Gases in Farming

Fertile Farms cause Environmental Harm?

Farmland; peaceful, green, natural, organic, healthy …but could farms actually be harming the environment by releasing nitrogen into the atmosphere?
In order to grow a lot of plants in a given area, farmers use fertilizer, most of the time, too much fertilizer. These fertilizers are composed primarily nitrogen, phosphorus and potassium. Because the plants can’t possibly use all this nitrogen, it gets released back into the atmosphere as nitrogen oxide compounds which, depending on what layer of the atmosphere they are in, either depletes ozone or produces ozone.

Figure \(\PageIndex{1}\) UV light from the sun is blocked by ozone in the atmosphere.

Credit: NOAA

In the stratosphere, nitrogen oxides deplete the ozone that absorbs high energy ultraviolet radiation from the sun. This protects life forms on earth because ultraviolet radiation has enough energy to break bonds in many biological molecules causing serious damage to organisms. In the troposphere these nitrogen oxide compounds produce ozone which is part of the irritating smog that affects respiratory systems and is considered a greenhouse gas because it traps heat from the sun.

Most of the gas phase reactions that take place in the atmosphere are reversible reactions. This means that as reactants form products, products breakdown into reactants. Some of the gases formed by reversible reactions are harmful to the atmosphere. Nitrous oxide, N2O, is one of the most harmful of these gases because it forms nitric oxide, NO which reduces the amount of ozone in the stratosphere.

\[2 NO_{(g)} + O_{3 (g)} \rightarrow 2 NO_{2 (g)} \nonumber\]

At equilibrium, the rate that new products are formed is equal to the rate that products break down into reactants; that is there is no net change in concentration of reactants and products. This relationship is described by an equilibrium constant, Keq:

The general equilibrium expression for a reaction:

\[\ce{aA + bB -> cC + dD}\nonumber\]

is written as:

\[\text{K}_c = \dfrac{[\text{C}]^c[\text{D}]^d}{[\text{A}]^a[\text{B}]^b}\nonumber\]

The brackets "[ ]" represent the concentration of the species (moles per liter or molarity). "a, b, c, and d" represent the coefficients used to balance the equation. The "c" in Kc indicates that the value of K is determined using the concentrations of each species. There are two cases when a species is not shown in the equilibrium expression:

• when it is a solid
• when it is a pure liquid or solvent1

The equilibrium position can change based on the concentrations of the chemicals involved, temperature or pressure if some of the chemicals a gas in a closed system. This shift is known as Le Chatelier’s Principle. Simply put; if a system in dynamic equilibrium is disturbed, it will react in a way to counteract that disturbance. If a substance is added, the reaction will shift to use it up. If a substance is removed, it the reaction will shift to replace it.

Exercise \(\PageIndex{1}\)

1. Consider the following endothermic reaction:

\[N_{2 (g)} + O_{2 (g)} \rightleftharpoons 2NO_{ (g)}\nonumber\]

1. If the concentration of oxygen increased, would the reaction shift to the product side or the reactant side?
2. If the temperature increased, which way would the reaction shift?
3. Nitrogen and oxygen are found in the atmosphere. At high temperatures, they combine to make nitric oxide; can you think of place this might occur?
4. Write the equilibrium expression for the equation reaction above.

2. Write the balanced chemical equation for the reaction between nitrogen dioxide gas and atmospheric oxygen to produce nitric oxide gas and ozone.

1. This reaction goes in the forward direction in the presence of sunlight. Would it be exothermic or endothermic?
2. Would you expect to find as much ozone produced during the day as at night? Why?
3. Do you think this reaction occurs in the stratosphere or the troposphere? Why?
4. What would happen to the equilibrium position of this reaction if the concentration of nitrogen dioxide were decreased?

3. Explain the importance of equilibrium position in determining the environmental impact of nitrogen oxides.

Answers:

1a. If the concentration of oxygen increased, the equilibrium would shift to the product side to use up the additional oxygen.
b. Since the reaction absorbs energy, the equilibrium would shift to the product side to use up the increased energy from the higher temperature.
c. High temperatures occur in the exhaust system of automobiles and smokestacks of manufacturing plants and provide the energy needed for the reactions of atmospheric nitrogen and oxygen to produce nitrogen oxides.
d. \(\text{K}_c = \dfrac{[\text{NO}_2]^2}{[\text{N}_2][\text{O}_2]}\)

2. \(\ce{NO2 (g) + O2 (g) <-> NO (g) + O3 (g)}\)
a. This reaction absorbs energy from the sun so it would be endothermic.
b. More ozone would be produced during the day because that is when sunlight would be available.
c. This reaction would most likely occur in the troposphere because more oxygen is available to form nitric oxide.
d. If the concentration of the nitrogen dioxide decreased, the equilibrium position would shift to the reactant side and less ozone would be produced.
3. The environmental impact of nitrogen oxides is affected by the concentration of nitrogen in the atmosphere and the temperature that these reactions occur because the reactions that produce nitrogen oxides are reversible and reversible reactions eventually establish an equilibrium. According to LeChatelier’s Principle, stress to the equilibrium (change in concentration of reactants and products or change in temperature at which the reaction occurs) can shift the equilibrium to produce more product or more reactant which in turn can produce more nitrogen oxide compounds. Nitrogen oxide compounds are considered greenhouse gases, produce hazy brownish yellow smog and can also combine with water in the atmosphere to produce acid rain.

From ChemPRIME: 9.0: Prelude to Gases

1. Professor Patricia Shapley, University of Illinois, 2011, http://butane.chem.uiuc.edu/pshapley/GenChem2/A12/1.html
2. Chemical Education Group, Purdue University, College of Science. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch16/equilib.html
3. Joy, Linda,
4. NOAA Study Shows Nitrous Oxide Now Top Ozone-Depleting Emission, August 27, 2009

Contributors and Attributions

• Ed Vitz (Kutztown University), John W. Moore (UW-Madison), Justin Shorb (Hope College), Xavier Prat-Resina (University of Minnesota Rochester), Tim Wendorff, and Adam Hahn.