Reducing Combustion Pollution with “Green” Pyrotechnics in Environmental and Green Chemistry

ACCM Concepts

This Exemplar will teach the following concept(s) from the ACS Examinations Institute General Chemistry ACCM:

• VI. C. 2. a. Enthalpies of reaction, such as those incorporated in a thermochemical equation, can be used to calculate energy flow for a given amount of reactant or product.
• VI. G. 2. Hess’s law takes advantage of enthalpy being a state function to allow the determination of enthalpy change for a reaction based on combinations of more readily measured, or tabulated, reaction enthalpies.
• IX. E. 2. Practical implications of laboratory reactions include the ability to generate species of interest (such as oxygen) or release energy (by combustion of fuels). 

I. Introduction to Pyrotechnics

Fireworks are an important staple to any Disneyland celebration, Fourth of July evening, and major boating event. The production of fireworks is an exothermic process that is produced from a combustion reaction. Although the first fireworks were originally mainly composed of gunpowder, modern fireworks contain various chemicals such as charcoal, sulfur, manganese, sodium oxalate, aluminum and iron dust powder, potassium nitrates, potassium chlorate, potassium perchlorate, strontium nitrate, and barium nitrate, etc. [1] However, these modern fireworks have a substantial environmental impact due to the release of several toxic and poisonous substances upon explosion. [2] 

Figure 1: Fireworks blooming in the night sky [3]

II. Thermochemical equations

In order to visualize and mathematically analyze the gains or loss of energy in reactions, chemists use thermochemical equations that depict changes in enthalpy. Enthalpy is described as the change in heat, usually measured in kJ/mol that can be altered by the destruction and formation of molecules through the process of reaction. If a reaction releases energy in the form of enthalpy (-ΔH), it is considered an exothermic reaction and if the reaction absorbs energy in the form of enthalpy(+ΔH), it is considered an endothermic reaction. A combustion reaction, like the one seen in fireworks, is an exothermic reaction that has a very large release in enthalpy.

2Mg_(s) +O_2(g)  ->  2MgO_(s)                         ΔH_m = -1203.2 kJ (25°C, 1 atm pressure) [4]

Numerical values of delta H can be calculated depending on what information is available.  However, the quantity of heat released or absorbed by a reaction is proportional to the amount of each substance consumed or produced by the reaction. The above equation shows that 1203 kJ of heat energy is given off for every mole of O2 which is consumed, for every 2 mol of Mg consumed, or every 2 mol of MgO produced. ΔH_m is a term that can be used as a conversion factor to enable the calculations of heat absorbed when a given amount of substance is consumed or produced. If q is the quantity of heat absorbed and n is the amount of substance involved, an equation used to calculate the change in enthalpy in a reaction is:

ΔH_m=q/n

When trying to calculate enthalpy change given two or more chemical equations, Hess’ Law can be used. Hess’ Law is a principle that states that in the general case it is always true that whenever two or more chemical equations can be added algebraically to give a net reaction, their enthalpy changes may also be added to give the enthalpy change of the net reaction.

Hess’ Law upholds the law of conservation of energy, as if it were not true, it would be possible to think up a series of reactions in which energy would be created but which would end up with exactly the same substances we started with. Hess’ law enables us to obtain ΔH_m values for reactions which cannot be carried out experimentally.

Often, Hess’ Law is used to calculate energy change using the addition of two or more standard enthalpies of formation (ΔH_f). The standard enthalpy of formation is the enthalpy change associated with the formation of 1 mol of a pure substance from its elements. This value is often tabulated at normal atmospheric pressure and a specific temperature (usually 25°C). The above equation can be manipulated as:

Mg_(s) +1/2O_2(g)  ->  MgO_(s)                               ΔH_f = -601.6 kJ (25°C, 1 atm pressure)

 

Example \(\PageIndex{1}\)

Magnesium burns in air to produce bright light and is often used in fireworks displays The combustion of magnesium can be described by the following thermochemical equation used in the above explanation:

2Mg_(s) +O_2(g)  ->  2MgO_(s)                         ΔH_m = -1203.2 kJ (25°C, 1 atm pressure)

 

Calculate the heat of reaction (in kilojoules) when the combustion of 6.54 g of magnesium occurs. Include the sign in your answer.

Solution

The given thermochemical equation tells us that the heat of reaction for the combustion of magnesium is -1203.2 kJ/mol. To find the heat of reaction when 6.54 g of magnesium is combusted, we need to convert the mass of magnesium to moles using its molar mass.

The molar mass of magnesium is 24.305 g/mol, so 6.54 g of magnesium is equal to:

(6.54 g)/(24.305 g/mol) = 0.2685 mol

Now we can use the molar heat of reaction to find the heat of reaction for the given mass of magnesium:

-1203.2 kJ/mol * 0.2685 mol = -323.6 kJ

Therefore, the heat of reaction when 6.54 g of magnesium is combusted is -323.6 kJ.

 

III. Perchlorates in Fireworks

One of the main sources of pollutants associated with fireworks comes from their use of perchlorate as oxidizers. Perchlorate, or ClO₄^-, is an inorganic anion that can combine with cations such as sodium, potassium, or ammonia to form perchlorate salts like sodium perchlorate, potassium perchlorate, and ammonium perchlorate. In fireworks, perchlorates contribute to the display of loud explosions, colored flames, and bright light. [5] Potassium perchlorate and ammonium perchlorate are the most common sources of perchlorates in fireworks due to their higher availability of oxygen per unit weight, higher stability, and lesser sensitivity to mechanical action. [6] During a firework display, there is a high likelihood that perchlorates from the fireworks are released into the air and spread over a wide area below the combustion site, as Wu et al. found that the yearly deposition of perchlorate from fireworks ranges from 0.7 to 2.6 kg in the USA. [7]

Figure 2: Fireworks are major contaminators of surrounding air, water, and soil [8]

Perchlorates released from pyrotechnics can contaminate the soil and water of its surrounding area. They can either leach into bodies of water or potentially be absorbed by nearby plants through soil moisture and accumulate in their tissues. Through water, it may also reach insects, mammals, fishes, etc. [9]  Perchlorates are also potent thyroid disruptors, as they competitively inhibit the uptake of iodide by the thyroid gland thereby decreasing the production of thyroid hormones and increasing the thyroid stimulating hormone levels in blood. A thyroid hormone deficiency may result in functional deficits such as difficulty in visual-spatial information processing as well as poor fetal and neonatal development from fetuses who obtain thyroid hormones from the mother. [10]

Example \(\PageIndex{2}\)

Ammonium perchlorate, a hygroscopic oxidizer, is sometimes used as an oxidizer for fireworks. If handled incorrectly, ammonium perchlorate can violently decompose according to the thermochemical equation below:

2NH₄ClO_4(s) -> N_2(g) + Cl_2(g) + 2O_2(g) + 4H₂O_(g)             ΔH_m = -375.6 kJ (25°C, 1 atm pressure)

The enthalpy of formation of H₂O_(g) is -241.8 kJ. Calculate the enthalpy of formation for ammonium perchlorate.

Solution

This question utilizes ideas about standard enthalpies of formation. In the given reaction, two moles of ammonium perchlorate decompose to form one mole of nitrogen gas, one mole of chlorine gas, two moles of oxygen gas, and four moles of water vapor. Diatomic gases (N_2(g), Cl_2(g), and O_2(g) here) have an enthalpy of 0kJ because they represent the elemental state. Since the enthalpy of formation of water vapor is given as -241.8 kJ/mol and diatomic gases have an enthalpy of formation= 0kJ, the enthalpy of formation of ammonium perchlorate can be calculated as follows:

ΔH_rxn = 2(ΔH_NH4CLO4) - 4(ΔH_H20) 

-375.6kJ =  2(ΔH_NH4CLO4) - 4(-241.8 kJ)

ΔH_NH4CLO4 = -295.8kJ

 

IV.  “Green” Pyrotechnics

As these pollutants are a major source of concern, firework producers are currently focusing on the development of “green” pyrotechnics. These devices substitute perchlorates for other more environmentally-friendly oxidizers such as periodates. The use of periodates in pyrotechnics reduces the amount of these oxidizers that are dispelled by the combustion reaction that occurs. [11] Pyrotechnics that are currently in development are now primarily using these periodates as their oxidizing agent instead of perchlorates, especially due to the heavy restriction on perchlorates by the Environmental Protection Agency.

Example \(\PageIndex{3}\)

The burning of fireworks releases a large amount of air pollutants, with two main products being sulfur dioxide (SO₂) and carbon dioxide (CO₂). Calculate the enthalpy of the following chemical reaction:

CS₂(l) + 3O₂(g) ---> CO₂(g) + 2SO₂(g)

Given: 

C(s) + O₂(g) ---> CO₂(g)             ΔH_f = −393.5 kJ/mol

S(s) + O₂(g) ---> SO₂(g)             ΔH_f = −296.8 kJ/mol

C(s) + 2S(s) ---> CS₂(l)               ΔH_f = +87.9 kJ/mol 

Solution

According to Hess's law, the net enthalpy of a reaction can be calculated using the enthalpies of formation of the reactants and products in the reaction.

the first principal equation remains unchanged as CO₂ is a product in the overall reaction. The principal enthalpy must be multiplied by 2 because there are 2 moles of SO₂ produced in the overall reaction. Lastly, the third principal reaction must be reversed because CS₂ is seen as a reactant for the overall reaction. After adding the total enthalpies in, the answer calculated for the chemical reaction is:

ΔH_rxn = (−393.5)  + (−593.6) + (−87.9) = -1074.2 kJ/mol

 

 

 

 

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Bibliography

1. Ravindra, K.; Mor, S.; Kaushik, C. P.;. Short-term variation in air quality associated with firework events: a case study. J. Environ. Monit. 2003, (5), 260-264.

2. Steinhauser, G.; Klapötke, T;. “Green” Pyrotechnics: A Chemists' Challenge. Angewandte Chemie International Edition. 2009,  47 (18), 3330-3347.

3. Popular Mechanics Editors. The Science Behind Your Favorite Fireworks. https://www.popularmechanics.com/sci...ireworks-work/.

4. National Institute of Standards and Technology. NIST Chemistry WebBook, SRD 69: magnesium oxide. https://webbook.nist.gov/cgi/inchi?ID=C1309484&Mask=2.

5. Strategic Environmental Research and Development Program (SERDP). Strategic Environmental Research and Development Program—alternative causes of wide spread, low concentration perchlorate impacts to groundwater. 2005. http://www.clu-in.org/download/contaminantfocus/perchlorate/Perchlorate-ER-1429-tr-1-2005.pdf.

6. Zoeller, R. T.; & Rovet, J.;. Timing of thyroid hormone action in the developing brain: clinical observations and experimental findings. J. Neuroendocrinol. 2004, 16 (10), 809-18.

7. Wu, Q.; Oldi, F. J.; Kannan, K;. Fate of perchlorate in a man-made reflecting pond following a fireworks display in Albany, NewYork, USA. Environ Toxicol Chem. 2011, 30 (11), 2449–2455.

8. Matthew Cappucci. How the weather affects fireworks displays and tips for viewing. https://www.washingtonpost.com/weath...works-weather/.

9. Ellington, J. J.; & Evans, J. J.; Determination of perchlorate at parts-per-billion levels in plants by ion chromatography.  J. Chromatogr A. 2000, 898 (2), 193-9.

10. Sijimol, M.R.; Mohan, M.; Environmental impacts of perchlorate with special reference to fireworks—a review. Environ Monit Assess. 2014, 186, 7203–7210.Brusnahan, J. S. ; Shaw, A. P.; Moretti, J. D.; Eck, W. S.; Periodates as Potential Replacements for Perchlorates in Pyrotechnic Compositions. Prop., Explos., Pyrotech. 2017, 42 (1), 62-70.

11. Brusnahan, J. S. ; Shaw, A. P.; Moretti, J. D.; Eck, W. S.; Periodates as Potential Replacements for Perchlorates in Pyrotechnic Compositions. Prop., Explos., Pyrotech. 2017, 42 (1), 62-70.