Presentation on theme: "Reaction Rates & Equilibrium Unit 12 - Chapter 18."— Presentation transcript:
Reaction Rates & Equilibrium Unit 12 - Chapter 18
Reaction Rate Reaction rate – how fast reactants disappear and how fast product appears
A B 13.1 rate = - [A] tt rate = [B][B] tt time
Reaction Rate Reaction Rate = ∆ [A] ∆ t Example: CO (g) + NO 2(g) CO 2(g) + NO (g) - at t = 4.0 min, [CO 2 ] =.12 mol/L - at t = 8.0 min, [CO 2 ] =.24 mol/L - reaction rate =.24 mol/L -.12 mol/L 8.0 min – 4.0 min = 0.030 mol/L. min Unit for reaction rate = conc. with some time unit Products have a (+) rate Reactants have a (-) rate
Collision Theory of Kinetics Kinetics is the study of the factors that affect the speed of a reaction and the mechanism by which a reaction proceeds. In order for a reaction to take place, the reacting molecules must collide into each other. Once molecules collide they may react together or they may not, depending on two factors - ¬Whether the collision has enough energy to "break the bonds holding reactant molecules together"; Whether the reacting molecules collide in the proper orientation for new bonds to form. orientation
Effective Collisions Collisions in which these two conditions are met (and therefore the reaction occurs) are called effective collisions. The higher the frequency of effective collisions the faster the reaction rate. When two molecules have an effective collision, a temporary, high energy (unstable) chemical species is formed - called an activated complex It is a transition state between reactant and product It has a very short lifetime (10 -13 s) Has to form for product to be formed
Activated Complex The difference in potential energy between the reactant molecules and the activated complex is called the activation energy, E a This is the minimum amount of energy that particles must have in order to react. The larger the activation energy, the slower the reaction The energy to overcome the activation energy comes from the kinetic energy of the collision being converted into potential energy, or from energy available in the environment, i.e. heat. Different reactions have different activated complexes and therefore different activation energies
Energy Diagram Energy of products is lower than energy of reactants energy lost, exothermic, -∆H Energy of products is higher than energy of reactants energy gained, endothermic, +∆H What is this called?
Factors Affecting Reaction Rate 1.Nature of the Reactants Cl 2(g) + CH 4(g) CH 3 Cl (g) + HCl (g) Cl 2 Cl + Cl(fast) Cl + CH 4 CH 3 Cl + H(slow) H + Cl HCl(very fast) individual steps = elementary steps all steps together = reaction mechanism the slowest step determines the rate of the reaction called the rate determining step Intermediates – product in one step, reactant in another
Factors Affecting Reaction Rate 2.Concentration The larger the concentration of reactant molecules, the faster the reaction will go. Increases the frequency of reactant molecule collisions 3. Particle Size (Surface Area) more particles on the surface = more particles available for collisions more collisions = more act. complex = more product smaller particles give you more surface area
Factors Affecting Reaction Rate 4. Agitation this puts more liquid/gas particles in contact with the solid = ↑ collisions = ↑ act. complex = ↑ product 5. Pressure ↑ pressure by ↓ volume – puts particles closer together = ↑ collisions = ↑ act. complex = ↑ product All of these factors are similar, in terms of explanation, to concentration!!!
Factors Affecting Reaction Rate 6. Temperature most effective at speeding up a reaction ↑ temp. = ↑ KE (particles moving faster) particles move faster leading to more collisions the collisions are also harder these harder collisions contain the needed energy to overcome the E a therefore the reaction rate will increase
Factors Affecting Reaction Rate 7. Catalyst substance that speeds up a reaction, but isn’t used up in the reaction provides a “different pathway” that requires lower E a lower E a = more collisions having the proper amount of energy = ↑ act. complex = ↑ product
Reaction Dynamics If the products of a reaction are removed from the system as they are made, then a chemical reaction will proceed until the limiting reactants are used up. However, if the products are allowed to accumulate; they will start reacting together to form the original reactants - called the reverse reaction. We show this reverse reaction by using a double arrow (H 2(g) + I 2(g) 2HI (g) )
Reaction Dynamics The forward reaction slows down as the amounts of reactants decreases because the reactant concentrations are decreasing At the same time the reverse reaction speeds up as the concentration of the products increases. Eventually the forward reaction is using reactants and making products as fast as the reverse reaction is using products and making reactants. This is called chemical equilibrium. rate forward = rate reverse Note: This equilibrium is dynamic
Chemical Equilibrium Dynamic Equilibrium can only be reached in a closed system!! When a system reaches equilibrium, the amounts of reactants and products in the system stays constant the forward and reverse reactions still continue, but because they go at the same rate the amounts of materials don't change. There is a mathematical relationship between the amounts of reactants and products at equilibrium
Equilibrium Expression Capital letters (A,B,C,D) – reactants or products Lowercase letter (a,b,c,d) – coefficients from the equation NOTE – products on top, reactants on bottom In this expression, K eq is a number called the equilibrium constant. ratio of product concentration to reactant concentration at equilibrium Do not include solids or liquids, only solutions and gases The value of K eq depends on temp. of the reaction – if temp. changes then the value of K eq changes. aA + bB cC + dD = [C] c [D] d [A] a [B] b K eq
Example – Determine the value of the Equilibrium Constant for the Reaction 2 SO 2(g) + O 2(g) 2 SO 3(g) ¬Determine the Equilibrium Expression Plug the equilibrium concentrations into to Equilibrium Expression ®Solve the Equation 3.503.00SO 3 1.251.50O2O2 2.00SO 2 [Equilibrium][Initial]Chemical So what does this K eq value tell us???
Position of Equilibrium The size of the equilibrium constant shows whether products or reactants are favored at equilibrium. K eq > 1, products are favored at equilibrium K eq < 1, reactants are favored at equilibrium
¬Determine the Equilibrium Expression Plug the equilibrium concentrations and Equilibrium Constant into the Equilibrium Expression ®Solve the Equation ?3.00SO 3 1.251.50O2O2 2.00SO 2 [Equilibrium][Initial]Chemical Example – If the value of the Equilibrium Constant for the Reaction 2 SO 2 + O 2 2 SO 3 is 4.36, Determine the Equilibrium Concentration of SO 3
More Equilibrium Practice 1.Write the equilibrium constant expression for the following reaction. 3H 2(g) + N 2(g) ↔ 2NH 3(g) 2.An analysis of an equilibrium mixture for this reaction in a 1.0 L flask at 300 o C gave the following results: 0.15 mol H 2, 0.25 mol N 2 and 0.10 mol NH 3. Calculate the K eq for this reaction. 3.2BrCl (g) ↔ Cl 2(g) + Br 2(g) The equilibrium constant for this reaction is 11.1. The equilibrium mixture contains 4.00 mol Cl 2 and 4.00 moles of Br 2. How many moles of BrCl are present?
Le Ch âtelier’s Principle Le Châtelier's Principle guides us in predicting the effect various changes have on the position of equilibrium it says that if stress is applied to a system in dynamic equilibrium, the system will change to relieve the stress. The position of equilibrium moves to counteract the change. Three common stressors: Concentration Temperature Pressure
Concentration Changes and Le Châtelier’s Principle A+ B ↔ C + D Adding a reactant – equilibrium shifts right Removing a reactant – equilibrium shifts left Adding a product – equilibrium shifts left Removing a product – equilibrium shifts right
Only affects a reaction involving gases with an unequal number of mole of reactants & products. Increasing the pressure on the system causes the position of equilibrium to shift toward the side of the reaction with the fewer gas molecules Decreasing pressure causes a shift toward the side with more gas molecules Example 3H 2(g) + N 2(g) 2NH 3(g) + 92kJ ↑ Pressure – shifts to the right ↓ Pressure – shift to the left Changing Pressure and Le Châtelier’s Principle
Increasing the temperature causes the reaction to shift away from the heat. For exothermic reactions - Think of heat as a product of the reaction Therefore shift the position of equilibrium toward the reactant side For endothermic reactions - Think of heat as a reactant The position of equilibrium will shift toward the products Cooling an exothermic or endothermic reaction will have the opposite effects. Changing Temperature and Le Châtelier’s Principle
Examples – Le Chatelier’s Principle What effect do the following changes have on the equilibrium position for the following reaction? 1.PCl 5(g) + heat ↔ PCl 3(g) + Cl 2(g) a. addition of Cl 2 b. increase in pressure c. removal of heat d. removal of PCl 3 as formed 2. C (s) + H 2 O (g) + heat ↔ CO (g) + H 2(g) a. Lowering the temperature b. Increasing the pressure c. Removal of H 2 as formed
Examples – Le Chatelier’s Principle 3. At 425 K – Fe 3 O 4(s) + 4H 2(g) ↔ 3Fe (s) + 4H 2 O (g) How would the equilibrium concentration of H 2 O be affected by the following: a. Adding more H 2 b. Adding more Fe (s) c. Decreasing the pressure d. Adding a catalyst