Rates and equilibrium Rates of reaction, Slides of Analytical Chemistry

Download Rates and equilibrium Rates of reaction and more Slides Analytical Chemistry in PDF only on Docsity! Rates and equilibrium Rates of reaction - collision theory Introduction:  The term rate is used to describe the speed of a reaction.  Different reaction will have different rates Rate of reaction: The rate of a reaction is defined as the change in concentration of a reactant or product in a given time Rate = Change in concentration Units: mole dm-3 = mol dm-3 s-1 Time s  When a reaction starts the concentrations of the reactants are at their highest, the rate is at its fastest.  As the reactants are used up the concentration decreases, so does the rate.  When one of the reactants is used up the reaction stops. The rate is 0/ Factors affecting the rate of a chemical reaction: 1. Temperature – An increase in temperature will increase the rate of reaction. 2. Pressure - Higher pressure increases the rate of a reaction 3. Concentration – An increase in concentration (or pressure) will increase the rate of reaction. 4. Surface area – An increase in surface are (which is a decrease in particle size) will increase the rate of reaction. 5. Catalyst – A suitable catalyst will increase the rate of a reaction. Collision theory:  This is used to explain how reacting molecules collide leading to a reaction  When 2 molecules collide a reaction only occurs if the conditions are right. 1) Activation energy.  Reacting molecules have to collide with enough energy to break the initial bonds, the activation energy.  The collision must be greater than or equal to the activation energy.  If the collision is less than the activation energy, they just bounce apart. 2) Orientation:  The molecules must collide in the correct orientation in order for the products to be made: The effect of concentration on reaction rate: Increase in concentration Increase in No particle per volume Increase in the collision frequency Increases the rate of the reaction  This means that in a certain amount of time, more collision will take place.  This means there will be more collisions with energy greater than or equal to the activation energy.  Which means the rate increases. The effect of pressure on reaction rate: Increase in pressure Decrease in the volume Increase in the No particle per volume Increase in the collision frequency Increases the rate of the reaction 2) Heterogeneous catalyst  The catalyst and reactants are in different phases (ie gas reactants, solid catalyst) Examples Cracking, isomerisation and reforming reactions in the oil industry. Haber process – Production of ammonia  Ammonia used to be produced by using an electrical discharge tube with nitrogen.  Ammonia is used in the production of fertillisers  Fritz Haber used a gas phase reaction between nitrogen and hydrogen: N2(g) + H2(g)  NH3(g)  The process uses a finely divided iron catalyst.  The iron catalyst weakly absorbs the nitrogen molecules weakening the very strong triple bonds. Questions 1 - 4 P 205 / Questions 1 - 4 P 207 The Boltzmann distribution  In a gas or liquid it is assumed that the molecules collide with elastic collisions.  This means that no energy is lost during a collision with the container or other molecules.  As the molecules are moving, they have kinetic energy.  In any sample of gas or liquid there will be a distribution of energies: Some will have high energy: This means they will move very fast Some will have low energy This means they will move very slow Most will have average energy Most will move with the average speed Important features of the Boltzmann distribution: The area under the curve is equal to the total number of molecules in the sample: The area does not change No molecules have zero energy: The curve starts at 0,0 There is no maximum energy for a molecule: The curve never touches the axis Only molecules with energy greater than the activation energy are able to react The activation energy is the minimum energy required to react The effect of temperature on reaction rate: Increase in temperature Increase in kinetic energy of molecules Increase in the speed of the molecules - Increases the collision frequency - Increases the No of collisions with energy > activation energy More successful collisions Increase in the rate The effect of a catalyst on reaction rate: Catalysts offer a route with a lower activation energy As more molecules will overcome the lower activation energy An increases in the number of successful collisions Increases the rate of the reaction Questions 1-3 P209 Chemical equilibrium Reversible reactions:  Consider the reaction: Mg(s) + H2SO4(aq)  MgSO4(aq) + H2(g)  The reaction stops when all of the limiting reagent has been used up.  The reaction is said to go to completion and this is indicated by   Some reactions are reversible though: 2SO2(g) + O2(g)  2SO3(g)  This reaction is reversible and can take place in the forward direction (above) or the reverse direction (below): 2SO3(g)  2SO2(g) + O2(g)  The reaction is described as reversible. Equilibrium is indicated by  Dynamic equilibrium:  When a reversible reaction has appeared to have stopped we say it has reached an equilibrium.  At equilibrium the rate of the forward and reverse reaction occurs at the same rate.  This means that the concentrations of the reactants and products are the same.  There is no observable change.  This is described as DYNAMIC.  The extent of how far the reaction has gone towards the products is called 'the position of the equilibrium'  Equilibrium is stable under fixed conditions and in an isolated system: Fixed conditions: Temperature and pressure remains the same. Isolated system: No material is added or removed from the system Factors affecting the position of equilibrium:  The position of an equilibrium can be altered by altering the conditions or adding / removing material from the system. 1) Changing the concentration of reactants or products 2) Changing the pressure (if gases involved) 3) Changing temperature  The effects of these changes can be predicted using: 3) Changing temperature: a) Ammonium chloride  The white solid vaporises on heating into NH3(g) and HCl(g).  Further up the tube when the vapour has cooled NH4Cl(s) reforms.  This is an example of a reversible reaction: NH4Cl(s)  NH3(g) + HCl(g) On heating NH3(g) + HCl(g)  NH4Cl(s) On cooling A change in temperature affects the position of the equilibrium NH4Cl(s)  NH3(g) + HCl(g) H = endothermic  Heating the reaction moves it to the products. The cold side – endothermic.  Cooling moves the reaction to the reactants. The hot side – exothermic. The equilibrium moves to oppose the change imposed upon it b) Nitrogen oxides:  The syringe contains 2 gases in equilibria: N2O4  2NO2 H = +58KJ Mol-1 Colourless Brown Endothermic Hot Cold  The hot syringe had a very dark brown colour.  Heat moves the reaction moves towards the products.  An increase in temperature increases the amount of NO2. The equilibrium moved in the forward direction (towards the cold side).  The cool syringe had a very pale brown colour.  Cooling moves the reaction moves towards the reactants.  A decrease in temperature increases the amount of N2O4. The equilibrium moved in the reverse direction (towards the hot side). A change in temperature affects the position of the equilibrium The equilibrium moves to oppose the change imposed upon it Hot Cold Increasing Temperature Temperature Summary:  If you increase the temperature, the equilibrium will move to the side that will reduce temperature.  If you decrease the temperature, the equilibrium will move to the side that will increase temperature. A change in concentration affects the position of the equilibrium The equilibrium moves to oppose the change imposed upon it The effect of a catalyst on an equilibrium system  A catalyst has no effect on the position of the equilibrium.  A catalyst speeds up the rate of a reaction so it will only increase the rate at which equilibrium is achieved.  This is true for the forward and reverse reaction. Equilibrium and industry: The Haber process  Ammonia is used as it is soluble and is readily oxidised into nitric acid by the ’Ostwald process’.  The reaction: N2(g) + 3H2(g)  2NH3(g) H = -92KjMol-1 Endothermic Exothermic 4 Moles 2 Moles high pressure Low pressure The problem: 1. The equilibrium lies over to the reactants. 2. Nitrogen has a triple bond which makes it very hard to break. This gives the reaction a high activation energy The solutions: Process Equilibria Rate Compromise Temperature Decrease temperature – will move equilibria to the exothermic side – the products Increase temperature – will increase the rate of reaction as activation energy is more likely to be overcome. Moderate temperature of 400 - 500 oC used. Pressure Increase pressure – Moves equilibrium to products with fewest moles of gas. Increase pressure – will increase the rate of reaction as more particles per volume. Cost of pumps and reaction vessel becomes very expensive – 200 atm Catalyst No effect Increases the rate Finely divided iron with metal oxide promoters Remove ammonia as it is formed Equilibrium is never reached so the rate doesn’t slow down. Recycle unreacted H2 and N2  The plants run continuously with a lower energy and less expensive to build than conventional plants.  80% goes to making fertillisers such as ammonium sulphate.  A proportion makes nitric acid which goes on to make explosives. The process:  Other industrial processes include the contact process - converting sulphur dioxide to sulphur trioxide. Questions 1-2 P211 / 1-2 P213 / 8-13 P215 / 2 P217 High pressure reaction vessel with iron catalyst. Expansion – mixture cools and ammonia liquefies. Unreacted H2 and N2 is recycled N2 from fractional distillation of air H2 from steam + hydrocarbon Ammonia