More from my trip into the archives earlier this week.
From Nov. 20, 1978:
SIR: We would like to alert persons to possible hazards involved with the rather common laboratory procedure of dissolving electrophoresis polyacrylamide gels with hydrogen peroxide, in order to measure radioactive species by scintillation counting.
Recently a very violent explosion occurred in one of our laboratories which caused complete destruction of a hood and moved a cinder-block wall located 30 feet from the blast. Kick-out panels and glass were blown out of the laboratory and chemicals on shelves in the adjacent laboratory were knocked to the floor. Fortunately the blast occurred when the labs were vacant, or otherwise severe injury or loss of life would certainly have occurred to personnel.
The blast apparently occurred due to the formation of explosive peroxides formed from the solubilization of polyacrylamide gels and subsequent counting procedures. The procedure used was basically the following: Polyacrylamide gels (1 cm2) were dissolved with the addition of 0.6 ml of 30% H202, and the resulting solution was added to a scintillation cocktail consisting of a 1:1 mixture of toluene and 2-ethoxyethanol along with scintillation fluors. After the samples were counted for 14C, the contents of all scintillation vials were pooled and concentrated over low heat on a hot plate in the hood. Eventually the radiological safety officer was to dispose of the material. Material had been accumulating in the hood for three to four weeks.
Addition of hydrogen peroxide to the polyacrylamide gels could result in the formation of peracids azo- or nitro-compounds. This mixture was then added to the toluene-ethoxyethanol cocktail, and hydrogen peroxide not used in solubilization of the gel could form explosive adducts with the ether. Tests of a commercial cocktail mixture showed that peroxides were present even before the hydrogen peroxide was added.
The procedure used for dissolving the gels is used by many laboratories and had been used for three years in our labs without incidence. We recommend that either alternate methods be used to solubilize the gels or that the peroxides be immediately destroyed after scintillation counting.
–Dennis W. Darnall, Professor of Chemistry, New Mexico State University, Las Cruces
And a response, from May 21, 1979:
SIR: This is in response to Prof. D. W. Darnall’s chemical safety letter on “Explosive peroxides” in C&EN, Nov 22, 1978, page 47.
It is by no means necessary to invoke the formation of an organic peroxide to understand the explosion described by Darnall. Mixtures containing strong hydrogen peroxide and soluble organic matter have been known to be explosive for many years. See, for example, E. S. Shanley and F. R. Greenspan, I&EC, 39, 1536 (1947). The relevant point is that solutions of combustible matter in strong oxidants, if metastable, may have explosive properties entirely analogous to those of organic compounds carrying strong oxidizing substituents. Thus, a stoichiometric quantity of any organic compound dissolved in strong hydrogen peroxide may yield a solution with explosive power and sensitivity equivalent to those of nitroglycerin.
In general, explosive behavior is noted with hydrogen peroxide compositions only when the weight ratio of H202 to H20 is more than one. Darnall reported the use of 30% H202; his solutions as made were, therefore, outside the range of explosive compositions. However, the evaporation step reported by Damall caused loss of water and solvents. If the rate of loss exceeded the decomposition rate of the hydrogen peroxide, the composition would sooner or later enter the explosive range. Initiation might have been thermal. As concentration and temperature increased, the rate of some oxidation reaction increased, finally leading to a self-accelerating, i.e., explosive reaction.
Good practice in dealing with peroxide-containing residues is based on destabilization and dilution of the peroxide. It is frequently adequate to pour such residues into an excess of sodium carbonate solution. The alkaline conditions tend to hydrolyze organic peroxides which may be present, and to favor orderly decomposition of hydrogen peroxide residues.
–Edward S. Shanley, Winchester, Ma.