a brownish solid – copper metal. Our G-Chem lab has the students run through what’s known as the “copper cycle” where they get to observe these changes and more – although in a Covid year, they will likely not be able to actually do this in person.

 

The metal displacement reaction I described above can be represented by the following chemical equation.

 

Zn(s) + CuSO₄(aq)  -->   Cu(s) + ZnSO₄(aq)

 

The students learn how to write the net ionic equation and identify sulfate as a spectator ion that does not participate in the reaction.

 

Zn(s) + Cu²⁺(aq)  -->  Cu(s) + Zn²⁺(aq)

 

We also break up this net ionic equation into its two half-equations to track the loss and gain of electrons in the redox reaction.

 

All this is accompanied by nice “molecular” level pictures of what’s going on in solution. Here’s one from the current textbook we’re using. 

 

Notice it juxtaposes the microscopic world pictures with the macroscopic picture you’d observe in a real-life demo. Interestingly, the particle picture of zinc metal colors the zinc atoms grey, and similarly the copper metal atoms have the reddish-brown of a copper metal. This is somewhat misleading. Bulk copper metal might be reddish brown, and bulk zinc metal might be greyish-silver, but the atoms certainly don’t have those colors.

 

The day after my class, I was reading another chapter from Multiple Representations in Chemistry (blogged about here and here), which discussed misconceptions students have about bulk properties such as color. This particular zinc-copper displacement experiment was one of several such examples. The chapter described a study probing the conceptual understanding of high school chemistry students navigating between the macroscopic, microscopic, and symbolic views of chemistry (known as Johnstone’s Triangle). Not surprisingly, the students didn’t really understand what gives rise to the color of a chemical substance that they can see – and many of them thought that it came from the color of individual atoms.

 

Horrors! I wonder if my G-Chem students have the same misconception. I certainly talked about the observed color changes. When copper precipitates out of solution, you can see the brownish metal. This and the disappearance of the blue color, which I attributed to Cu²⁺ ions in aqueous solution, was connected to the reduction half-reaction Cu²⁺(aq) + 2e^- --> Cu(s). Similarly the dissolving of zinc metal and the solution turning colorless was because of the oxidation half-reaction Zn(s) --> Zn²⁺(aq) + 2e^-, and I probably said that Zn²⁺ ions in aqueous solution are colorless.

 

I think the students comprehend that individual atoms in the solid metal are not colored per se. We went over this the very first day of class when we discussed atom representations, and I occasionally make reference to it when molecular pictures show up again in chemical bonding. “Atoms aren’t colored! This is just a standard representation chemists use!” But I hadn’t discussed the ions. And the textbook picture “colors” the dissolved aqueous ions the same color as the atoms in the bulk solid. Needless to say, they are not blue and colorless for Cu²⁺(aq) and Zn²⁺(aq) ions respectively. I’m not sure what the students were thinking about the different colors of aqueous solutions because I didn’t emphasize this point. I should check but haven’t had time to do so, as my class sessions are usually very tightly scripted.

 

In one of our early G-Chem labs, the students make a standard curve for a red dye solution, and then use it to measure the concentration of red dye in a Kool-Aid solution. There is some discussion about what the absorption spectra of the dye looks like as we discuss Beer’s Law and how to set the UV-visible spectrometers to “lambda-max”. I’m not teaching lab this year, and because of remote teaching, I didn’t do my usual activity of having the students use hand-held spectroscopes and look at colored solutions to see what light gets through and what is absorbed. So I’m not sure my students have a good notion of why they observe color in a solution. The textbook picture doesn’t quite help where the molecular level picture of CuSO₄(aq) has a darker blue background than the ZnSO₄(aq) with a lighter blue background, supposedly representing the “blue” of water. I have mentioned to the students a couple of times that water is not blue even as I use my blue marker to draw beakers of solution on the white board. I might be adding to student confusion without knowing it.

 

What am I learning from all this? Next year when I teach first-semester G-Chem again, I’m going to pay attention to this issue of color and try to address potential misconceptions systematically by building it into several places in my syllabus. And I need to make a stronger connection to what students are seeing and doing in lab. I haven’t taught the lab for a number of years, and that’s caused me to “forget” to make some of those stronger connections between lecture and lab courses.

 

On top of all this, I think the students might further trip up when seeing the simple picture of metallic bonding. The chapter I read discussed students viewing the metal “atoms” in the bulk solid as metal ions – with a sea of separated electrons in the nooks and crannies between the atom/ion spheres. And these pictures have colors. Who knows what the students are actually thinking? I’ll need to pay attention to this again next year. But right now I’m just trying to finish this Covid semester without burdening the students with more details. As to the specific zinc-copper metal-displacement, I have another shot at it next semester when we cover electrochemistry in detail, so I’ll try to design an activity to probe what the students are thinking and how we can correct any misconceptions.