Reaction Rates and Equilibrium

1. Reaction Rates and Equilibrium

2. What is meant by the rate of a chemical reaction? • Can also be explained as the speed of he reaction, it is the amount of time required for a chemical reaction to come to completion. • Different reactions take different times • Burning • Aging • Ripening • Rusting

3. Collision Theory • Atoms, ions and molecules can react to form products when they collide provided that the particles are orientated correctly and have enough kinetic energy. • The relative orientations of the molecules during their collisions determine whether the atoms are suitably positioned to form new bonds. • Particles lacking necessary kinetic energy to react bounce apart when they collide. • Imagine clay balls… • The minimum amount of energy that particles must have in order to react is called the activation energy

4. Activated Complex

5. Activation Energy • Swedish Chemist Svante Arrhenius explored activation energy, Ea • He found that most reaction rate data obeyed an equation based on three factors: • The fraction of molecules possessing the necessary activation energy or greater • The number of collisions occurring per second • The fraction of collisions that have the appropriate orientation • Arrhenius equation k = Ae-Ea/RT

6. Factors that Affect Reaction Rates • Temperature (Alkaseltzer activity) • Raise temperature, faster reaction rate • Lower temperature, slower reaction rate • Higher temperatures make molecules move faster because they have more kinetic energy so reaction is more likely • Particle Size/Surface Area (lycopodium demo) • The smaller the particle size the larger the surface area. • An increase in surface area, increases the amount of reactant exposed, which increases the collision frequency. • Can be dangerous… coal powder or gas particles reacting to our lungs.

7. Catalyst (demo) • A substance that increases the rate of a reaction without being used up itself during the reaction. • They help reactions to proceed at a lower energy than is normally required. • Enzymes catalyze reactions in our body • An inhibitor interferes with the action of a catalyst

8. Concentration • The number of reacting particles in a given volume also affects the rate at which reactions occur. • Cramming more particles into a fixed volume increases the concentration of reactants, the collision frequency and therefore, reaction rate.

9. Rate Law: Effect of Concentration on Rate • One way of studying the effect of concentration on reaction rate is to determine the way in which the rate at the beginning of a reaction depends on the starting concentrations. a A + b B  c C + d D Rate = k[A]m[B]n • The exponents m and n are called reaction orders and depend on how the concentration of that reactant affects the rate of reaction.

10. N2(g) + 2 H2O(l) NH4+(aq) + NO2−(aq) Concentration and Rate If we compare Experiments 1 and 2, we see that when [NH4+] doubles, the initial rate doubles.

11. N2(g) + 2 H2O(l) NH4+(aq) + NO2−(aq) Concentration and Rate Likewise, when we compare Experiments 5 and 6, we see that when [NO2−] doubles, the initial rate doubles.

12. Concentration and Rate • This means Rate  [NH4+] Rate  [NO2−] Rate  [NH4+] [NO2−] which, when written as an equation, becomes Rate = k [NH4+] [NO2−] • This equation is called the rate law, and k is the rate constant. Therefore,

13. Rate Laws • A rate law shows the relationship between the reaction rate and the concentrations of reactants. • The exponents tell the order of the reaction with respect to each reactant. • Since the rate law is Rate = k [NH4+] [NO2−] the reaction is First-order in [NH4+] and First-order in [NO2−].

14. Rate Laws Rate = k [NH4+] [NO2−] • The overall reaction order can be found by adding the exponents on the reactants in the rate law. • This reaction is second-order overall.

15. Reversible Reactions • Until now, most of the reactions we have examined have gone completely to products. Some reactions are reversible meaning the reaction from reactants to products also goes from products to reactants at the same time. • Consider the following reaction: 2SO2 + O2 2SO3 In this reaction, SO2 and O2 are placed in a container. Initially, the forward reaction proceeds and SO3 is produced. The rate of the forward reaction is much greater than the rate of the reverse reaction. As SO3 builds up, it starts to decompose into SO2 and O2. The rate of the forward reaction is decreasing and the rate of the reverse reaction is increasing. Eventually, SO3 decomposes to SO2 and O2 as fast as SO2 and O2 combine to form SO3 (see Fig. 19.10 from text). When this happens, the reaction has achieved chemical equilibrium. • Chemical equilibrium is when the rate of the forward reaction = the rate of the reverse reaction. It says nothing about the amounts of reactants and products at equilibrium. Simulation

16. Equilibrium Constants • When a system reaches equilibrium there is a mathematical relationship between the concentrations of the products and the concentrations of the reactants. • aA + bB cC + dD • Kc = [C]c * [D]d [A]a * [B]b • Ch 19 pg 545-548: Practice problems 8-13

17. Factors Affecting Equilibrium: Le Chatelier’s Principle • If a stress is applied to a system at dynamic equilibrium, the system changes to relieve the stress. • Concentration • Temperature • Pressure

18. Concentration • Increasing the concentration of a substance causes the reaction to favor one reaction direction over the other to remove some of the extra substance. • Decreasing the concentration of a substance causes the reaction to shift to produce some of the substance that was removed. • For example: 2SO2 + O2 2SO3 • a) If [SO2] increases by adding more, what happens the concentrations of all of the substances? [SO2] goes up; [O2] goes down; [SO3] goes up. • Adding SO2 causes that to go up. The system tries to get rid of that extra SO2 so some of the additional will react with O2 causing the O2 to decrease. This speeding up of the forward reaction causes more SO3 to be produced. Adding SO2 causes the forward reaction to be favored to use up excess SO2.so reaction shifts to the right. • b) If [SO3] decreases by removing some, what happens the concentrations of all of the substances? [SO2] goes down; [O2] goes down; [SO3] goes down. • Removing SO3 causes the reaction to favor the forward reaction to try to produce more so reaction also shifts to the right.

19. Temperature • Increasing the temp. causes the equilibrium position to shift in the direction that absorbs heat (favors the endothermic reaction). Decreasing the temp. causes the equilibrium position to shift in the direction that produces heat (favors the exothermic reaction). • For example: 2SO2 + O2 2SO3 + heat • a) If temp increases it will cause the reaction to shift left (favor the reverse reaction) to try to remove the extra heat so: [SO2] goes up; [O2] goes up; [SO3] goes down.

20. Pressure • Affects systems containing gas particles since gasses are most affected by pressure. If pressure is increased, the reaction will shift to the side that contains the fewest number of gas particles and vice-versa (favor the reaction where most gas particles are reactants). • For example: 2SO2(g) + O2 (g) 2SO3(g) + heat • a) If pressure is increased, the reaction will shift to the side with the fewest gas particles so it will shift to the right and: [SO2] goes down; [O2] goes down; [SO3] goes up. • Chapter 19 pg 544: Practice problems: 6,7