Reaction Rates & Chemical Equilibrium PDF

Lecture II Dr. Omer 1  The rate of a reaction is the change in concentration of a reactant (or product) per unit time.  Some chemical reactions take place rapidly. Net ionic equation: Ag+(aq) + Cl‾(aq) ➙ AgCl(s)  Some chemical reactions take place very slow.  The st...

Lecture II Dr. Omer 1  The rate of a reaction is the change in concentration of a reactant (or product) per unit time.  Some chemical reactions take place rapidly. Net ionic equation: Ag+(aq) + Cl‾(aq) ➙ AgCl(s)  Some chemical reactions take place very slow.  The study of reaction rates is called chemical kinetics. EX:  The concentration might increase from 0 to 0.12 mol/L over a period of 30 min.  The rate of the reaction is the change in the concentration of iodomethane divided by the time interval: This unit is read “0.0040 mole per liter per minute.” During each minute of the reaction, an average of 0.0040 mol of chloromethane is converted to iodomethane for each liter of solution. 2  Collision Theory A B  Figure 2 The energy of molecular collisions varies. (a) Two fast-moving molecules colliding head-on have a higher collision energy than (b) two slower- moving molecules colliding at an angle. The collision theory indicates that a reaction takes place only when molecules collide with the proper (1)orientation and (2)sufficient energy The energy comes from the collision between A and B. If the energy of the collision is large enough, bonds will break and a reaction will take place. If the collision energy is too low, the molecules will bounce apart without reacting. GREATER THAN ACTIVATION ENERGY The minimum energy necessary for a reaction to occur 3  Collision Theory For this reaction to take place, the molecules must collide in such a way that the H of the HCl hits the O of the water, as shown in Figure. A collision in which the Cl hits the O. cannot lead to a reaction, even if sufficient energy is available. H2O(l) + HCl(g) ➙ H3O+(aq) + Cl‾(aq) The collision theory indicates that a reaction takes place only when molecules collide with the proper (1)orientation and (2)sufficient energy Even if two molecules collide with an energy greater than the activation energy, a reaction may not take place if the molecules are not oriented properly when they collide. Consider, for example, the reaction between H2O and HCl 4  Collision Theory [Summary] More example, N2+ O22NO. To form NO product, the collisions between N2 and O2 molecules must collide in a specific orientation and energy: 5  Activation Energy Activation energy: the minimum amount of energy required to break the bonds between atoms of the reactants.  The activation energy of a reaction must be supplied for the reaction to proceed.  Each reaction has its own activation energy. In an exothermic reaction (Eproduct Ereactant). < Might expect the reaction to take place rapidly. However, the reactants cannot be converted to products without the necessary activation energy. Six Covalent Six Covalent Bond break Bond form 6  The activation energy is inversely related to the rate of the reaction.  The lower the activation energy, the faster the reaction;  the higher the activation energy, the slower the reaction. When the reacting molecules reach this point, one or more original bonds are partially broken, and one or more new bonds may be in the process of formation. 7 Activation energy is analogous In the same way, a collision to climbing a hill. To reach a must provide enough destination on the other side, energy to push the you must have the energy reactants to the top of the needed to climb to the top of hill. the hill. Activation Energy Once at the top, its easy to run Then the reactants down the other side. The may be converted to energy needed to get us from products. the starting point to the top of the hill would be the activation energy. 8  Factors affecting reaction rate A. Nature of the Reactants:  Ionic comp in aq. Sol. Has faster reaction than covalent comp.(whether in aqueous solution or not): Because, activation energies for these reactions are very low because usually no covalent bonds must be broken. Many of Covalent comp reactions require 15 min to 24 h or longer B. Concentration:  The rate of a reaction increases when the concentration of the reactants increases: Because, increasing the number of effective collisions. That’s why chemicals in solid phase (as a powder) has faster rate of reaction Increase surface area.  For reactions in the gas phase, an increase in pressure usually increases the rate.  Example: A patient having difficulty breathing is given a breathing mixture with a higher oxygen content than the atmospheric concentration: Hb+ O2HbO2 9  Factors affecting reaction rate C. Temperature: Reaction rates increase with increasing temperature.  Because, at higher temperatures, the kinetic energy of the reactants increase, making them move faster and therefore collide more often and it provides higher % of collisions with more energy for activation. Reactions almost always go faster at higher temperatures. To cook food faster, we raise the temperature. When body temperature rises, the pulse rate, rate of breathing, and metabolic rate increase. We slow down a reaction by decreasing the temperature. We refrigerate food to slow molding and decay. In some heart surgeries, the body temperature is decreased to 25°C (82°F) so the heart can be stopped and less O2 is required by the brain. 10  Factors affecting reaction rate D. Catalyst: Any substance that increases the rate of a reaction without itself being used up. Some that increase the rate of only one reaction and others that can affect several reactions Heterogeneous catalyst: A catalyst in a separate phase from the reactants. EX: Homogeneous catalyst: A catalyst in the same phase as the reactants. EX: Enzymes in body tissues.11  Chemical Equilibrium In earlier chapters, we considered the forward reaction in an equation and assumed that all of the reactants were converted to products. However, most of the time, reactants are not completely converted to products because a reverse reaction takes place in which products collide to form the reactants. When a reaction proceeds in both a forward and reverse direction, it is called reversible. That means there are 2 reaction rates:  the rate of the forward reaction  the rate of the reverse reaction The rate of the forward reaction gradually decreases, while the rate of the reverse reaction (which began at zero) gradually increases. Eventually the two rates become equal. At this point, the process is in dynamic equilibrium (or just equilibrium). 12 Chemical Equilibrium - Example 13 13  Equilibrium Constants At equilibrium, the concentrations can be used to set up a relationship between the reactants and the products. Equilibrium constant expression for a reversible chemical reaction: The equilibrium constant is different for every reaction. Some reactions have a large K ; others have a small K. 14 14 The equilibrium constant for a given reaction remains the same no matter what happens to the concentrations, but it is not true for changes in temperature. The value of K changes when the temperature changes 15  Equilibrium with a Large Kc  When a reaction has a large Kc, it means that the forward reaction produced a large amount of products when equilibrium was reached.  The equilibrium mixtures contains more products than reactants. 16  Equilibrium with a Small Kc  When a reaction has a small Kc, it means that the forward reaction produced a small amount of products when equilibrium was reached.  The equilibrium mixtures contains less products than reactants. 17  Changing Equilibrium Conditions: Le Châtelier's Principle  When we alter any of the conditions of a system at equilibrium, the rates of the forward and reverse reactions will no longer be equal. This is called that stress is placed on the equilibrium.  Then the system responds by changing the rate of the forward or reverse reaction in the direction that relieves the stress to reestablish equilibrium. Le Châtelier's Principle: When a stress (change in conditions) is placed on a reaction at equilibrium, the equilibrium will shift in the direction that relieves the stress. 18 19  Changing Concentrations  When reactant is added, the concentration of reactants increases and the forward rate will increase to make more products to balance.  When reactant is removed, the concentration of reactants decreases and the reverse rate will increase to replace the lost reactant.  When product is added, the concentration of products increases and the reverse rate will increase to make more reactants to balance.  When product is removed, the concentration of products decreases and the forward rate will increase to replace the lost reactant. 20 Changing Temperature Endothermic reactions  Increase temperature shifts right to remove the heat  Decrease temperature shifts left to remove the heat Exothermic reactions  Increase temperature shifts left to remove the heat  Decrease temperature shifts right to remove the heat 21 Changing Pressure A change in pressure influences the equilibrium only if one or more components of the reaction mixture are gases Increasing pressure:  All the concentrations will increase.  The system will shift toward producing the side with fewer moles. Decreasing pressure:  All the concentrations will decrease.  The system is free to move toward producing the side with more moles. In the above reaction, Increasing pressure, the equilibrium will shift to the left. And vice versa 22 Effect of a Catalyst on Equilibrium  When a catalyst is added to a reaction, the activation energy is lowered.  Therefore speeding up the reaction. As a result, both the forward and reverse reactions increase, but the same ratios of products and reactants are attained. Therefore, catalysts do not effect equilibrium. 23 24 25  THANKS! Do you have any questions? Dr. Omer Saeed 733152363 26 27