# Vive la science!

I just had a moment. Not a moment, but a MOMENT.

A super-moment-of-science moment.

It all started with a pretty lacklustre, results-wise, experiment that I — and most of us (probably) — have our students perform. It concerns the determination of the empirical formula of magnesium oxide. I have my students do this experiment free of a cookbook. Preparation involves a discussion during the previous class, followed by some pre-lab questions. Students place a strip of freshly cleaned and weighed magnesium in a crucible, and heat over a Bunsen burner flame until the magnesium reacts completely. I circulate during the experiment, gently raising each crucible lid to admit more oxygen without making it look like a new pope has been elected.1

There are more involved procedures for this experiment,2 but I don’t use them. They are heavily cookbook-based; most students wouldn’t know what they were doing or why. Afterwards, students carry out calculations; we put class data into a spreadsheet. We analyze our less-than-perfect results; we propose sources of error, which we investigate to see if they make sense.

Analysis of the class’s data reveals a ratio significantly different from Mg1O1. We have a productive discussion about possible sources of error and how our empirically determined values may or may not support these hypotheses.

Year after year, something continued to nag at me about this reaction. The so-called magnesium oxide produced in the crucible didn’t look nice and white, as magnesium oxide should. It was grey, with a noticeable amount of black. Could this be magnesium nitride, I thought? But Mg3N2 is a greenish-yellow powder.3

So I did what many busy teachers do. I forgot about it. Until recently, that is.

My post-lab discussion happened to coincide with a reading of Yehoshua Sivan’s excellent article in a recent Chem 13 News.4 You really should read his paper, but Yehoshua reminds us that magnesium burns in CO2 according to the reaction represented by:

2 Mg(s)  +  CO2(g)  →  2 MgO(s)  +  C(s)

The magnesium that is heated in the crucible over a Bunsen burner flame is surrounded by carbon dioxide, a combustion product of methane. No wonder the material left in the crucible after the reaction looks grey-ish, with black — it contains carbon!

Two beakers provide a makeshift combustion chamber.

I confirmed this by repeating Sivan’s burning of magnesium in air and in the CO2-rich flame of a Bunsen burner.

I took things one step further by modifying the apparatus for the determination of the empirical formula of magnesium oxide, as shown. I weighed two empty 1-L beakers and a cleaned piece of magnesium ribbon. Immediately after igniting the magnesium, I inserted it into the opening between the 1-L beakers. This makeshift combustion chamber filled with pure white magnesium oxide. After combustion, I taped over the opening and let the magnesium oxide “smoke” condense and settle. I’ll make a long story short: Mg1O1.08.

I plan to keep the original, albeit flawed, experiment, but I will perform the “new” experiment as part of a post-lab discussion.5

Vive la science!6

## References and notes (website accessed March 2015)

1. Magnesium oxide, the desired reaction product, can escape as white smoke. This traditionally indicates that the conclave of cardinals has elected a new pontiff. (Apologies to non-Catholics.)
2. Clancy et al, Chemistry 11, McGraw-Hill Ryerson, Toronto, 2011, page 282.
3. https://en.wikipedia.org/wiki/Magnesium_nitride
4. Yehoshua Sivan, “Burning magnesium in a Bunsen blame and other flame experiments”, Chem 13 News, February 2015, pages 12 – 13.
5. Interested readers may contact the author for a student- and teacher-ready lab handout.
6. And vive le Chem 13 News!