A Matter Of Balance Biology Essay

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During 1885, a French chemist by the name of Henri Louis Le Chatelier created a rule for use to predict what effects occur in a reaction system when the reaction is at equilibrium. It stated that; if a dynamic equilibrium is disturbed by changing the conditions, then the position of equilibrium moves to counteract that change (Clark, 2002). This is directly applied to a wide range of chemical reactions today, specifically industrial chemistry reactions, where the most efficient means of production must be used to minimise expenses. By using this principle, industrial chemists are able to provide a constant condition for reactions that yield the highest possible amount of products, while keeping the cost to a minimum. Throughout this article, equilibrium as a whole, both physical and chemical changes, will be researched and applied using the Haber process.

The Haber process is used to synthesise ammonia by combining both Nitrogen gas and Hydrogen gas. The Haber process directly links to Le Chatelier's Principle via the need to produce an efficient and cost effective system for industrial producers by means of temperature, pressure, and catalysts to alter the equilibrium (Clark, 2002). The Haber process combines Nitrogen gas, extracted from the air, with Hydrogen gas usually obtained via Methane. The chemical equation for this system is shown below.

The above system produces heat as a product, therefore the system is considered to be exothermic, meaning that as Ammonia is produced, then simultaneously heat is given off as a by-product. By referring to Le Chatelier's Principle, it is shown that the heat produced by this system would drive the reaction backwards and therefore away from the products, and synthesizing less Ammonia. Although this is a downfall, this heat change also causes a faster rate of reaction through the system which is desirable to industrial chemists.

Industrial chemists that use the Haber process for the production of ammonia are constantly searching for ways to become economically efficient. That is to decrease the production time and maximise profit. This is all done through the use of Le Chatelier's principle. One method by increasing the reaction time is through the use of a catalyst. An iron based catalyst is the most commonly used substance in the Haber process. This iron based catalyst is very useful from economical and equilibrium points of view as the catalyst is not consumed by the reaction nor does it alter the equilibrium constant (K), it only speeds up the reactions of the system.

Because equal volumes of equal gases at the same temperature occupy the same space, according to Avogadro's Law, it can be assumed that pressure will affect the equilibrium of a system containing only gases, such as the Haber process (Clark, 2002). By using the mole ratios of the Reactants in the Haber process, it can be deduced that the ratio of Hydrogen gas to Nitrogen gas is 1:3. This means that for every mole of Nitrogen gas there are three molecules of Hydrogen gas. This information is also valuable to industrial chemists as it allows them to ensure that the more expensive of the two reactants, in this case Hydrogen gas, is completely consumed by adding excess Nitrogen gas, the less expensive of the two reactants, into the system.

In order for industrial chemists to apply Le Chatelier's principle to the Haber process, they can change, increase or decrease, one or more of three variables; temperature, pressure, or concentration.

In order to alter the equilibrium through the use of a temperature change, an industrial chemist must find the most efficient position, so that all of the reactants are consumed and produce Ammonia. Too much heat and the system will favour the reverse reaction, however too little heat and the rate of reaction won't be high enough. It is because of this, that a range of 400-450°C is used as a medium in order to maximise the amount of synthesised Ammonia through the system.

In order to alter the equilibrium through the use of a pressure change, there must be gases present in the system, which is fortunate for the Haber process, making this method very valuable to industrial chemists. Le Chatelier's principle states that if overall pressure is increased on a system, then the equilibrium will shift to favour the production of fewer molecules (Clark, 2002). The Haber process contains four molecules of reactants (1 Nitrogen gas and 3 Hydrogen gas) and only 2 molecules of the product (2 Ammonia). This means that if the overall pressure on the system is increased, then the equilibrium will shift to favour the synthesis of products. This pressure will also cause the rate of reaction to increase as the molecules are pushed closer together, allowing for an easier reaction. This means that the higher the pressure that can be achieved, the better. However, in order to be more economically efficient as well as have the highest reaction rate as possible, the pressure of the system is recommended to be set at 200 atmospheres (Clark, 2002). This compromise pressure is the most economic for industrial chemists as higher pressures require significantly more expensive pipes and vessels in order to cope with higher pressures and lower pressures reduce the reaction rate.

Example

Below is an example demonstrating the affects that the above equalising method/s have when applied to the Haber process at the set temperature of 400ËšC.

If the concentration of nitrogen gas is 3mol/L and the concentration of hydrogen gas is 4mol/L, the K value being , then the concentration of ammonia would be:

This shows that 2.38% of ammonia is synthesised in the above system. In order to increase the percentage yield of the same system, the pressure/concentration can be increased. Because the set pressure to be used on the Haber process is 200, we will instead look at the effects that changing the concentration of one of the species involved in the system has.

The concentration of nitrogen in the system is increased to 6.2mol/L, therefore the Q value is calculated using the steps below.

Because the Q value of the system is smaller than the abovementioned K value, then the equilibrium will be driven towards the products, in this case the synthesis of ammonia, and hence producing a larger yield.

The information to this point has outlined the causes and effects that the stresses of Le Chatelier's Principle have upon the Haber process. It is because of these effects that industrial chemists find it vital to equilibrium equations such as this. It allows them to predict how the system can be made most efficient and in the most economic method, and also allows them to fix possible problems within the system or the environment surrounding it.

Complication in the Manufacturing Process

One of the complications that industrial chemists may face during the execution of the Haber process is the failure of a thermostat or temperature regulating system. This is an issue when it comes to efficiency of the system as this failure can cause an increase or decrease in temperature and therefore altering from the efficient set temperature of 400ËšC. If a rise in temperature were to occur, then the reaction would be driven towards the production of the reactants, in this case hydrogen and nitrogen gases, because the system will absorb the excess heat in order to regulate the temperature. If the temperature of the system were to decrease, then the reaction rate would decrease dramatically also, however the theoretical yield of the products, in this case ammonia, would be high. The simplest method to restore the system back to an efficient equilibrium would be to change the temperature back to 400ËšC, however, if there is damage to the temperature regulation system/thermostat then this cannot be achieved. The next advised method of restoring equilibrium is to alter the concentration of nitrogen gas. Theoretically, increasing the amount of nitrogen within the system should drive the reaction towards the synthesis of products. Proof of this is shown below.

Because the equilibrium constant is unknown, it must first be calculated. The previous amounts of species in the system were noted prior to the disruption caused to the system and are shown in the table below.

ICE

N2

H2

NH3

Initial

4.4mol/L

6.2mol/L

Change

0.0478mol/L

0.0478mol/L

Equilibrium

4.3522mol/L

6.1522mol/L

0.0478mol/L

Table 1.1 - An ICE table showing the concentrations of the species involved prior to the disruption

In order to once again increase the yield of the system, the equilibrium constant, or K value, must first be calculated.

Because the ideal percentage of ammonia produced is 15%, the percentage yield of the system after failure is calculated.

Because the percentage yield is only 0.451%, it can be seen that the amount of ammonia produced has decreased as the temperature of the system increased. However, in order to resolve this issue, a reaction quotient must be calculated. The industrial chemist will aim for a reaction quotient, or Q value that obtains a suitable yield of 15% ammonia.

2.4478 mol/L excess N2 is added to the system

Because the Q value is less than the K value, it proves that increasing the concentration of Nitrogen gas will drive the reaction towards the synthesis of products, in this case ammonia. Further increasing the amount of nitrogen gas will also increase the amount of ammonia produced, however also the pressure, which in this case cannot be done due to the 200 atmospheres limit of the containment vessel. Please note that this is only a temporary solution to the thermostat problem. Once the thermostat is fixed, the temperature can be restored to 400ËšC and the ideal yield of 15% re-established with more ease.

The theoretical concentration of nitrogen gas required to reach the ideal yield, of 15% ammonia, can also be calculated.

Because 0.0478mol/L is equal to 0.451% yield of ammonia, the concentration of ammonia at 15% yield is

Therefore, in order to discover the concentration of nitrogen gas required, substitute

into the below equation.

Because nitrogen and hydrogen are both gases, they are in proportion with each other and therefore the concentrations of each can be replaced with

This means that in order to drive the reaction forward enough, in this system, to reach a yield of 15% ammonia, the total amount of nitrogen gas would need to be approximately 32.5mol/L. however, because of the 200 atmospheres limit of the containment vessel, raising the yield to this level would be very impractical. It is for that reason that the concentration of nitrogen gas would be increased by a significantly lesser amount to that, and would only be a temporary fix to the thermostat issue.

Although the purpose has changed, the production of ammonia is still executed to this date through the Haber process. Le Chatelier's principle enables industrial chemists to alter the temperature, concentration and pressure, allowing them to alter the rate of reaction and the yield of a desired product and, as shown above, can also be used to find temporary fixes to issues in the environment surrounding the system. The equilibrium constant (K) and reaction quotient (Q) can be compared to each other, allowing industrial chemists to determine whether or not a change in the system will favour the synthesis of products or reactants. Because of the wide variety of industrial chemists utilising the Haber process, in the future it will be improved so that the pressure limit of the containment vessel increases and therefore the entire system can be made even more efficient and cost effective.

Works Cited

Clark, J., 2002. Le Chatelier's Principle. [Online]

Available at: http://www.chemguide.co.uk/physical/equilibria/lechatelier.html

[Accessed 24 Fabruary 2013].

Clark, J., 2002. The Haber Process. [Online]

Available at: http://www.chemguide.co.uk/physical/equilibria/haber.html

[Accessed 26 February 2013].

Deb Smith, S. M. M. G. R. S., 2006. Chemistry In Use Book 2. s.l.:McGraw-Hill Australia.

Mombourquette, M., 2012. Chapter 12: Chemical Equilibrium. [Online]

Available at: http://www.chem.queensu.ca/people/faculty/mombourquette/firstyrchem/equilibrium/

[Accessed 7 March 2013].

Wallace, R., King, J. & Sanders, G., 1983. In: Biosphere: The Realm of Life. New York: Oxford Uni Press, pp. 523-536.

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