Explaining the Factors That Affect the Rate of Reaction - Assignment Example

Task A. Explaining the factors that affect the rate of reaction Temperature Raising the temperature of a chemical reaction causes the reaction rate to rise due to the increased amount of high energy collisions between the molecules involved. It is only such collisions that end up in a reaction. One requirement for the collisions to occur is that they should have an energy content that is at least equal to the activation energy of the reaction. These collisions can be explained by the collision theory. The theory was put forward by Max Trautz in 1916 and then again in 1918 by William Lewis. It aims to explain how chemical reactions happen under various physical and chemical conditions. The theory states that when appropriate molecules of a reactant collide against each other, it is only a definite proportion of the collisions that result in a perceptible chemical change (Goldberger and Watson, 2004). These collisions are termed successful collisions, and possess activation energy. The idea of activation energy was introduced by Svante Arrhenius in 1889, and is the amount of energy needed to be gained by the reactant molecules to form the product. During the exact instant of collision, the pre-existing bonds are broken and new bonds formed. This results in the formation of the products of the reaction. The Maxwell-Boltzmann distribution law explains how the number of molecules that exist at a specific temperature is spread over a range of kinetic energies (Guggenheim, 2007): The higher the temperature the higher the rate at which its constituent particles move. This increased particular motion increases the chances of the particles colliding, resulting in more successful collisions that lead to bonds being broken and new ones forming, resulting in the formation of new products, and thereby increasing the reaction rate. This is proven in the graph above, which shows that the molecules that have higher kinetic energies resulting from high temperatures have higher reaction rates, since the increased motion between the molecules results in increased molecular collisions that raises the reaction rates. 2. Pressure Pressure has an effect on the rate of various chemical reactions, only involving gases. Raising it leads to a corresponding increase in the rate of the reaction. Raising or lowering the pressure of a reaction system involving liquids or solids leads to no change in the rate of the reaction. For a specific amount of gas, to raise the pressure of the gas one would have to compress the gas so that it would be contained in a lesser volume. Doing so would mean the same amount of gas is found in a much smaller volume, resulting in a higher concentration. Since the gas particles are closer to each other, their random motions result in more frequent collisions. These increased collision frequencies result in the breaking of bonds between the gas molecules, and new ones forming, resulting in an increased reaction rate. This can be illustrated as: 3. Reactant concentration The reactant concentration has an effect on the reaction rate of chemical reactions involving liquids (solutions) only. For gases, this is done by altering the pressure. For a reaction involving molecules of the same or different elements, the molecule have to collide, regardless of whether both molecules are in liquid form or if one is in solid form. Having a higher concentration of a reactant means that there is an increased chance of the molecules colliding, thereby increasing the chances of formation of new bonds (and products) (Burton, 2000). To hasten the formation of one product that is obtained from combining two liquids, one would have to increase the concentration of the limiting reagent. Increasing this reagent implies there is more of the reagent per unit volume that will react. This high volume reagent per unit volume means there are more successful collisions possible between the particles in the solution, increasing the reaction rate: The concentration of the reactants has no effect on the rate of chemical reactions involving catalysts, as the reaction will only proceed as fast as the catalyst permits. 4. Using a catalyst Catalysts are important in several chemical reactions. In a typical chemical reaction, the formation of products from the reactants follows the path that consumes the least amount of energy. In some cases, the activation energy required for this step could be high, resulting in the reaction to proceed slowly. The use of catalysts is explained by the transition state theory that states that there exists a transition state between the two states where reactant molecules and product molecules are found (Levine, 2005). In the transition state, the reactant molecules come together to form an activated complex. The concentration of this complex, the rate at which its bonds are broken, and the manner in which the bonds are broken will establish whether the reaction will occur. In the transition state theory, two pathways are possible: one involves the activated complex dissociating and the reactant molecules recombining, and the other involves the rearrangement of the bonds in the activated complex to yield the products. In the transition state, catalysts raise the stability of the activated complexes, thereby favouring the formation of the products, thus raising the reaction rate. The catalyst is not used up. Catalysts bind to the molecules of the reactant resulting in the reorganization of the electron densities in the reactant molecules. This leads to the weakening of the bonds that are supposed to be broken, thus hastening the reaction since the weakened bonds more easily broken. When the reactant molecule onto which the catalyst is bound reacts, the catalyst molecule ends up being released and is thus capable of binding to another reactant molecule and repeating this whole process. By introducing a catalyst into a slow reaction, an alternative path with a much lower activation energy is followed, hastening the reaction. The alternative path with the lower activation energy will quicken the reaction rates for reactions that would normally occur slowly. Task B. Catalysts and their roles in industries 1. Homogenous and heterogeneous catalysts A homogeneous catalysis is made up of a series of reactions involving the reactants and the catalyst being in the same phase (liquid, solid, gas), either dissolved in a solvent with the reactants, or as a gas. An example of a homogenous catalysis involves the oxidation of iodide ions (I-) by peroxodisulphate ions (S2O82-) using Fe2+ ions as the catalyst (Lister and Renshaw, 1999). Adding Fe2+ ions to the same solution containing the two ions yields: S2O82- (aq) + 2Fe2+ (aq) ------------> 2SO42- (aq) + 2Fe3+ (aq) The Fe3+ then oxidizes the iodide ions to form iodine, regenerating Fe2+: 2Fe3+ (aq) + 2I-(aq) -------------> 2Fe2+ (aq) + I2 (g) A heterogeneous catalysis employs the use of a catalyst in a phase that is different from that of the reactants. It mostly involves using a catalyst in solid form and the reactant in either liquid or gaseous form. In such a reaction, one or both reactants get adsorbed onto the catalyst’s surface, binding it at the active site. After binding, the interaction between the catalyst molecules and the reactant molecules leads to weakening of the bonds in the molecule followed by successful collisions between the reactant molecules that lead to the weakened bonds being broken and new products formed. The product molecules then undergo desorption, thus regenerating the catalyst. An example of a heterogeneous catalysis is the use of nickel in the oxidation of sulphur dioxide in oxygen to sulphur trioxide by passing the gaseous mixture over solid vanadium (V) oxide(Louie, 2005): SO2 (g) + V2O5 (s) -----------------> SO3 (g) + V2O4 (s) The vanadium (IV) oxide then undergoes oxidation by the oxygen, regenerating V2O­5: V2O4 (s) + O2 (g) ----------------> V2O5 (s) 2. Role of catalysts in the Haber-Bosch process Synthesis of ammonia by the Haber-Bosch process follows the equation: The Haber-Bosch yields of ammonia at various pressures and temperatures are shown below: From the graph above, it is evident that at 200°C and a pressure of 750 atmospheres, all of the reactants are converted into ammonia. To sustain such a high pressure in large scale is very difficult. Lowering the pressure to around 250-350 atmospheres and raising the temperature to between 400-500°C while passing the gases over four catalyst beds was found to overcome the problem with containing the gases at a very high pressure (Smil, 2004). Each pass over one catalyst bed leads to the conversion of around 15% of the mixture. The process makes use of an iron oxide catalyst contained in an ore called magnetite. Exposing the ore to hot hydrogen eliminates oxygen, thus reducing the iron. During the reduction, the catalyst retains a large proportion of its volume, resulting in a material that is very absorbent and with a large surface area. The increased surface area greatly increases the rate of the conversion (Leigh, 2007). The use of the iron catalyst provides an alternative pathway with a lower activation energy for the breaking of the bonds between N2 and ­H2. The N2 and H2 get adsorbed onto the catalyst’s surface, forming an iron-enzyme complex: The adsorbed N2 and H2 is combined at a ratio of 3H:1N and the combined mixture is adsorbed onto the catalyst and then cooled, turning it into liquid ammonia. The H2 for the Haber-Bosch process is obtained by adding steam to methane in the presence of a nickel catalyst at a temperature of 700°C: References Burton, G. (2000) The Effect of Concentration on Rate, in Salters Advanced Chemistry: Chemical Ideas, Heinemann Educational Publishers. Goldberger, M. and Watson, K. (2004) Collision Theory, Dover Publications. Guggenheim, E. (2007) Boltzmanns Distribution Law, North-Holland Pub. Co. Leigh, G. (2007) Catalysts for Nitrogen Fixation, Haber-Bosch and Other Industrial Processes, vol. 1, pp. 33-54. Levine, R. (2005) Structural Considerations in the Calculation of Reaction Rates, in Molecular Reaction Dynamics, Cambridge University Press. Lister, T. and Renshaw, J. (1999) Kinetics, in Understanding Chemistry for Advanced Level, Nelson Thornes. Louie, D. (2005) Contact Section, in Handbook of Sulphuric Acid Manufacturing, DKL Engineering, Inc. Miller, F., Vandome, A. and McBrewster, J. (2009) Ideal Gas Law: Equation of State, Ideal Gas, Gas, Boyles Law, Charless Law, Kinetic Theory, Rudolf Clausius, State Function, Pressure, Volume, Amount of Substance, Mole (unit), Gas Constant, Kelvin, Alphascript Publishing. Moritmer, M. and Taylor, P. (2002) Chemical Kinetics and Mechanism, Royal Society of Chemistry. Smil, V. (2004) Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production, The MIT Press. Read More