Chemical Kinetics: Factors That Affect Reaction Rate - Assignment Example

DateChemical Kinetics: Factors That Affect Reaction Rate Factors That Affect Reaction Rate Chemical kinetics is the study of chemical processes and the factors that affect them. It involves the critical analysis of molecules and atoms and how they behave when certain experimental conditions are altered. The developments in chemical kinetics were initiated by Peter Waage and Cato Gulberg, a duo who formulated the law of mass action. The law states that the speed of a chemical reaction is proportional to the quantity of the reacting substances. The rate of a chemical reaction is affected by the following factors. Temperature Temperature is the measure of the kinetic energy in particles. Increasing the temperature in a chemical reaction increases the kinetic energy of the molecules. As the kinetic energy increases the activation energy required for the reaction to happen is acquired easily therefore increasing the rate of reaction. Activation energy is the energy required to be reached for a reaction to take place (Boundless Chemistry). The increase in temperature increases the collisions hence the reaction rate will be increased. Pressure Pressure affects the rate of gaseous reaction. The increase in pressure means that the molecules are being compacted together hence increasing the concentration in the moment. When the pressure is increased the gas molecules collide more often and so the reaction rate increases. Liquids and solids are not affected by increase or decrease in pressure. On top of this some gaseous reactions are not affected by change in pressure (Prigigine & Wildom 2001). The increase in pressure is equivalent to increase in concentration as when the volume is reduced, the number of molecules per given volume increases. Reactant Concentration An increase in the concentration of the reactants results to an increased number of collisions which eventually lead to an increased reaction rate. The increased number of collisions is attributed to the increased number of particles per unit volume which makes the reaction rate to increase (Prigigine & Wildom 2001). Using a Catalyst A catalyst is defined as a substance that increases the reaction rate without being consumed in the reaction. A catalyst lowers the activation energy for the reaction to take place. Activation energy is the energy that needs to be overcome for a reaction to take place. The activation energy is lowered in a number of ways which include;- increasing the frequency of reactions, changing the relative orientation of the molecules in the reaction, donating electron density to the reacting molecules, reducing the intra-molecular bonding within the molecules and providing an alternative pathway for the reaction (Prigigine & Wildom 2001). Diagram showing the effect of a catalyst on the rate of a chemical reaction, noting the effect on the activation energy of the reaction Understanding Catalysts and their Role in the industry We noted earlier that catalysts are substances that speed up the rate of a chemical reaction without being consumed in the reaction. This is done by ;- increasing the frequency of reactions, changing the relative orientation of the molecules in the reaction, donating electron density to the reacting molecules, reducing the intra-molecular bonding within the molecules and providing an alternative pathway for the reaction (Prigigine & Wildom 2001). In this section we will further look in the action of catalysts in chemical reactions and the differences between homogenous and heterogeneous catalysis. Difference between Homogenous and Heterogeneous Catalysis Homogenous Catalysis Homogenous catalysis involves a series of chemical reaction where the catalyst used is in the same phase as the reactants. The phase here refers to solid, liquid or gas. A good example of the industrial homogenous catalysis is the oxidation of Sulfur IV Oxide (SO2) to Sulfur VI Oxide (SO3) using Nitrogen II Oxide (NO) as the catalyst. This was the first industrial catalyzed reaction and an ideal example of a homogenous catalysis reaction (Fogler & Gurmen 2007). SO2 SO3 Acid Catalysis is another example of a homogenous catalysis reaction. In this reaction an acid is the homogenous catalyst. The water used facilitates the formation of protons by its self ionization. This way acid are used to catalyze the hydrolysis of esters. CH3CO2CH3 + H2O CH3CO2H + CH3OH Heterogeneous Catalysis In heterogeneous catalysis the phase of the reactants is different from the phase of the catalyst. The bulk of heterogeneous catalysts are solids. They are mainly used to catalyze gaseous and liquid reactants. Heterogeneous catalysis involves the process the adsorption where molecules of different phases bind. The molecule that is binding is called an adsorb ate. The catalyst mainly acts as the adsorbate in heterogeneous catalysis. An example of heterogeneous catalysis is the Haber Bosch process in the Formation of Ammonia where iron promoted with potassium II oxide, calcium II oxide, Silicon IV oxide and Aluminum VI oxide is used. Differences between Homogenous and Heterogeneous Catalysis Homogenous Catalysis Heterogeneous Catalysis In homogenous catalysis it is difficult and expensive to recover the catalyst used. It is easy and cheap to recover the catalyst used in a heterogeneous reaction Thermal stability is poor Thermal stability is good It has good selectivity due to its single active site Its selectivity ranges from good to poor due to the presence of multiple active sites. Role of the Catalyst in the Haber - Bosch process Iron catalyst is used in Haber – Bosch process to lower the activation energy needed for the breakdown of nitrogen and hydrogen (Frederic, Vandome & McBrewster 2010). N2 N2 + 3H2 2NH3 The use of iron catalyst lowers the activation energy and so the reaction becomes faster and economical. The pressure in the chamber is also raised to compress the gases together and hence increase the collisions and the reaction rate concurrently (Frederic, Vandome & McBrewster 2010). Bibliography Prigigine, I & Wildom, B. 2001. Collision Theory of Chemical Reactions. John Wiley and Sons. Fogler & Gurmen. 2007. Rate Law and Stoichimetry. University of Michigan. Boundless Chemistry. The Collision Theory. Available at. Frederic, P. M., Vandome, A. F. & McBrewster, J. 2010. The Haber process. VDM Publishing, 2010. Read More