Science in Pictures: The Oxidation of Ammonium Dichromate


You don't see this reaction demonstrated in schools anymore because chromium compounds are cancer-suspect agents. It's a shame because not only is it visually appealing to students, but it also demonstrates several concepts that are covered in physical science and chemistry courses.


In the first picture we see an orange powder: that's ammonium dichromate [(NH4)2Cr2O7]. To get it to react we have inserted a magnesium strip into the orange pile. When we light the magnesium, it will react and release heat, which will activate the breakdown of ammonium dichromate. On its own, solid (NH4)2Cr2O7 cannot absorb enough energy from the environment, so it has to be "kick-started" into reacting. Eventually the second reaction will release even more heat than the oxidation of magnesium, but nothing will happen if it is not first supplied with the required activation energy.



Now the match is lit, which will ignite the magnesium. While we wait for the main reaction to get going, we will investigate why ammonium dichromate breaks down in such a spectacular fashion. Cr2O7-2 , the dichromate ion, is rich in oxygen, so the bonded metal is said to be in an unstable and high oxidation state. If given the opportunity to snatch electrons, the dichromate ion will not hesitate and become a more stable ion such as Cr+3. Now it just so happens that another part of the same ammonium dichromate crystal has just what the dichromate seeks. We’re talking about ammonium: NH4+1. It would be energetically better off as plain N2. The price? Simple. Just lose electrons to Cr2O7-2.


The excess energy that Mg and oxygen release in reacting appears as a blinding flash of light. There is also heat released, and this will get the solid ammonium dichromate to react. Soon the dichromate ions will oxidize the ammonium within the orange powder, and the chromium atoms will be reduced to ions of a lower oxidation state:

The overall reaction can be simplified to :


 (NH4)2Cr2O7à 4 H2O(g) + N2(g)  + Cr2O3(s)


Unfortunately the balanced equation does not reveal that each chromium ion receives three electrons from each ammonium ion.


The reddish core suggests that the main reaction has begun. There are sparks in the crater of the chemical volcano, which spews out a greenish ash. The latter is chromium III oxide, containing the more stable ion and non carcinogenic ion of chromium. Notice that the side of the container, which we have used to prevent the carcinogenic ash from contaminating the room, is now steamy. Water is one of the products of the reaction, but the intense heat vapourizes it, and the increase in pressure caused by the escaping steam propels the ash upwards. Although invisible, nitrogen is also emitted from the hot crater.

(If the lid is almost completely covering the container, the steam will drive air out. Upon cooling, the condensing steam will cause a signinficant reduction in pressure if the lid is placed over the container's lip shortly after the steam is seen coming out. It will then be difficult to remove the lid as atmospheric pressure's push will not be countered by the lower pressure inside the container.)



Sadly, the reaction is complete. The glass top has been carefully removed, as it is extremely hot to the touch. There is no orange dichromate left over; all that remains is Cr2O3. Although it weighs less than the starting material ( because, after all, we have “lost” the mass of the escaping gases), it seems more voluminous than the starting material, creating the illusion that our mountain has grown like a true volcano. The truth is that the green crystals are just larger, and allow more air space in between them. This reaction needed a little spark to get going, but it is one of nature’s classic spontaneous reactions: it is highly exothermic and a servant of entropy, the universe’s measure of disorder.