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Chemical Monitoring and Management
The Chemical Monitoring and Management Module is part of the New South Wales, Higher School Certificate (HSC) Chemistry course studied by Secondary Students in their final year of schooling (Year 12). Students study four modules, 3 compulsory, and 1 of the 5 elective modules.
The 3 compulsory modules are:
The five option modules, of which one may be studied are:
The module "Chemical Monitoring and Management" is designed to teach students studying Chemistry:
The syllabus, created by the New South Wales Board of Studies, for the HSC Chemical Monitoring and Management module can be found here.
The key aspects of the Chemical Monitoring and Management module are outlined below, under generalised headings.
Additional recommended knowledge
Section 9.4.1: Monitoring and Managing Reaction
Chemistry as a career
The roles of Chemists in the 21st century are becoming increasingly diverse and important. Professional societies such as the Royal Australian Chemical Institute and the Institution of Engineers, Australia along with universities, companies and the Australian Government are increasing their efforts to inform society of the important role of chemists. Here is one example of a role that a chemist may be employed in.
New Product Development in Paints
The chemist would be required to develop new paints as the paint market environment changes. This chemist would need to have qualifications such as a Bachelor of Science followed by polymer science or Colloid science. The chemist would need to have a good underlying understanding of organic chemistry, solvents, polymers and colloids. The chemist would need to work as part of a larger team of chemists and scientist that would participate in the development and testing of new products such as outdoor or indoor paints.
Other Chemical Occupations
Apart from this, there are a wide variety of Chemical occupations that chemists may be employed in, Including;
Chemical Monitoring and Management of Combustion
Sometimes during chemical reactions, different products can be produced under different conditions. Chemists need to monitor these reactions so that the right products may be formed. One example of this is combustion. If fuel is burnt in a plentiful supply of oxygen, the reaction products are carbon dioxide and water. If however, the fuel is burnt in a limited supply of oxygen, the products are quite different. If combustion is incomplete carbon monoxide and carbon can also be formed. Consider this reaction happening in a car engine. It is extremely important to monitor the reaction because incomplete combustion is not only a waste of fuel, it also produces carbon monoxide, a poisonous gas that replaces oxygen in the haemoglobin of blood causing asphyxiation.
Complete combustion of octane
Incomplete combustion of octane
Complete combustion of methane
Incomplete combustion of methane
Section 9.4.2: Maximise Production
Industrial Uses of Ammonia
While toxic in gas form, Ammonia contributes significantly to the nutritional needs of the planet. In the 21st century Ammonia is used for a wide variey of purposes such as in:
The synthesis of ammonia
At pressures of around 250 atmospheres and temperatures of around 400oC, ammonia (NH3) can be synthesied from its component gasses nitrogen and hydrogen in the ration of 1:3 through an equilibrium reaction using the Haber process. The chemical equation for this reaction is
N2 + 3H2 2NH3
The synthesis of ammonia is a reversible reaction that will reach equilibrium. The forward reaction is where ammonia is produced from nitrogen and hydrogen and the reverse reaction is where nitrogen and hydrogen are produced from Ammonia. Equilibrium is said to be reached when the rate of the forward reaction is equal to rate of the reverse reaction. The forward reaction of Nitrogen and Hydrogen to form Ammonia is exothermic, releasing 92kJ/mol.
As the temperature rises, the particles move more quickly, having a higher kinetic energy. Since chemical reactions require particles to collide with sufficient energy to break their original bonds, the higher temperature (and therefore higher particle energy) helps to accelerate these reactions. The number of collisions between the reactant particles is also increased as the temperature rises and the paricals are moving far faster. By increasing the pressure of the reaction vessel to 250 atmospheres, there are physically more particles in the same space. This also increases the rate of reaction as the particles re far more likely to collide with one another.
Le Chatelier's Principle
The yield of ammonia produced through the Bosch/Haber process is influenced by the conditions inside the reaction vessel in accordance with Le Chatelier's principle which states:
"If a chemical system at equilibrium experiences a change in concentration, temperature, or total pressure; the equilibrium will shift in order to minimize that change"
In the production of ammonia, the forward reaction of nitrogen and hydrogen to form ammonia is exothermic. Therefore, the reactant molecules have a higher bond energy than the product molecules. According to Le Chatelier's Principle, as the temperature increases, the equilibrium will shift towards the side of the reaction with the higher bond energy, in an attempt to minimize this change. In the production of ammonia, this shift will result in larger amounts of hydrogen and hitrogen being produced. As the main aim of the synthesis of ammonia is to produce large amounts of ammonia for little cost, having a reaction vessel heated to 400oC seems a contradiction. However, since the reaction between nitrogen and hydrogen occurs so slowly at low temperatures, it is actually more viable to increase the temperature (and therefore the reaction rate) whilst trading off some of the total yield.
As ammonia, nitrogen and hydrogen are all gases at room temperature, one mole of each will occupy the same space, at the same temperature and pressure, no matter what their respective atomic masses are. In the production of ammonia, it takes one mole of nitrogen and three moles of hydrogen to create one mole of ammonia. Therefore ammonia, when in its gaseous form, takes only one quarter of the space required by the nitrogen and hydrogen required to form it when at the same temperature and pressure. According to Le Chateliers principle, an increase in pressure will cause the reaction to shift towards the side that takes less space in order to minimise the change. In the production of ammonia this causes the reaction to shift towards the side of ammonia. So an increase in pressure inside the reaction vessel will increase the yield of ammonia.
The Haber Process patented by Fritz Haber in 1908 requires large amounts of monitoring and management as it is a delicate balancing act between raction energy, reaction rate and equilibrium. The temperature and pressure inside the reaction vessel needs to be monitored by chemists. If temperatures are increased too much the yield of ammonia will drop in accordance with Le Chatelier's Principle. At high temperatures, the catalyst used to speed the reaction will also be damaged. If the temperature is too low, the reaction will not occur fast enough to be economically viable. Whilst increasing the pressures inside the reaction vessel will increase both the reaction rate and yield, it increates the likelihood of an explosion should the vessel suffer a fault. In the Haber Process, ammonia is drawn from an area of the reaction vessel cooled below its boiling point (-33°C) after it has formed, thus helping improve the efficiency as more Ammonia will need to be created to retain the equilibrium. The addition of more nitrogen and hydrogen into the reaction vessel also has the same effect. Both of these processes need to be monitored and managed by chemists to ensure peak efficiency is achieved. Contaminants in the reaction vessel can cause a variety of undesired effects; from degrading the efficiency of the catalyst to causing an explosion as oxygen reacts with hydrogen.
When the temperature is increased inside the reaction vessel, reaction rates increase, but the overall yield goes down. The Addition of the catalyst Fe3O4 increases the reaction rate to an economically viable level without increasing the temperature so high that the yield is greatly affected. Thus, this catalyst enable the Haber Process for producing ammonia to operate far more efficiently.
The pressure of the products inside the reaction vessel is increased to around 250-300 atmospheres. This has a dual effect on the reaction. The increased pressure will casue the equilibrium reaction to shift towards the side with a smaller gaseous volume. As there are only 2 moles of ammonia produced from 3 moles of hydrogen and 1 mole of nitrogen, the reaction will shift toward the ammonia side, thus increasing the efficiency of the whole reaction. The increase of pressure inside the reaction vessel also increases the number of collisions between the reactant molecules, enableing them to react faster, increasing the overall reaction rate.
Development of the Haber Process
The Haber Process was developed in Germany by Fritz Haber in 1912. Germany was preparing of World War One, so Haber's research was of great interest to the German chemical industry, as ammonia was needed for explosives and fertilizers. At this time Germany was highly dependent on imports, with its primary source of Ammonia being from Chile.
The British naval blockade of Germany blocked Germany's supply of ammonia from Chile. However the Haber process was able to produce ammonia independent of outside supplies and so was able to assist in the production of munitions, nitric acid and fertilizers essential for food production. All of these were essential for the war effort and enabled Germany to keep fighting in World War One for far longer than would have been otherwise possible. In fact, Haber has been held responsible for prolonging the war.
|This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Chemical_Monitoring_and_Management". A list of authors is available in Wikipedia.|