Unit 9: Kinetics and Equilibrium

1. Unit 9: Kinetics and Equilibrium Chapter 12 General Chemistry I Edmond North High School

2. Collision Theory • The following three statements summarize the collision theory. • 1. Particles must collide in order to react. • 2. The particles must collide with the correct orientation. • 3. The particles must collide with enough energy to form an activated complex, which is an intermediate particle made up of the joined reactants. • An effective collision is one that results in a reaction; new products are formed.

3. The Collision Model • The Collision Model is used to explain current known characteristics of reaction rates. • According to the Collision Model: • Molecules must collide to react. • Concentration affects rates because collisions are more likely. • Must collide with enough energy. • Temperature and rate are related. • Only a small number of collisions produce reactions.

4. Activated Complex • The Activated Complex is an intermediate particle formed when reactants collide and stick together. • Old bonds are breaking while new bonds are forming. • The minimum amount of energy colliding particles must have in order to form an activated complex is called the activation energy. • Particles that collide with less than the activation energy cannot form an activated complex.

5. Reaction Coordinate Diagrams • It is helpful to visualize energy changes throughout a process on a reaction coordinate diagram. • It shows the energy of the reactants and products (and, therefore, H). • The high point on the diagram is the transition state. • The species present at the transition state is called the activated complex.

6. Exothermic Reactions Review • An exothermic reaction releases heat, and an endothermic reaction absorbs heat.

7. Endothermic Reactions Review • The endothermic reaction absorbs heat because the products are at a higher energy level than the reactants.

8. Orientations

9. Kinetics • Kinetics is the study of reaction rates and the factors that effect them. • Why is it important? • When we understand reaction rates we can control chemical reactions and use them for specific purposes.

10. Kinetics • As a reaction occurs: • concentration of reactants decreases • concentration of products increases • Kinetics studies the rate at which a chemical process occurs. • Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

11. Reaction Rate • Reaction Rate is the number of atoms, ions or molecules that react in a given time to form products • We measure either the rate of disappearance of the reactants or the rate of appearance of one of the products • In simpler terms: • The average rate is the change in a given quantity during a specific period of time.

12. Measuring Reaction Rates • Change in Electrical Conductivity • Change in Color • Change in Pressure • Change in Volume

13. Reaction Rate Graph • For the reaction H2 + I2 → 2HI • What is happening to the concentration of the reactants? • What is happening to the concentration of the products?

14. Reaction Rates & Stoichiometry • We use coefficients to give us a ratio of rates between species in a reaction. • In the reaction 3H2 + N2 2NH3, since N2 has a coefficient of 1 and NH3 has a coefficient of 2, the rate N2 is used = 2(rate NH3 is created) • Likewise, since H2 has a coefficient of 3 and N2 has a coefficient of one, it is used 3 times as fast as N2.

15. Factors Affecting Reaction Rates • The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction. • An important factor that affects the rate of a chemical reaction is the reactive nature of the reactants. As you know, some substances react more readily than others. • The more reactive a substance is, the faster the reaction rate.

16. Factors That Affect Reaction Rates • Physical State of the Reactants • In order to react, molecules must come in contact with each other. • The more homogeneous the mixture of reactants, the faster the molecules can react. • Concentration of Reactants • As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

17. Factors that Effect Reaction Rate • Particle Size (surface area) • For solids, breaking up big pieces increases surface area, increasing rate by having more places for the molecules to interact • Ionic compounds have more surface area than covalent

18. Factors that Effect Reaction Rate • Pressure – Gases only! • As pressure increases the concentration increases, so you will have more collisions

19. 11-2 Factors That Affect Reaction Rates • Temperature • At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy. • Presence of a Catalyst • Catalysts speed up reactions by changing the mechanism of the reaction. • Catalysts are not consumed during the course of the reaction.

20. Catalysts • Catalysts increase the rate of a reaction by decreasing the activation energy of the reaction. • Catalysts change the mechanism by which the process occurs. • One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break.

21. Enzymes • Enzymes are catalysts in biological systems. • The substrate fits into the active site of the enzyme much like a key fits into a lock.

22. Factors that Effect Reaction Rate • Inhibitors • Compounds added to a reaction that slow it down • Examples: Lead in diesel, preservatives in food.

23. Reaction Rate Laws • Rate Laws relate the reaction rate and reactant concentration • We write the rate laws for elementary reactions from the equation. • They include only the gaseous or aqueous reactants! • Solids & liquids are not included because they have a constant concentration that does not increase or decrease

24. Rate Law • A Rate Law is an expression that shows how the rate depends on concentration of reactants • Example: • Rate laws use reactants only • k is proportionality constant (called the rate constant) • n is the order (exponent) of the reactant determined experimentally • n is generally the coefficient of the chemical in the balanced equation HOWEVER experimental data overrides the coefficient. • Reaction is xth order in A • Reaction is yth order in B • Reaction is (x +y)th order overall aA + bB cC + dD Rate = k [A]x[B]y

25. Order of Reaction • Order of Reaction • For the reaction with the rate = k[A]2[B] • Rate is 2nd order with respect to A • Rate is 1st order with respect to B • Rate is 3rd order overall (add exponents) • How does concentration affect reaction rate? • If [A] doubles what happens to the rate? • If [B] triples, what happens to the rate?

26. Temperature and Rate • Generally, as temperature increases, so does the reaction rate. • This is because k is temperature dependent.

27. Indicators of Completed Reactions • So how do you know when a reaction has gone to completion? • A gas is produced and escapes (open container) • A precipitate forms from two aqueous solutions. • A covalent product is formed (usually water)

28. What is Equilibrium? • When a reaction results in almost complete conversion of reactants to products the reaction goes to completion. • Most reactions, however, do not go to completion. They appear to stop. • The reason is that these reactions are reversible. • A reversible reaction is one that can occur in both the forward and the reverse directions.

29. Chemical Equilibrium • Chemical equilibrium occurs when two opposite reactions occurring at the same time and rate A + B ↔ AB • Forward & Reverse reactions don’t stop at equilibrium, just looks that way because concentration remains constant • Rateforward reaction = Ratereverse reaction

30. Equilibrium Example • You have a bridge between 2 cities. The number of cars going in either direction on the bridge are equal (at equilibrium) • The populations of the cities on either side of the bridge do not have to be equal!

31. Equilibrium Position • The equilibrium position of a reaction is determined by: • Initial concentrations • Energy of reactants and products • Degree of organization of reactants and products

32. Progression of Equilibrium • As a reaction progresses: • [A] decreases to a constant, • [B] increases from zero to a constant. • When [A] and [B] are constant, equilibrium is achieved. • In a system at equilibrium, both the forward and reverse reactions are being carried out; as a result, we write its equation with a double arrow.

33. A System at Equilibrium • Once equilibrium is achieved, the amount of each reactant and product remains constant. • Rates are equal; concentrations are not. • The concentrations do not change at equilibrium.

34. Equilibrium Positions • If equilibrium lies “to the left”: • There are more reactants and less products. • If equilibrium lies “to the right”: • There are less reactants and more products. • If reactants are mixed and concentrations do not change: • The reaction could already be at equilibrium. • Reaction rates are so slow that change is too difficult to detect.

35. Equilibrium Expressions • Equilibrium expressions relate concentrations of reactants to those of the products • These can be written from balanced equations • They look similar to rate laws! BE CAREFUL! • Uses both reactants & products • Include only the gaseous or aqueous phases as with rate laws

36. [NO2]2 [N2O4] Keq = kf kr = Deriving an Equilibrium Expression • Forward reaction: N2O4 (g)  2NO2 (g) • Rate law: Rate = kf [N2O4] • Reverse reaction: 2NO2 (g)  N2O4 (g) • Rate law: Rate = kr [NO2]2 • Therefore, at equilibrium: Ratef= Rater • kf [N2O4] = kr [NO2]2 • Rewriting this, it becomes

37. General Equilibrium Expressions • A (s) + 2B (g) ↔ 2C (g) + 3D (g) Keq = [C]2[D]3 [B]2 • Keq = equilibrium constant (capital K) • Numerical value of the ratio of product concentrations to reactant concentrations • [ ] = concentration in M (mol/L) • Order = coefficient becomes exponent

38. PbCl2 (s) Pb2+ (aq) + 2 Cl−(aq) Solids and Liquids Are Constant • Both can be obtained by dividing the density of the substance by its molar mass—and both of these are constants at constant temperature. • Therefore, the concentrations of solids and liquids do not appear in the equilibrium expression. Kc = [Pb2+] [Cl−]2

39. What Are the Equilibrium Expressions for These Equilibria?

40. Equilibrium Constant, K • If we know the value of K, we can predict: • The tendency of a reaction to occur. • If a set of concentrations could be at equilibrium. • The equilibrium position, given initial concentrations. • K will always have the same value at a certain temperature. • No matter what amounts are added, the ratio at equilibrium will always be same

41. Equilibrium Constant (Keq) • Tells you whether the products or reactants are favored • Keq > 1 • Products are favored (forward rxn); lots of product is made • Keq< 1 • Reactants are favored (reverse rxn); not much product made • The size of K and time needed to reach equilibrium are NOT related

42. The Units for K • Are determined by the various powers and units of concentrations. • They depend on the reaction. • At any temperature. • Temperature affects rate. • Equilibrium position is a set of concentrations at equilibrium.

43. The Reaction Quotient (Q) • To calculate the reaction quotient or Q, one substitutes the initial concentrations on reactants and products into the equilibrium expression. • Q gives the same ratio the equilibrium expression gives, but for a system that is not at equilibrium. • It is used to tell if a reaction is at equilibrium or not.

44. Reaction Quotient: Comparing Q and K • The relationship between Q and K tells which way the reaction will shift • Q = K: at equilibrium, no shift • Q > K: too large, forms reactants, shift to left • Q < K: too small, forms products, shift to right

45. Le Chatelier’s Principle • Le Chatelier’s Principle states that if stress is added to a system at equilibrium, the reaction will speed up in the direction that will relieve the stress. • Once the stress is relieved, equilibrium is re-established and no further changes are noticed. • 4 Types of stress • Concentration • Temperature • Pressure • Volume

46. Stress Factors 1. Concentration • Shifts away from an increase or addition • Shifts toward a decrease or subtraction

47. Stress Factors 2. Temperature (treat heat as a reactant if endothermic or product if exothermic) • Shifts away from an addition • Shifts toward a subtraction

48. Change in Temperature • Temperature affects the rates of both the forward and reverse reactions. • Doesn’t just change the equilibrium position, changes the equilibrium constant.

49. Stress Factors 3. Pressure (only effects gases) • Increase pressure, shifts to side with lower total moles of gas • Decrease pressure, shift to side with higher total moles of gas CO (g) + H2(g) ↔ CH4(g) + H20 (g)

50. Stress Factors 4. Changes in the volume • Suppose the volume of the reaction vessel for the system is decreased, resulting in an increase in pressure. • The equilibrium will shift to relieve the stress of increased pressure.