Sublimation of iodine: Rise and fall of a misconception

For years I believed that solid iodine, I₂, only underwent sublimation when heated at atmospheric pressure. I was so wedded to this idea that I performed the demonstration for years. A small beaker containing a few grams of iodine, when heated in a fume hood, appeared to produce plenty of dense, purple vapour, which obscured the (purportedly) solid iodine at the bottom.

My faithful students listened to my schtick, noted the observation and grouped solid iodine with notable sublimators, such as solid CO₂ — dry ice.

I was happy to continue in this vein until I happened across the phase diagram for iodine when preparing a PowerPoint lesson. Either the diagram was wrong, or iodine does not undergo sublimation at standard pressure. (A phase diagram of iodine is shown below).

Given that empirical evidence is the trump card, I tried the demonstration after school. This time, after heating iodine crystals over a hot plate for a few minutes and observing the purple fumes, I picked up the beaker and quickly poured the contents onto a sheet of white paper.

Surprise, surprise! The iodine formed a splatter pattern, indicative of the presence of a liquid. The phase diagram was correct — and I was wrong.

For years I led my students to believe my misconception. A quick search — YouTube included — revealed that I am not the first person to realize this misconception.¹

Aside from wanting to send an apology letter to over a thousand people, what I learned about the non-sublimation of iodine shows me the need to constantly re-evaluate, with a critical eye, what I do and how I do it.

I am not perfect — and never will be.²

Notes and references

1. Marina Stojanovska, Vladimir M Petrusevski, Bojan Soptrajanov
   1. The concept of sublimation — Iodine as an example, Emergent Topics in Chemistry Education, March 2012.
   2. Addressing Misconceptions about the Particulate Nature of Matter among Secondary-School and High-School Students in the Republic of Macedonia, Creative Education, 2012. Vol.3, No.5, pages 619-631.PDFs of both articles can be found with a Google search. Article (a) has a short video-clip worth watching.
2. If you doubt this, let me put you in contact with my spouse.

Michael P Jansen <mjansen@crescentschool.org>
Crescent School, Toronto ON

Phase diagram of iodine

Chem 13 News asked Lew Brubacher, past editor, to comment on Mike’s discovery. Lew prepared a phase diagram below (Fig. 1) along with comments on the following page. Since this diagram is relevant to both articles, we have included it here. Before you read Lew’s comments on the next page, take a while to study the phase diagram and consider what it reveals.

When I was asked to comment on Mike’s article (above), I looked for a phase diagram of iodine. Among the few I found, two had significant errors. So I prepared the one shown in Mike’s article at left. As Mike notes, iodine can indeed be a liquid at atmospheric pressure between 113.7 ^oC and 184.3 ^oC.

But as I thought about phase diagrams and Mike’s article — and having not worked with phase diagrams for more than 20 years(!) — a new, admittedly naïve, question arose. In Mike’s experiment, all three phases are present simultaneously, but how can this be when the phase diagram shows that this can happen only at the triple point? It slowly came back to me that the conditions for which a phase diagram applies are not the conditions of Mike’s experiment.

Phase diagram of iodine.

A phase diagram describes the situation in which one is observing:

1. a single chemical substance,
2. in a closed system,
3. under equilibrium conditions.

Mike’s experiment

1. involves not a single substance (it also contains air),
2. is an open system, and
3. is not at uniform (equilibrium) temperature.

To elaborate: When Mike puts the flask containing solid iodine on a hot plate (at one atmosphere pressure), the temperature will not be uniform throughout the flask. At the bottom, if the temperature is 113.7 ^oC or higher, there will be liquid iodine, though, as Mike notes, it might be difficult to see it through the intensely coloured iodine vapour. In any event, it is not appropriate to try to apply the phase diagram, in detail, to the behaviour of the iodine in this flask. Here, we can have solid, liquid and vapour present simultaneously, even though the system is not at the triple point — because the system is not a single component, it’s not at equilibrium, nor at uniform temperature and pressure throughout.

But having disposed of that red herring, we are left with an excuse to recall what a phase diagram does tell us. First, a word about constructing the phase diagram in Fig. 1 at left. A search showed that there is some disagreement on the basic data.

I chose to use these data from Wikipedia:

• triple point: 113.5 ^oC at 12.1 kPa (point 1 on phase diagram)
• melting point: 113.7 ^oC at 101.3 kPa (point 2)
• boiling point: 184.3 ^oC at 101.3 kPa (point 3)
• critical point: 546 ^oC at 11,700 kPa

Also, from the old International Critical Tables, I found these densities — and the corresponding molar volumes, V_m:

• solid: 4.93 g/cm³ at 20 ^oC (V_m = 51.5 cm³/mol), and 4.92 g/cm³ near the melting point (V_m = 51.6 cm³/mol).
• Liquid: 4.00 g/cm³ near the melting point (V_m = 63.5 cm³/mol).

Of interest also is the vapour pressure of solid iodine (Table 1).

1. From International Critical Tables, Vol 3, page 201
2. This is the melting point, according to this source

A phase diagram of a single substance describes the behaviour of that substance as the pressure and temperature are changed. Imagine a sample of solid iodine in an idealized syringe that can withstand any temperature and pressure, and is transparent, so you can see what is happening inside. No other substance is present — no air, water, etc. The system must be at equilibrium, with a uniform temperature and pressure throughout. Suppose that you start with the syringe at 101.3 kPa (1 atm) and 20 ^oC, which is point 4 in Fig. 1. All of the iodine is in the solid state. Now slowly warm the sample, adding energy at a constant rate, while the pressure stays constant. When the temperature gets to 113.7 ^oC (point 2), the temperature stops rising as the added heat is used to melt solid iodine. When melting is complete the temperature resumes its rise until it gets to 184.3 ^oC (point 3), and the liquid begins to evaporate. When vaporization is complete, the temperature again begins rising. In this system, we see iodine vapour only when the temperature is at 184.3 ^oC or higher.

In contrast, when we observe a sample of solid iodine in the bottom of an unsealed glass flask, at 20 ^oC and 1 atm pressure, the system is different from the idealized system just described. There is air in the flask, and we see a faint colour of iodine vapour. Although the vapour pressure of iodine is very low (0.027 kPa, or 0.20 mm Hg), the vapour is coloured intensely enough that we can see it. Most of the 1 atm pressure is due to the air — approximately 79 kPa N₂, 21 kPa O₂, 1 kPa Ar, 0-2 kPa H₂O(g), 0.04 kPa CO₂ and 0.027 kPa I₂. If the flask is unsealed, the iodine vapour can escape. Eventually, the iodine sublimes away, just like an ice cube in a freezer.

In another application of the diagram, imagine starting slightly to the right of point 2, and raising the pressure, on a line rising vertically. Initially, the iodine is liquid, but will change to solid at some point. That’s to be expected, by Le Châtelier’s principle, because the molar volume of solid iodine is smaller than that of the liquid. In the case of water, in contrast, the molar volume of the solid (ice) is greater than of the liquid; thus, in water’s phase diagram the solid/liquid line has a negative slope.

Lew Brubacher
Department of Chemistry
University of Waterloo, Waterloo ON