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The reaction, $\ce{H2 + I2 <=> HI}$ is often used as an example for equilibrium reactions. It is however a difficult reaction to show in a school laboratory: we are not equipped to work with flammable gas at $\pu{450 ^\circ C}$ and high pressure.

It would be great if I could show my students what this process looks like in real life. I have a description, with drawings, by Max Bodenstein, from 1896 -- slightly dated.

Does anyone have pictures or videos of the gases involved and of the equipment used to make $\ce{HI}$ from the elements and to measure reactants and products?

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  • $\begingroup$ Synthesis from elements is an industrial method. In the lab you probably want to use hydrolysis of $\ce{PI3}$ instead. $\endgroup$ – andselisk Mar 1 at 17:38
  • $\begingroup$ What's the point? It won't make a great display anyway. $\endgroup$ – Ivan Neretin Mar 1 at 18:01
  • $\begingroup$ andselisk: Thanks, but I want to show this particular equilibrium reaction. I don't need the hydrogen iodide. $\endgroup$ – Martin Mar 3 at 13:47
  • $\begingroup$ Ivan: Students will encounter this reaction in textbooks and/or exams. I think it's important to show that it's a real thing that happens to actual substances, rather than just theoretical juggling with symbols. "Great display:" there's a purple gas that turns transparent, and there must be cool gear involved. $\endgroup$ – Martin Mar 3 at 13:51
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To show an example for equilibrium reactions to your students, you can use a reaction which has visual change as shifting the equilibrium position such as color. My best bet is, use of cobalt(II) chloride equilibrium with water: $$\ce{[CoCl4]^2-(aq) + 6H2O(l) <=> [Co(H2O)6]^2+(aq) + 4Cl-(aq)}$$

The beauty her is The ion, $\ce{[Co(H2O)6]^2-(aq)}$ is pink while $\ce{[CoCl4]^2-(aq)}$ is blue. According to the Le Chatelier’s principle, at relatively low chloride concentrations (e.g., diluting the solution or adding $\ce{AgNO3(aq)}$ to remove $\ce{Cl-}$ ions), the equilibrium shifted to the right and, therefore, the solution becomes pink. However, if there is a large concentration of excess chloride (e.g., adding $\ce{Cl-}$ ions by means of adding $\pu{12 M}~\ce{HCl}$ drop wise or adding $\ce{CaCl2}$ beads), the equilibrium tends to shift to the left, making the solution blue.

In addition, similar to its sensitivity to the concentration of solutes, the equilibrium is also sensitive to temperature as well (See picture below). At colder temperatures, the equilibrium tends to shift to the right(pink color at LHS of the picture) while at warmer temperatures, it lies to the left (blue color at RHS of the picture):

TempSensitivity

Reading:

  1. http://www.rsc.org/learn-chemistry/resource/res00000001/the-equilibrium-between-two-coloured-cobalt-species?cmpid=CMP00005957
  2. https://eic.rsc.org/cpd/understanding-equilibrium-a-delicate-balance/2000012.article
  3. http://research.chem.psu.edu/mallouk/ilab/Equilibrium_and_Le_Chatelier.htm
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  • $\begingroup$ Thanks Mathew, this is nice! I'll try that sometime. I like to show [Fe(SCN)2]+ -- might be less toxic? Or starch with iodine. $\endgroup$ – Martin Mar 3 at 14:00
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Well you could create extra pressure by adding both gases in a container close the container with a moving top . As long as you pressure the top and the container's volume is decreased according to the law of thermodynamics (Boyle) you can increase the pressure at a significant level .

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    $\begingroup$ In theory, yes. In a school lab? Er, well... $\endgroup$ – Ivan Neretin Mar 1 at 18:00

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