Letters to the editor

●  A view from a across the pond: In September 2011, you had a letter from Shawn McGovern concerning storage of chemicals and waste. The reply from Greg Friday was full of good sense. Here are some more points from the UK:

• Expiry dates are put on chemicals to inform the user that the assay cannot be guaranteed after that date. For school use, we can use many chemicals well after their expiry date although experienced chemists will know that there are problems with certain chemicals, such as hydrogen peroxide, sodium hypochlorite, hydrated calcium chloride, iron(II) salts etc. These can be handled with a yearly inspection.
• Shawn is correct in saying that storage of waste is a safety issue. Several explosions have been recorded in stores of organic waste. One that involved collecting silver residues led to a prosecution. We now advise that Tollens’ reagent be made “in situ” and the remains poured down the foul water drain immediately afterwards with lots of water.
• The foul water drain is any drain that goes to the sewage plant, as opposed to run-off drains which go to local streams or just soak away into soil; no chemicals should be disposed of down these drains. If one estimates the percentage of chemical waste from schools compared to what comes from houses, restaurants, inns and small businesses, it is very small indeed. When connecting and priming the pump in an inn, beer goes down the sink — think of it as a dilute ethanol solution. And consider all the triglyceryl stearate coming from restaurants! Oil and fat is a big problem in the sewers of central London.
• While visiting a German secondary school for the children of the British Army, I came across a complex waste system in which effluent from the science laboratories, the technology rooms, the art room and the domestic science room was collected. The process included adjusting pH and the separating of non-water soluble chemicals. It often overflowed, went wrong, and finally the company that installed it went out of business. Several thousand pounds lost to the British tax payer!
• In light of the Olympic Games, all schools and businesses with chemicals have to keep them under lock and key. The bombings on the London Underground were caused by chemicals that one could get from a school. CLEAPSS has been saying that security is important for years.

We can reduce concentrations and amounts. But let us not reduce the teaching of practical chemistry! CLEAPSS videos can be found at on the CLEAPSS Youtube channel.
(CLEAPSS is an advisory service providing support in science and technology to school and colleges in the UK.)

Regards,

Bob Worley, Lead Chemistry Adviser 
(who retired on the day he composed this and would love to come to Canada and run a session on these experiments!)

CLEAPSS The Gardiner Building, Brunel Science Park, Kingston Lane, Uxbridge, UK  www.cleapss.org.uk

●  With regard to “Is magnesium nitride formed in the combustion of magnesium?” (Chem 13 News, March 2012)

Two thoughts came to mind when reading this article: firstly, that the empirical formula looks wrong (it isn’t, but I still do a double-take whenever I see it), and secondly that magnesium nitride was crucial in confirming the existence of argon in air. Neither thought helps answer the questions raised by the article; an online search of chemical abstracts databases, however, turned up a number of early references, including an 1878 article by J. W. Mallet concerning the same demonstration.¹ In particular, Mallet notes that:  “The chief part of the metal burned is, of course, converted into the oxide, with [sufficient heat to] induce the remainder, when the supply of oxygen is limited, to unite with nitrogen.” —(Emphasis added.)

Additional papers from the early 1900s clarify the conditions required for formation of the nitride in air. For example, Eidemann and Moeser were able to obtain high yields of Mg₃N₂ by heating powdered magnesium mixed with iron or various metal oxides under carefully controlled conditions in air.² Similarly, Matignon and Lassieur reported that while the oxide formed more readily on porous magnesium (1000 times faster at 670^oC), nitrogen penetrated “deeper into the mass ... forming a pure nitride”.³ Later experiments by Makolkin et al using finely powdered magnesium confirm the different rates of reaction, reporting activation energies for nitride formation in pure N₂ between three and seven times those for oxide formation in either air or pure O₂.^4,5

It should be clear from this that the formation of Mg₃N₂ from heating or burning magnesium in air will depend on the temperature, the form of magnesium used, and the experimental configuration. Mallet, for example, first observed Mg₃N₂ using magnesium ribbon but subsequently used magnesium fillings.¹ Similarly, Eidemann and Moeser contrast results for an open crucible with one covered by a sealed lid having a small hole.² With regard to the colour of Mg₃N₂, the earliest papers appear to be consistent in describing it as a “yellowish-green”, but Makolkin et al describe both “grey” and “yellow” films depending on the ratio of MgO to Mg₃N₂ in the sample.⁴ Colour may also depend on temperature since Cotton and Wilkinson described Mg₃N₂ formed at 300^oC in pure N₂ as forming colourless crystals.⁶

In short, “results may vary” !

David Stone, Senior Lecturer
University of Toronto, Toronto ON

References

1. Reported in J. W. Mallett, Chemical News & Journal of Industrial Science, 38, 1878, 39ff  (Available as a free Google e-book, http://books.google.ca/books?id=ugYAAAAAMAAJ&pg=PA39).
2. W. Eidemann and L. Moeser, Berichte Deutschen Chemischen Gessellschaft, 34, 1901, pages 390-393.
3. C. Matignon and C. Lassieur, Chemiker-Zeitung, 36,1912, page 30.
4. I. A. Makolkin et al, Zhurnal Prikladnoi Khimii, 33, 1960, pages 824-831.
5. I. A. Makolkin et al, Trudy Moskovskii Institut Narodnogo Khozyaistva, 46,1968, pages 43-50.
6. F. A. Cotton & G. Wilkinson, Advanced Inorganic Chemistry, 4th edition, John Wiley & Sons, New York, page 283.