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Ch. 12-13

Ch. 12-13. Reaction Kinetics and Equilibrium. Reaction Kinetics. Looks at the reaction process and the factors that help us predict reactions. Stability. Thermodynamically Stable: Reaction does not spontaneously occur Kinetically Stable: Spontaneous

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Ch. 12-13

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  1. Ch. 12-13 Reaction Kinetics and Equilibrium

  2. Reaction Kinetics • Looks at the reaction process and the factors that help us predict reactions

  3. Stability • Thermodynamically Stable: • Reaction does not spontaneously occur • Kinetically Stable: • Spontaneous • Reaction is occurring so slow it is undetected ( but things are still reacting) • Ex. Decomposition of H2O2 (needs brown bottle)

  4. Reaction Mechanism • Rxn occurs in a series of steps • Reaction Mechanism: • Series of reaction steps that must occur for a reaction to go to completion • Each step has 2 particles colliding

  5. Ex. A + B -> C (step 1) C + D -> E (step 2) E + F -> G (step 3) Total Rxn : A+ B + D + F  G

  6. There were intermediates in between that you never saw (C, E) • you only see the original reactants and the final products • Intermediates: • Something that appears in the series but not in the final product

  7. Ex. N2O  N2 + O (step 1) N2O + O  N2 + O2 (step 2) Total rxn. : 2N2O  2N2 + O2 (O was an intermediate)

  8. Clock Reactions • Reaction Mechanism- teaching it

  9. Homogeneous Reaction • All reactant(s) and product(s) are in the same phase

  10. Heterogeneous Reaction • Reaction that takes place at the interface between 2 phases • Zn (s) + HCl (aq)  H2 (g) + ZnCl2 (aq) ( HCl bubbles on the surface of the Zn)

  11. Collision Theory In order for a rxn. to occur their particles must collide & those collisions must result in interactions • Collisions must: • Collide w/ enough energy • Have particles positioned in a way that enables them to react • Rate of reaction song

  12. Factors that affect reaction rates: • Nature of reaction: • Dependent on the type of bond involved • Ionic reaction rates, faster than covalent • Stirring: • Molecules in faster motion increase probability that the particles will hit & collide w/ enough energy

  13. Crushing: • Smaller pieces increase the surface area so there are more possible sides for collisions LycopodiumSmall scale- creamer Mythbusters-creamer • Concentration: (video) • Quantity of matter that exists in a unit of volume • Increasing concentration increase # of collisions therefore increasing rate • Ex. Double the concentration  4x the collisions

  14. Pressure (works for gases) • Increase pressure, decrease volume • So you have the same # of molecules in a smaller space, more molecules per unit of volume (i.e. higher concentration)  more collisions that could occur  increase rate

  15. Temperature: • Measure of average kinetic energy (frequency of collisions) • Increase that frequency , the collisions increase • Increase temperature does 2 things: • Heating up molecules, moves them faster, more chances for collisions • More kinetic energy in molecules increase the motion of particles, easier to get over that activation energy, rate of reaction will increase

  16. Commercial Break What is this?

  17. A “Cattle List” (get it, a catalyst) Ha, ha, ha, ha, ha, ha, ha,ha, ha, ha, ha (I crack myself up)

  18. Catalyst • A chemical that increases the speed of the reaction but remains chemically unchanged • Doesn’t change the normal position of the equilibrium • Same amounts of product will be formed w/ or w/out the catalyst –just takes longer • Types: homogeneous & heterogeneous • Sugar/sulfuric

  19. Heterogeneous Catalyst • Surface catalyst • Ex. metal oxides, platinum • Works by adsorption – the adherence of one substance to the surface of another • Catalyst has specific lumps that hold the chemicals in the right position to react (increase the chance of them coming together)

  20. Catalytic converter: • Platinum honeycomb structure (more surface area) • Pollution  SO2, CO2, NO • Converter lets H2O react w/ gases to convert them to weak acids (more complete combustion)

  21. Homogeneous Catalyst • In same phase as reactants • Forms an intermediate or activated complex • Reactant reacts better w/ the catalyst than the other reactant • Ex. Sulfuric acid in ester reaction enzymes Colbalt chloride

  22. Entropy • Chemical systems tend to achieve the lowest possible energy state (more stable) • Law of Disorder – states that things move spontaneously in the direction of maximum chaos • Entropy • Can be thought of as– measure of the disorder of the system or the randomness (more stable)

  23. Entropy • More exact definition- measure of the number of possible ways that the energy of a system can be distributed; related to the freedom of the system’s particles to move and the number of ways they can be arranged (energy dispersal)

  24. Misconceptions about Entropy • This view of the second law of thermodynamics is very popular, and it has been misused. Some argue that the second law of thermodynamics means that a system can never become more orderly. Not true. It just means that in order to become more orderly (for entropy to decrease), you must transfer energy from somewhere outside the system, such as when a pregnant woman draws energy from food to cause the fertilized egg to become a complete baby, completely in line with the second line's provisions.

  25. Entropy of gas is greater than liquid or solid • Entropy increase when a substance is divided into parts • Entropy increase w/ increase in temperature

  26. Inhibitors • Prevents reaction from happening for a certain length of time (delays reaction) • Not opposite of catalyst • Ex. Lemon juice on apples • A + B  AB • w/ inhibitor: A-inh + B  no rxn. • Once inhibitor is used up then: A + B  AB

  27. Energy Diagrams • Activation Energy: • Energy required to start a chemical reaction • High activation energy  few collisions have enough energy for a reaction  get slow undetected reaction • Activated Complex: • Product formed when reactants have collided w/ sufficient energy to meet activation energy requirement

  28. Energy Diagram: Exothermic Rxn -releases energy, lower energy after rxn.

  29. Energy diagram: Endothermic Rxn. -absorbed energy, higher energy after rxn.

  30. Endo thermic/exothemic song

  31. Energy Diagram: Catalyst -w/ catalyst product formed faster -lowers the activation energy requirement

  32. Reaction Rate: • Rate of disappearance of one of the reactants or rate of appearance of one of the products • Unit: (mole/L)/s • Change in molarity per second

  33. Reaction rate song (second time)

  34. Rate Law: • Rate is dependent on the concentration of the reactants • Expression relating the rates of reaction to the concentration of reactants • [ ] = concentration

  35. A + B  AB • Rate = k [A] [B] • k= specific rate constant (proportionality constant relating to concentration – value changes depending on rxn.)

  36. Ex. H2 + I2 => 2HI rate= k [H2 ] [I2 ] Exp. 1- [H2 ] = 1.0M [I2 ] = 1.0M rate= .20 M/s k=? .20= k [1.0M] [1.0M] k = .20

  37. Exp. 2 - [H2] = .5 M [I2] = .5 M rate = ? k= .20 rate= k [H2 ] [I2 ] rate = .20 [.5] [.5] rate= .05 M/s

  38. The rate law for elementary reactions is just the product of the reactants, reactions that have more than one step you would need to figure out the order of reaction.

  39. Rate Determining Step • The step or reaction in the series that is slower than all the others  the reaction rate is dependent on this Ex. Person going 45 in the left lane on I-94

  40. Reaction Order or Order of Reaction • Changing the concentration of substances taking part in a reaction usually changes the rate of reaction • A rate equation shows this effect mathematically • Orders of reaction are a part of this rate equation (helps us describe the reaction ) • Orders of Reaction are always found by doing experiments

  41. Elementary Reactions • A reaction with no intermediate steps (very rare) – not a reliable way to determine order • One can determine the order with the coefficients • Rate is proportional to the concentration of the reactants raised to the power of the coefficients Rate expressed as: aA + bB cC + dD Rate = k [A]a [B]b ( a and b are the coefficients)

  42. Reaction Order • Can determine reaction order experimentally or graphically • Experimentally: • Gather data and see what happen to rates if you change the concentration (1st order- double [ ] doubles rate, 2nd order – double[ ] quadruples rate, zero order- rate constant with any [ ] )

  43. Graphically: • Plot concentration vs. time – identify which graph gives you a linear graph • Zero Order: Linear Graph [A] vs time • 1st order: Linear graph ln[A] vs time • 2nd order: Linear graph 1/[A] vs. time

  44. Sum of the power to which all the reactant concentrations are raised (always defined in terms of reactant concentrations (no products)) • Overall order = a + b (exponents added together)

  45. Finding overall Order Ex. Rate = k [A] [B]2 Rate is 1st order for reactant A Rate is 2nd order for reactant B Overall order =(a + b) = 3rd order -if you double [A] = doubles rate -if you double [B] = quad. rate

  46. Practice Problems • Rate Law: • Reaction: • 2NO(g) + Cl2(g) 2 NOCl (g)

  47. Using the following data, calculate the rate law and constant.

  48. What is the rate law? • Rate = k [NO]2 [Cl2] • What is the order of the reaction with respect to NO? • 2nd order • What is the order of the reaction with respect to Cl2 • 1st order

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