Reaction Rates and Equilibrium Pre AP

1. Reaction Rates and EquilibriumPre AP Chapter 19 C.Smith

2. I. Rates of ReactionA. Collision Theory • 1. Rate is defined as the measure of speed change that occurs over a period of time. • 2. In chemistry, rates of chemical change usually are expressed as the amount of reactant changing per unit time. • 3. Visible changes caused by chemical reactions are related to changes in the properties of individual atoms, ions, and molecules through a model called collision theory. • 4. Atoms, ions, and molecules can react to form products when they collide provided that the particles have enough kinetic theory.

3. I. Rates of ReactionA. Collision Theory • 5. The minimum amount of energy that particles must have in order to react is called activation energy. • 6. Activation energy is a barrier that reactants must cross to be converted into products • a. During a reaction, unstable particles that are neither products or reactants are temporarily formed. • b. The particle is called an activated complex and exist for a very short time. (approximately 10-13 sec.) • c. This period is called the transition state in where the activated complex can either reform reactants or form products.

4. I. Rates of ReactionA. Collision Theory 7. Collision theory explains why some naturally occurring reactions are immeasurably slow at room temperature and essentially at a rate of zero. 8. If the activation energy of the products is lower than the energy of the reactants, the reaction is exothermic. 9. If the activation energy of the products is higher than the energy of the reactants, the reaction is endothermic.

5. I. Rates of ReactionA. Collision Theory Addison-Wesley Chemistry p.535

6. I. Rates of ReactionB. Factors Affecting Reaction Rates • 1. Temperature a. At higher temperatures, molecules move faster and more chaotic and has an increase in kinetic energy. b. The effect increases the number of collisions between molecules and the more colliding particles are energetic enough to slip over the activation energy barrier.

7. I. Rates of ReactionB. Factors Affecting Reaction Rates • 2. Concentration a. Increasing the number of particles increase the number of collisions between particles. b. Increasing collisions between particles increases the frequency of particles that become activated complexes and possible products. c. Cramming more particles into a fixed volume increases the concentration of reactants, the collision frequency, and therefore the reaction rate.

8. I. Rates of ReactionB. Factors Affecting Reaction Rates • 3. Particle Size a. The smaller the particles size, the larger is the surface area for given mass of particles. b. An increase in surface area increases the amount of the reactant exposed for reaction, which further increases the collision frequency and the reaction rate. c. One way to increase the surface area of solid reactants is to dissolve them. d Another way to increase the surface area is to grind solids into fine powder. (explosion hazard)

9. I. Rates of ReactionB. Factors Affecting Reaction Rates • 4. Catalysts a. A catalyst is a substance that increases the rate of a reaction without being involved in or used up in the reaction. b. Catalysts lower the activation energy of a reaction which allows more particles to form products. • 5. Inhibitors a. An inhibitor interferes with the action of a catalyst. b. An inhibitor “poisons” a catalyst and can slow and even stop a reaction.

10. Question? • What does a catalyst do when a reaction is at EQUILIBRIUM? • A catalyst increase the reaction rate. At equilibrium the increases the rate is equal in both the forward and reverse directions equally. 

11. Potential Energy Diagrams • Activation Energy • "Ea"-The Energy required to initiate a chemical reaction. Both endothermic and exothermic reactions require activation energy.

12. Example- Factors Affecting Reaction Rates • How do catalysts affect the rate of a reaction? • Catalysts speed up chemical reactions. Only very minute (tiny) quantities of the catalyst are required to produce a dramatic change in the rate of the reaction. This is really because the reaction proceeds by a different pathway when the catalyst is present. Adding extra catalyst will make absolutely no difference. The only other important thing you need to remember about catalysts is that they are not consumed in the course of the reaction. • How does a catalyst work? A catalyst provides a surface on which the reaction can take place. This increases the number of collisions between the particles of the substances that are reacting. • Ammonia oxidized on the surface of the hot Pt wire according to the equation: • 4NH3 + 5O2 =Pt,heat=> 4NO + 5H2O

13. II. Reversible Reactions and EquilibriumA. Reversible Reactions • 1. Reactions sometimes do not go to completion. • 2. Some reactions are reversible. • 3. In reversible reactions, the formation of products occurs along with the decomposition of products and reforming of reactants. • 4. This is represented by a double arrow. • 5. When the forward and reverse reactions occur at the same rate in a reaction, a chemical equilibrium is reached.

14. II. Reversible Reactions and EquilibriumA. Reversible Reactions • 6. At a chemical equilibrium, there is no net change in the actual amounts of the components of the systems. • 7. The equilibrium position of a reaction is given by the relative concentration of the system’s components. • 8. The equilibrium position indicates whether the components on the left or right side of a reversible reaction are at a higher concentration. a. If the higher concentration is on the right side of the reaction, it is a forward reaction. b. If the higher concentration is on the left side of the reaction, it is a reverse reaction.

15. II. Reversible Reactions and EquilibriumA. Reversible Reactions • 9. Example: 2SO2 + O2 2SO3 Forward reaction if more SO3 present Reverse reaction if more O2 or SO2

16. II. Reversible Reactions and EquilibriumB. Factors Affecting Equilibrium 1. A delicate balance exist in a system at equilibrium. 2. Any change in the system will cause it to shift to restore equilibrium. 3. Henri Le Châtelier studied shifts in equilibrium and concluded the following principle: If stress is applied to a system in dynamic equilibrium, the system changes to relieve the stress.

17. II. Reversible Reactions and EquilibriumB. Factors Affecting Equilibrium • 4. Stresses that upset equilibrium can be concentration of reactants or products, changes in pressure, and changes in temperature. • 5. Changes in concentration a. If products are removed, the reaction will shift to replace them and removal of reactants will cause a shift to replace them. b. If products are added, the reaction will shift to form reactants and addition of reactants will cause a shift to form products.

18. II. Reversible Reactions and EquilibriumB. Factors Affecting Equilibrium • 5 c. How does concentration affect the rate of a reaction? • Increasing the concentration of the reactants will increase the frequency of collisions between the two reactants. There are more reacting particles in a given volume. So there are more collisions and therefore more "EFFECTIVE COLLISIONS".

19. II. Reversible Reactions and EquilibriumB. Factors Affecting Equilibrium • 5 d. How does surface area affect a chemical reaction? • Increasing surface area has NO EFFECT on the energy of the particles. It does however increase the number of sites available for a reactions to occur. So more collisions can occur and therefore more "EFFECTIVE COLLISIONS" can happen. • Solids with a smaller particle size (e.g. powders or small chips) react more quickly than solids with a larger particle size (e.g large chips). Note the perimeter of the large cube has 12 sides for a possible reaction. The Small cubes combined have 36 (9*4)sides for a reaction  to occur. 

20. II. Reversible Reactions and EquilibriumB. Factors Affecting Equilibrium • 6. Affects of Temperature a. Increasing temperature will cause a shift in the direction that absorbs heat. b. Decreasing temperature will cause a shift in the direction that releases heat. • 7. Affects of Pressure - Changes in pressure only affect equilibrium that have an unequal number of moles of gaseous reactants and products. a. Increase in pressure will shift to side of the reaction with the least number of moles. b. Decrease in pressure will shift to side the reaction with the greatest number of moles.

21. Question? • Does increasing pressure count as increasing reaction rate? • I would say YES. My reasoning...the pressure is directly related to the concentration of the gases. Therefore, an increase in pressure leads to an increase in reaction rate. Lets file that under concentration.

22. II. Reversible Reactions and EquilibriumC. Equilibrium Constants • 1. Chemist represent equilibrium as numerical value. • 2. This value relates the amount of reactants to products at equilibrium. • 3.This is called the equilibrium constant Keq. • Keq = [products] [reactants] aA + bBDcC + dD Keq = [C]c[D]d [A]a[B]b

23. II. Reversible Reactions and EquilibriumC. Equilibrium Constants • 4. If the equilibrium constant (Keq) is greater than one, the products are favored at equilibrium. • 5. A Keqless than one means the formation of reactants is favored at equilibrium. • 6. Example: In the reaction below, at the equilibrium point, dinitrogentetroxide has a concentration of 0.0055M and nitrogen dioxide has a concentration of 0.025M. What is the equilibrium constant for the reaction? • N2O4D 2NO2

24. Example 6. Example: In the reaction below, at the equilibrium point, dinitrogentetroxide has a concentration of 0.0055M and nitrogen dioxide has a concentration of 0.025M. What is the equilibrium constant for the reaction N2O4D 2NO2 Keq = [NO]2 [N2O4] Keq = [.0055]2 [.025] Keq = .00121

25. III. Determining Whether a Reaction Will Occur A. Free Energy and Spontaneous Reactions • 1. Some chemical and physical processes release energy that can be used to bring about other changes. • 2. This energy is called free energy. • 3. Free energy is the energy that is available to do work. • 4. Spontaneous reaction is a reaction that will occur naturally and favors the formation of the product(Keq>1). • 5. All spontaneous reactions release free energy!

26. III. Determining Whether a Reaction Will Occur A. Free Energy and Spontaneous Reactions e Energy • ΔG = ΔH - TΔS • ΔG= Gibbs Free Energy      ΔH = ENTHALPY T = temperature in Kelvin   ΔS = ENTROPY • Physics Definition-Gibbs free energy, the amount of thermodynamic energy in a fluid system which can be converted into non-mechanical work at a constant temperature and pressure. • What you need to know....If ΔG is negative, the a reaction is spontaneous (it can occur). Spontaneous doesnot imply that it is instantaneous. It could happen over the course of years or decades (I hope you get the point). • ΔG < 0 , the reaction is spontaneousΔG > 0, the reaction is nonspontaneousΔG = 0, the reaction is at equilibrium

27. III. Determining Whether a Reaction Will Occur A. Free Energy and Spontaneous Reactions • 6. Nonspontaneous reaction is a reaction that does not favor formation of the product and does not occur naturally. • 7. Reactions can be spontaneous for one set of conditions but not for another. • 8. Sometimes nonspontaneous reactions can occur when combined with spontaneous reactions. • 9. This does NOT refer to speed of reaction, only whether or not it naturally occurs.

28. III. Determining Whether a Reaction Will Occur A. Free Energy and Spontaneous Reactions 10. We define enthalpy change (DH) to be the measure of change in heat content for a reaction. 11. Since all spontaneous reactions release free energy, you would think that these reactions or processes would always be exothermic (release heat = - DH). 12. This is NOT the case with the melting of ice to water which requires energy (endothermic), but it happens spontaneously. 13. Therefore, enthalpy change is not the only factor that determines whether a reaction will be spontaneous.

29. III. Determining Whether a Reaction Will Occur B. Entropy 1. The law of disorder states that all processes occur in such a way that they move towards maximum disorder (randomness). 2. Entropy is defined as the measure of the amount of disorder in a system. 3. A reaction that absorbs energy (endothermic) can be spontaneous if entropy increases.

30. III. Determining Whether a Reaction Will Occur B. Entropy • 4. How does entropy increase? a. Changing to a less organized phase 1. Solid changing to liquid or gas 2. Liquid changing to gas b. Substance is divided into parts: for example NaCl is separated to Na+ and Cl- in solution c. The product has more molecules than the reactants : for example 2H2O a 2H2 + O2 • 5. Entropy also increases as heat is applied to a system.

31. III. Determining Whether a Reaction Will Occur B. Entropy • This is considered to be the degree of disorder of a system. Gas particles have random motion  have high entropy values. Liquids have much lower entropies and solids lower still. Solutions have a greater entropy than pure liquids as the particles in a solution are more separated and solvent molecules separate the solute particles. • Solids<  Liquids < Aqueous Solutions <  Gases

32. III. Determining Whether a Reaction Will Occur B. Entropy • 6. You have to look at both the change in enthalpy (heat) and the change in entropy (disorder) to determine if a reaction is spontaneous. a. Exothermic + entropy increase = Spontaneous reaction b. Endothermic + entropy increase = Spontaneous reaction ONLY if entropy change is more than heat absorbed c. Exothermic + entropy decrease =spontaneous reaction ONLY if heat change is more than increased order d. Endothermic + entropy decrease = nonspontaneous reaction

33. III. Determining Whether a Reaction Will Occur B. Entropy • 7. The entropy change of a reaction is given the symbol ΔS, units when used are J/mol K (not kJ). • 8. Entropy increases will be observed when any one or more of the following occur: • A reaction breaks up a larger molecule into smaller molecular fragments. • A reaction occurs in which there is an increase in the moles of gas in the product. • A process where a solid changes to a liquid or gas or a liquid changes to a gas. (phase change)

34. Entropy

35. IV. Calculating Entropy and Free EnergyA. Calculating Entropy Change 1. Entropy is a quantitative measure of the disorder of a system. 2.The symbol for entropy is S. 3. Standard entropy (entropy at 25oC and 101.3kPa) is So DSo = Soprod – Soreact 4. Example; Calculate the standard entropy change that occurs when 1 mol H2O(g) condenses to 1 mol H2O(l).

36. IV. Calculating Entropy and Free EnergyB. Calculating Free Energy • 1. In every spontaneous process, some energy becomes available to do work. • 2. This is called Gibbs free energy. • 3. Gibbs free energy is DG. • 4. Because free energy is released in a reaction, when DG is negative (free energy is released) then the reaction is spontaneous. DG = DH – TDS (T in Kelvin) • 5. At STP, 25oC & 101.3kPa, DGo = DHo – TDSo

37. IV. Calculating Entropy and Free EnergyB. Calculating Free Energy 6. Also, if ΔH and ΔS are unknown, you can use the change for the formation of a substance from their elements; DGo = DGoprod – DGoreact 7. Practice Problem: Tell whether or not the following reaction are spontaneous at 25oC: C (s) + O2(g) a CO2 (g)

38. V. The Progress of Chemical ReactionsA. Rate Laws 1. The rate of reaction depends in part on the concentration of reactants. 2. If the reaction is AaB, then the rate of reaction is dependent on the concentration of A with time. 3. The equation is an example of a rate law. 4. The letter k is called a specific rate constant and is proportionality constant relating the concentration of reactants to rate. 5. The equation, Rate=k [ A ]a , is a first order reaction.

39. V. The Progress of Chemical ReactionsA. Rate Laws 6. Or, for A + B a C the rate is first order for A and first order for B and second order for the overall reaction when coefficients a and b each equal 1 Rate=k [ A ]a [ B ]b 7. Ideally, the superscript “a” and “b” would correspond to the coefficient of A and B, but because of the complex steps that many reactions undergo forming intermediates, this is not usually the case.

40. V. The Progress of Chemical ReactionsB. Reaction Mechanisms 1. The simplest reaction progress curve would be obtained for an elementary reaction in which reactants are converted to products in one single step. 2. Most reactions consist of a number of elementary reactions. 3.The reaction mechanism includes all elementary reactions of a complex reaction. 4. There are intermediate product which is a product of one reaction becomes the reactant for another reaction. 5. Intermediates are more stable than activated complex but still highly reactive.