Haber Process Essay Example

Introduction

This process involves the production of ammonia. The process uses hydrogen and nitrogen as reactants. The Haber process is an artificial method of nitrogen fixation. The process obtains hydrogen from natural gas while the air acts as the source of nitrogen. The reaction of these two gases is reversible and leads to the production of heat. The reaction is as shown in the equation below.

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It is essential to note that the forward reaction of the Haber process is exothermic while the reverse reaction of the same process is endothermic. The Haber process is commercially used to produce ammonia. Vojvodic et al. (2014) elucidate that nitrogen and hydrogen are mixed in the ratio of 1:3 to be in line with Avogadro's law which elucidates that at the same pressure and temperature, equal gas volumes contain equal molecule number. Imperatively, equilibrium and reaction rates are two key factors that determine the amount of ammonia produced in this process.

Temperature and pressure are useful conditions that help in shifting the equilibrium to favor production of more ammonia. The temperature required for this reaction is between 400 and 5000C. This is to shift equilibrium position to the right so as to produce maximum ammonia. Also, a pressure of 200 atmospheres is needed to make the amount of ammonia produced to be as high as possible. It is worthy to note that without the use of a catalyst, the reaction is slow. Therefore, the reaction involved in the Haber process uses finely divided nickel or iron catalysts to increase the surface area over which hydrogen and nitrogen gases react. For this reason, the increased reaction surface helps in speeding up the reaction rate of this process. However, it is worthy to understand that catalysts have no effects on equilibrium. The Ammonia produced is then separated from hydrogen and nitrogen in the reactor by the liquefying it under high pressure.

Le Chatelier's Principle and Reason for Decreased Ammonia Yield due to Increased Temperature

Le Chatelier's Principle

This principle is also known as the equilibrium law. This law enables the prediction of change effect on chemical equilibrium conditions. Normally, a chemical reaction's equilibrium state depends on factors such as concentration, pressure, and the temperature that are present within that system. According to Owen (2015), "When any system at equilibrium for a long period of time is subjected to change in concentration, temperature, volume, or pressure, then the system readjusts itself to partly counteract the effect of the applied change and a new equilibrium is established." Thus, it is imperative to note that any condition that results in status quo change prompts a reaction that opposes the responding system. This process aids in the manipulation of reversible reaction outcomes, usually with the aim of enhancing the yield of these reactions. Le Chatelier's Principle finds application in Haber process, a large-scale process for manufacturing ammonia gas.

Explaining why Ammonia Yield Decreased as a Result of Increased Temperature

The reverse reaction of Haber process has more molecules than that of the forward reaction. Since the forward reaction in Haber process is exothermic, an increase in temperature raises the system's heat content. This implies that the system consumes more heat energy as it tries to shift its equilibrium to the left. This leads to the production of less ammonia as the system favors a shift of equilibrium to the reactant side of the chemical reaction equation. This explains a decrease in ammonia when Haber increased the temperature in the reactor. Another explanation for the decrease in ammonia yield is that increase in temperature increases the volume of the gas inside the vessel. This consequently leads to a decrease in the pressure of the gas in the vessel. The resultant low pressure favors the reverse reaction, thus leading to a decrease in ammonia yield.

How the Use of a Catalyst Speed up the Production of Ammonia and the Effect on Quantity of Ammonia at Equilibrium.

Catalysts are used in a chemical process to increase reaction rates. They provide an alternative path to the formation of the end product. One of the paths towards the formation of the end product is by lowering activation energy required in the non-catalyzed reaction. An iron catalyst is used in the Haber process to help increase the rate of production of ammonia. The catalyst in the Haber process increases the reaction rate so that ammonia and nitrogen gases while in the reactor are able to reach the dynamic equilibrium faster. The iron catalyst ensures the hydrogen and nitrogen molecules do not need their respective translational degrees of freedom. The effect is a reduced activation energy hence the forward reaction is faster (Vojvodic et al., 2014). The hydrogen and nitrogen molecules can thus dissociate more easily before recombining to form ammonia. N2 and H2 bonds are broken more easily in the presence of a catalyst. The process can occur at lower temperatures. However, extremely low temperatures are not desirable because a little energy is required to dissociate nitrogen and hydrogen molecules even when there is a catalyst. Essentially, the catalyst speeds up the reaction rate by lowering the energy required to attain equilibrium position, otherwise referred to as activation energy. Importantly, the net enthalpy change for the reaction that leads to the formation of ammonia does not vary in the presence of a catalyst. The catalyst has no influence on the equilibrium position. The catalyst only speeds up the reaction rate, but do not lead to the production of more ammonia in the equilibrium state. A catalyst does not, therefore, alter the equilibrium state of the reaction, but simply increases the reaction rate. Thus, the quantity of ammonia at equilibrium does not change in the presence of the iron catalyst.

How the Vessel's Pressure Affect Equilibrium

The pressure change in the vessel can be due to changes in the volume of the gas inside the vessel. As a matter of fact, according to Boyle's law, the volume and pressure have an inverse relationship. The pressure change inside the vessel will result in a shift in the equilibrium of the reaction between nitrogen and hydrogen. To begin with, decreasing the pressure in the vessel shifts the reaction's equilibrium to the left since product side has fewer molecules as compared to the reactant side. Thus, to counteract the pressure decrease, the system shifts its equilibrium to the side exerting greater pressure. On the other side, high pressure inside the vessel makes the equilibrium to shift to the right. According to the law of equilibrium, "When a chemical system at equilibrium is disturbed, it returns to equilibrium by counteracting the disturbance." Thus, Le Chatelier's Principle explains that an increase in pressure makes the system to favor the reaction that generates fewer molecules while a decrease in pressure will favor the side of the reaction that generates more molecules (Lan & Tao, 2013). Consider the reaction that produces ammonia shown below.

It is worthy to note that the equation's left-hand side has 4 molecules for every 2 molecules on the equation's right-hand side. For this reason, since the product of the forward reaction, ammonia, has fewer molecules, the system will produce more of it as it tries to respond to the pressure increase. However, it is essential to understand that extremely high pressures can make vessels to burst. Moreover, building a vessel that withstands high pressures is expensive. Thus, the compromise pressure for the Haber process is 200 atmospheres. However, it also worthy to mention that a change in pressure has no effect on a gas-phase reaction in which both sides of the equation have the same number mole number.

Impacts of the Haber process on the Earth's Population Size

The Haber process is primarily used in the production of ammonia. Ammonia produced in this process is used to produce various items that both plants and animals consume. Basically, ammonium products directly or indirectly affect the human population. Ammonia is utilized in the production of fertilizers such as ammonium nitrate, urea, and anhydrous ammonia. These fertilizers are among inputs, together with pesticides, used to increase the productivity of the land. A higher increase in land productivity ensures a higher amount of food available for human consumption. The Haber process can thus be regarded as a means of raising land productivity to feed a large population. The human population will therefore increase because of more food availability.

Nevertheless, the Haber process has also contributed to some poor climate conditions which have necessitated a call for green revolution. According to Gregorich, Janzen, Helgason, and Ellert (2015), ammonium compounds have led to the production of synthetic products which have vastly affected the atmospheric composition. People have since moved from the more climate-friendly organic agriculture to the new inorganic mechanism. The release of excess ammonia in form of gases from fertilizers may be deposited on natural habitat which can bring unwelcome consequences. The opportunity cost of opting for inorganic farming will soon catch up with population when the climate change will become unbearable. Thus, the Haber process has partly contributed to a reduction in population.

Moreover, the Haber process affects bacterial life which indirectly affects human population. The production of fertilizers results in lowering of soil acidity which affects some plant lives. When plants are grown in too acidic soils, their productivity reduces which further affects the lives of organisms dependent on such plants. The population can thus reduce due to the presence of leached fertilizers which are indirect products of the Haber process. Furthermore, ammonium products like ammonium nitrate can be used to produce explosives. Explosives can be used in wars or terrorist attacks which would reduce the human population.

The Haber process and Aquatic Life

Haber process helps in nitrogen fixation by converting unreactive nitrogen into chemicals that organisms use for protein building. Usually, bacteria fix nitrogen using enzymes at room pressure and temperature. However, the industrial manufacture of ammonia needs lots of energy, leading to global warming. In addition, the Haber process results in the production of oxides of nitrogen. The use of fertilizer to fix nitrogen for plant use results in eutrophication pollution which reaches water bodies. The ramification for this is that eutrophication favors the growth of algae which die off soon after they have exhausted all the available nutrients. Aerobic bacteria then decompose the dead algae. It is prudent to understand that the decomposition process uses a lot of oxygen. This implies that aerobic bacteria deprive the aquatic life of oxygen. Consequently, this leads to the death of aquatic life, hence a collapse of the aquatic ecosystem.

Also, the Haber process leads to the formation of fertilizers which negatively affect aquatic life. When fertilizers leach onto drains and water bodies, they make those water bodies too toxic to accommodate aquatic organisms. The toxicity of such water bodies leads to the growth of more algae. The rapid growth of algae will likely block sunlight reaching the aquatic ecosystem. When there is inadequate light entering water bodies, the aquatic plant life will be at risk as plants will not be able to produce their food. Aquatic animals will also likely be extinct because they mainly depend on plants as their primary producers. Basically, the Haber process can negatively affect aquatic life by virtue of contributing to the formation of fertilizers which can result in the death of aquatic organisms. Additionally,...