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Reaction Rate How Fast Does the Reaction Go?. Collision Theory l In order to react molecules and atoms must touch each other. l They must hit each other.

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Presentation on theme: "Reaction Rate How Fast Does the Reaction Go?. Collision Theory l In order to react molecules and atoms must touch each other. l They must hit each other."— Presentation transcript:

1 Reaction Rate How Fast Does the Reaction Go?

2 Collision Theory l In order to react molecules and atoms must touch each other. l They must hit each other hard enough to react. – Must break bonds l Anything that increases how often and how hard will make the reaction faster.

3 Energy Reaction coordinate Reactants Products

4 Energy Reaction coordinate Reactants Products Activation Energy - Minimum energy to make the reaction happen – how hard

5 Energy Reaction coordinate Reactants Products Activated Complex or Transition State

6 Activation Energy l Must be supplied to start the reaction l Low activation energy – Lots of collision are hard enough – fast reaction l High Activation energy – Few collisions hard enough – Slow reaction

7 Activation energy l If reaction is endothermic you must keep supplying heat l If it is exothermic it releases energy l That energy can be used to supply the activation energy to those that follow

8 Energy Reaction coordinate Reactants Products Overall energy change

9 Things that Affect Rate l Temperature – Higher temperature faster particles. – More and harder collisions. – Faster Reactions. l Concentration – More concentrated molecules closer together – Collide more often. – Faster reaction.

10 Things that Affect Rate l Particle size – Molecules can only collide at the surface. – Smaller particles bigger surface area. – Smaller particles faster reaction. – Smallest possible is molecules or ions. – Dissolving speeds up reactions. – Getting two solids to react with each other is slow.

11 Things that Affect Rate l Catalysts- substances that speed up a reaction without being used up.(enzyme). l Speeds up reaction by giving the reaction a new path. l The new path has a lower activation energy. l More molecules have this energy. l The reaction goes faster. l Inhibitor- a substance that blocks a catalyst.

12 Energy Reaction coordinate Reactants Products

13 Pt surface HHHH HHHH l Hydrogen bonds to surface of metal. l Break H-H bonds Catalysts

14 Pt surface HHHH Catalysts C HH C HH

15 Pt surface HHHH Catalysts C HH C HH l The double bond breaks and bonds to the catalyst.

16 Pt surface HHHH Catalysts C HH C HH l The hydrogen atoms bond with the carbon

17 Pt surface H Catalysts C HH C HH HHH

18

19 Reversible Reactions Reactions are spontaneous if G is negative. If G is positive the reaction happens in the opposite direction. 2H 2 (g) + O 2 (g) 2H 2 O(g) + energy 2H 2 O(g) + energy H 2 (g) + O 2 (g) 2H 2 (g) + O 2 (g) 2H 2 O(g) + energy

20 Equilibrium l When I first put reactants together the forward reaction starts. l Since there are no products there is no reverse reaction. l As the forward reaction proceeds the reactants are used up so the forward reaction slows. l The products build up, and the reverse reaction speeds up.

21 Equilibrium l Eventually you reach a point where the reverse reaction is going as fast as the forward reaction. l This is dynamic equilibrium. l The rate of the forward reaction is equal to the rate of the reverse reaction. l The concentration of products and reactants stays the same, but the reactions are still running.

22 Equilibrium l Equilibrium position- how much product and reactant there are at equilibrium. l Shown with the double arrow. l Reactants are favored l Products are favored l Catalysts speed up both the forward and reverse reactions so dont affect equilibrium position.

23 Equilibrium l Catalysts speed up both the forward and reverse reactions so dont affect equilibrium position. l Just get you there faster

24 Measuring equilibrium l At equilibrium the concentrations of products and reactants are constant. l We can write a constant that will tell us where the equilibrium position is. l K eq equilibrium constant l K eq = [Products] coefficients [Reactants] coefficients l Square brackets [ ] means concentration in molarity (moles/liter)

25 Writing Equilibrium Expressions l General equation aA + bB cC + dD l K eq = [C] c [D] d [A] a [B] b l Write the equilibrium expressions for the following reactions. l 3H 2 (g) + N 2 (g) 2NH 3 (g) l 2H 2 O(g) 2H 2 (g) + O 2 (g)

26 Calculating Equilibrium l K eq is the equilibrium constant, it is only effected by temperature. l Calculate the equilibrium constant for the following reaction. 3H 2 (g) + N 2 (g) 2NH 3 (g) if at 25ºC there 0.15 mol of N 2, 0.25 mol of NH 3, and 0.10 mol of H 2 in a 2.0 L container.

27 What it tells us l If K eq > 1 Products are favored – More products than reactants at equilibrium l If K eq < 1 Reactants are favored

28 LeChâteliers Principle Regaining Equilibrium

29 LeChâteliers Principle l If something is changed in a system at equilibrium, the system will respond to relieve the stress. l Three types of stress are applied. – Changing concentration – Changing temperature – Changing pressure

30 Changing Concentration l If you add reactants (or increase their concentration). l The forward reaction will speed up. l More product will form. l Equilibrium Shifts to the right Reactants products

31 Changing Concentration l If you add products (or increase their concentration). l The reverse reaction will speed up. l More reactant will form. l Equilibrium Shifts to the left Reactants products

32 Changing Concentration l If you remove products (or decrease their concentration). l The reverse reaction will slow down. l More product will form. l Equilibrium reverseShifts to the right Reactants products

33 Changing Concentration l If you remove reactants (or decrease their concentration). l The forward reaction will slow down. l More reactant will form. l Equilibrium Shifts to the left. Reactants products l Used to control how much yield you get from a chemical reaction.

34 Changing Temperature l Reactions either require or release heat. l Endothermic reactions go faster at higher temperature. l Exothermic go faster at lower temperatures. l All reversible reactions will be exothermic one way and endothermic the other.

35 Changing Temperature l As you raise the temperature the reaction proceeds in the endothermic direction. l As you lower the temperature the reaction proceeds in the exothermic direction. Reactants + heat Products at high T Reactants + heat Products at low T l H 2 O (l) H 2 O(s) + heat

36 Changes in Pressure l As the pressure increases the reaction will shift in the direction of the least gases. At high pressure 2H 2 (g) + O 2 (g) 2 H 2 O(g) At low pressure 2H 2 (g) + O 2 (g) 2 H 2 O(g) l Low pressure to the side with the most gases.

37 Three Questions l How Fast? – Depends on collisions and activation energy – Affected by Temperature Concentration Particle size Catalyst l Reaction Mechanism – steps

38 Three Questions l Will it happen? – Likely if ΔH is negative – exothermic Or ΔS is positive – more disorder – Guaranteed if ΔG is negative ΔG o f Products – Reactants Or ΔG = ΔH -T ΔS

39 Three Questions l How far? – Equilibrium Forward and reverse rates are equal Concentration is constant – Equilibrium Constant One for each temperature – LeChâteliers Principle

40 Thermodynamics Will a reaction happen?

41 Energy l Substances tend react to achieve the lowest energy state. l Most chemical reactions are exothermic. l Doesnt work for things like ice melting. l An ice cube must absorb heat to melt, but it melts anyway. Why?

42 Entropy l The degree of randomness or disorder. l Better – number of ways things can be arranged lSlS l The First Law of Thermodynamics - The energy of the universe is constant. l The Second Law of Thermodynamics - The entropy of the universe increases in any change. l Drop a box of marbles. l Watch your room for a week.

43 Entropy Entropy of a solid Entropy of a liquid Entropy of a gas l A solid has an orderly arrangement. l A liquid has the molecules next to each other but isnt orderly l A gas has molecules moving all over the place.

44 Entropy increases when... l Reactions of solids produce gases or liquids, or liquids produce gases. l A substance is divided into parts -so reactions with more products than reactants have an increase in entropy. l The temperature is raised -because the random motion of the molecules is increased. l a substance is dissolved.

45 Entropy calculations l There are tables of standard entropy (pg 407). l Standard entropy is the entropy at 25ºC and 1 atm pressure. l Abbreviated Sº, measure in J/K. The change in entropy for a reaction is Sº= Sº(Products)-Sº(Reactants). Calculate Sº for this reaction CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O(g)

46 l For CH 4 Sº = J/K-mol l For O 2 Sº = J/K-mol l For CO 2 Sº= J/K-mol l For H 2 O(g) Sº = J/K-mol

47 Spontaneity Will the reaction happen, and how can we make it?

48 Spontaneous reaction l Reactions that will happen. l Nonspontaneous reactions dont. l Even if they do happen, we cant say how fast. l Two factors influence. l Enthalpy (heat) and entropy(disorder).

49 Two Factors l Exothermic reactions tend to be spontaneous. – Negative H. l Reactions where the entropy of the products is greater than reactants tend to be spontaneous. – Positive S. A change with positive S and negative H is always spontaneous. A change with negative S and positive H is never spontaneous.

50 Other Possibilities l Temperature affects entropy. l Higher temperature, higher entropy. l For an exothermic reaction with a decrease in entropy (like rusting). l Spontaneous at low temperature. l Nonspontaneous at high temperature. l Enthalpy driven.

51 Other Possibilities l An endothermic reaction with an increase in entropy like melting ice. l Spontaneous at high temperature. l Nonspontaneous at low temperature. l Entropy driven.

52 Gibbs Free Energy l The energy free to do work is the change in Gibbs free energy. Gº = Hº - T Sº (T must be in Kelvin) l All spontaneous reactions release free energy. So G <0 for a spontaneous reaction.

53 G= H-T S H G Spontaneous? --+ At all Temperatures S ?++ At high temperatures, entropy driven ?-- At low temperatures, enthalpy driven ++- Not at any temperature, Reverse is spontaneous

54 Problems l Using the information on page 407 and pg 190 determine if the following changes are spontaneous at 25ºC. 2H 2 S(g) + O 2 (g) 2H 2 O(l) + S(rhombic)

55 2H 2 S(g) + O 2 (g) 2H 2 O(l) + 2S l From Pg. 190 we find H f ° for each component – H 2 S = kJO 2 = 0 kJ – H 2 O = kJS = 0 kJ l Then Products - Reactants l H =2 ( kJ) + 2(0 kJ) - 2 (-20.1 kJ) - 1(0 kJ) = kJ

56 2H 2 S(g) + O 2 (g) 2H 2 O(l) + 2 S l From Pg. 407 we find S for each component – H 2 S = J/K O 2 = J/K – H 2 O = J/K S = 31.9 J/K l Then Products - Reactants l S= 2 (69.94 J/K) + 2(31.9 J/K) - 2(205.6 J/K) J/K = J/K

57 2H 2 S(g) + O 2 (g) 2H 2 O(l) + 2 S l G = H - T S l G = kJ - 298K ( J/K) l G = kJ J l G = kJ kJ l G = kJ l Spontaneous l Exergonic- it releases free energy. l At what temperature does it become spontaneous? l G = kJ J

58 Spontaneous l It becomes spontaneous when G = 0 l Thats where it changes from positive to negative. l Using 0 = H - T S and solving for T l 0 - H = - T S l - H = -T S l T = H = S = 1290 K kJ J/K = J J/K

59 Theres Another Way l There are tables of standard free energies of formation compounds.(pg 414) Gº f is the free energy change in making a compound from its elements at 25º C and 1 atm. for an element Gº f = 0 l Look them up. Gº= Gº f (products) - Gº f (reactants)

60 2H 2 S(g) + O 2 (g) 2H 2 O(l) + 2S l From Pg. 414 we find G f ° for each component – H 2 S = kJO 2 = 0 kJ – H 2 O = kJS = 0 kJ l Then Products - Reactants l G =2 (-237.2) + 2(0) - 2 (-33.02) - 1(0) = kJ

61 Does ice melt? l For the following change – H 2 O(s) H 2 O(l) ΔH° =6.03 kJ and ΔS° =22.1 J/K At what temperature does ice melt?

62 Reaction Mechanism l Elementary reaction- a reaction that happens in a single step. l Reaction mechanism is a description of how the reaction really happens. l It is a series of elementary reactions. l The product of an elementary reaction is an intermediate. l An intermediate is a product that immediately gets used in the next reaction.

63 l This reaction takes place in three steps

64 EaEa First step is fast Low activation energy

65 Second step is slow High activation energy EaEa

66 EaEa Third step is fast Low activation energy

67 In this case the second step is rate determining It is slowest Highest activation energy

68 Intermediates are present

69 Activated Complexes or Transition States

70 Mechanisms and rates l Intermediates are stable -they last for a little time l Activated complexes dont l There is an activation energy for each elementary step. l Slowest step (rate determining) must have the highest activation energy.

71 l The mechanism for the decomposition of hydrogen peroxide is l Which is the rate determining step? l What are the intermediates? l Sketch the potential energy diagram. Slow Fast


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