Thermite reaction

One of my first projects in high school was creating thermite entirely from household ingredients. This started my interest in hands-on science, and prompted me to become interested in chemistry and physics instead of my then-intended career of teuthology.

Thermite is a specific case of a general class of aluminothermic reactions. The premise of the reaction is that finely ground iron oxide, in the presence of aluminum or magnesium powder, will undergo an exothermic single replacement reaction in which the iron is reduced and the aluminum is oxidized. Because aluminum oxide is far more stable than iron oxide, the reaction releases a lot of energy– enough that the iron produced is molten. The reaction has the general form:

$Fe_2 O_3 + 2Al \rightarrow 2Fe + Al_2 O_3$

The exact stoichiometry (ie, the correct coefficients and thus relative ratios of molecules) will vary depending on which oxide of iron is used.

There are many exciting (but altogether dubious) stories reported about applications of this reaction—people have reported using the reaction for all manners of vandalism, including melting through a car’s engine block (later proven possible on Mythbusters) to breaking through padlocks and deadbolts (later simulated in an early episode of Breaking Bad). I recall watching a clip of the British television programme “Brainiac” in which two somber men in lab coats gleefully use thermite to melt through a dilapidated car.

Besides household rust, high-purity iron oxide can be obtained by isolating naturally-occurring magnetite, which I was capable of acquiring at a beach near my house. At the suggestion of various internet sources (and to the bewilderment of the lifeguards) I started wandering the beach during the day, dragging a magnet behind me in the hopes of picking up chunks of magnetite from the sand. I gradually built up my supply of iron oxide over several months. In order to generate aluminum powder, in the evenings I would “borrow” a mortar and pestle and sit in front of the TV grinding aluminum foil.

By the fall of my freshman year, I had accrued enough of each reagent to fill a film canister with thermite (roughly 100 grams). At the time I lacked a proper analytical balance, and so I recall mixing the powders volumetrically using estimates of their powdered densities and packing fractions. Fortunately, thermite is fairly robust to imperfections in stoichiometry, and so the final mixture (roughly 2:1 by volume if I remember correctly) was viable.

Thermite is notoriously difficult to light—while the reaction is very exothermic, a very large initial heat (really, the activation energy) is required to convince the reactants to move into states in which they are willing to react. The two most common methods involve adding lit magnesium ribbon and covering a layer of potassium permanganate with glycerin. The latter method is favored by demonstrators because it creates the illusion of fire coming from water—however it is also riskier, since the reaction can have a delayed start, making it unclear whether a failed demonstration can be safely approached. I decided that the best bet was to buy some magnesium ribbon online, which led me to the amusing webpage of United Nuclear.

The first ten or so times I attempted the reaction in my driveway, it sputtered and fizzled. But I found more success when I cut shavings of Magnesium ribbon and placed a pile of them on top of the thermite, and then stuck a long Mg “fuse” into the mixture that extended down the full length of the reaction vessel (a clay flower pot). The idea was that, if the reaction failed to propagate down, the still-burning magnesium backbone would help re-ignite it (I later learned that this is similar to how some of those irritating no-blow birthday candles work—Magnesium powder in the wick prevents them from cooling below the flash point of the cotton).

I could tell the reaction had succeeded by the vicious hissing that began to emanate from the flower pot. Soon fire was starting to spit out over the rim, depositing small droplets of molten metal that slowly sunk into the asphalt of my driveway. By the time I returned, a large chunk of formerly-molten iron was glowing in the charred remnants of the pot.

I later dropped my souvenir piece of iron and found that a large bubble had formed inside it. Had the iron remained molten long enough for this bubble to reach this surface and pop, a substantial quantity of molten iron could have been thrown from the flower pot because the pot was undersized relative to the amount of thermite. Inside the chunk of iron was a nice set of crystals—a sign that the iron produced by the reaction was somewhat pure and that it had cooled slowly.

I later understood that it was a lucky decision to use magnetite for the reaction—rather than being a single oxide of iron, like red ferric(iron III) or black ferrous(iron II) oxide, magnetite is a crystalline complex of both, denoted by IUPAC as FeO FeO·Fe2O3 or iron(II,III) oxide. The crystal produces the hottest thermite of any type, yet it can be isolated from beach sand.

Some details of my reaction, as well as pictures and video, can be found here.