The Chemistry of fireworks is a very interesting field. A look at an amusement park or a baseball game, or such events such as New Year’s Eve or Christmas Eve are just a few ways to show how much fun that comes our way from fireworks. This however is an intricate science requiring application of physical science. For instance, to produce a red chrysanthemum spray and an accompanying explosion requires certain components and materials such as an oxygen producer, color ejector, binder, fuel and propellants. There are three conspicuous and identifiable forms of energy produced by fireworks. These are heat, light and sound. The loud sound experienced in such events is attributed to the rapid release of energy into the atmosphere; hence the air expands at a greater rate than the speed at which sound travels. Therefore, a sonic boom which is a shock wave is produced.
The chemistry behind fireworks is a series of oxidation and reduction reactions which result in the desired sound and light. This happens as propellants push the firework into the sky. Oxidation reactions ensure that the oxygen needed to exhaustively burn the mixture of reducing agents and excite the atoms in the light-emitting compounds is produced. Oxidizers used such as chlorates, nitrates and percolates and reducing agents such as carbon and sulfur are available of the shelf for home-made users. The combination of reducing agents with oxygen is there responsible for the energy dissipated during the reaction. Black powder, which mainly contains nitrates, is the most used oxidizer (Conkling, 1985). A look at an explosion under the use of potassium nitrate so as to provide nitrate ions (NO3-) after decomposition can be represented as: Potassium nitrate potassium oxide + nitrogen gas + oxygen gas.
2KNO3 K2O +N2 + 2.502
The reaction is more controlled since when reacting, nitrates only release two in every three oxygen atoms, hence, the reaction is not exhaustive and vigorous since not all the oxygen atoms are actively used up. However, nitrates do not provide enough power to propel the firework into the sky and also ignite the package. Therefore, they cannot be used in star explosions since they cannot produce temperatures high enough to energize most color metal salts.
Star reactions need a temperature ranging from 1700 to 2000°C. This was enabled by the Italians in the 1830’s whereby they came across more explosive oxidizers, chlorates (ClO3-), which give up all their committed oxygen atoms upon reaction. This can be illustrated by the equation below which is highly spectacular, vigorous and releases more energy.
2KClO3 2KCl + 3 O2
Potassium chlorate potassium chloride + oxygen gas.
However, chlorates have the major demerit of being highly unstable, hence they can be dangerous to handle. On the merit side, chlorate can be easily ignited. For instance, dropping them on the ground can lead to a major explosion. This is since chlorates have the maximum potential of bonding with four oxygen atoms but they however bond with three. The fourth oxygen atom is left free, unsaturated and reactive. This makes chlorates better oxidizing agents. Further, in comparison to the slow-burning rate previously availed by nitrates, chlorates provide a faster reaction leading to a loud and exceedingly dangerous explosion. This was solved by the use of perchlorates which are more stable when releasing oxygen. The oxygen atoms in perchlorates are fully bonded hence stable. When reacting, perchlorates are able to release all their oxygen atoms. (Russell, 2009)
KClO4 KCl + 2O2
Potassium perchlorate Potassium Chloride + oxygen gas
So, perchlorates are not only more stable, but more oxygen-rich than chlorates. They, like chlorates, produce more vigorous reactions which produce hot, rapidly expanding oxygen atoms than nitrates in their star compartments.
Carbon and sulfur in charcoal are the most common reducing agents. They are contained in black powder and react to produce carbon dioxide and sulfur dioxide respectively.
Oxygen + sulfur Sulfur (IV) Oxide gas
O2 (g) + S(s) SO2 (g)
Oxygen + carbon Carbon (IV) Oxide gas
O2 (g) + C(s) CO2 (g)
The magnanimous amount of energy released in these reactions and the hot rapidly expanding oxygen gas provide a basis for propulsion and consequent explosion.
The chemistry of fireworks, so as to come up with a varying degree of colors has generated a lot of interest. Color is generated through two mainstream ways: Incandescence and luminescence. Incandescence entails the production of light by means of heat. When a substance glows as a result of heat, it first emits infrared wavelengths, then red light. Orange light is then produced as the object becomes progressively hotter, followed by yellow and finally white light. Under a controlled environment, the glow of reducing agents such as charcoal can be regulated at a certain temperature, hence emitting a particular color at the desired time. Temperature regulators that are most common are magnesium, aluminum and titanium. Luminescence is the production of light through other means other than heat. These can therefore occur at colder temperatures lower than room temperature since it is independent of any heat. An electron in an atom is first excited and destabilized by absorption of energy. The atom is then relegated to a lower energy state hence releasing the energy within via photons, the basics of light. The energy possessed by these photons consequently determines its wavelength or color. A major challenge in producing color through luminescence is that some salts used are unstable at room temperatures such as Barium chloride. Therefore, this problem must be solved by use of a combination of these salts with more stable compounds such as chlorinated rubber. For instance, in the combustion of the pyrotechnic composition between barium chloride and chlorinated rubber, a green color is produced. Other salts such as copper chloride which gives a blue color must be regulated not to attain high temperatures yet the brightness of the resultant blue color must be achieved. (Pressroom, 2010)
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