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Unit 10: Kinetics, Thermodynamics, & Equilibrium PowerPoint Presentation
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Unit 10: Kinetics, Thermodynamics, & Equilibrium

Unit 10: Kinetics, Thermodynamics, & Equilibrium

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Unit 10: Kinetics, Thermodynamics, & Equilibrium

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  1. Unit 10: Kinetics, Thermodynamics, & Equilibrium Essential Questions: How do chemical reactions happen in the first place? How fast do they go? Unit Exam is Friday May 8th

  2. Kinetics = study of the rate or speed at which reactions occur • A REACTION is the breaking and reforming of bonds to make entirely new compounds as products

  3. Reaction Mechanism = step by step process needed to make a product; how you get from “a” to “b” (like a recipe) • REACTANTS  PRODUCTS • Just like when we bake a cake we must follow directions • CAN’T OMIT any STEPS! • CAN’T CHANGE THE ORDER of the steps! • CAN’T OMIT any REACTANTS (ingredients)

  4. Determine whether each of the following chemical reactions is an example of aslow or fast reaction. Explain why knowing this relative rate of rxn is significant. • Rusting • Alka seltzer in water • Styrofoam decomposing • Weathering of rocks • Bleach removing color

  5. What determines the rate of a reaction? • 1. Number of Steps = more steps can mean a slower reaction • 2. Rate Determining Step = the SLOWEST step of the reaction; most important factor influencing reaction rate

  6. Collision Theory: In order for a reaction to occur, reactant particles must collide and have the following when doing so: • 1. Proper amount of energy • 2. Proper alignment/direction/orientation • *Only when particles collide with these two conditions are met will there be an EFFECTIVE COLLISION, resulting in a reaction. • https://www.youtube.com/watch?v=OttRV5ykP7A • https://www.youtube.com/watch?v=eSInI1xHvh4

  7. THE 6 factors Affecting Rate of Reaction:

  8. INCREASE concentration, increaserxn rate The MORE PARTICLES in a given space, the LESSSPACE b/w particles  MORE COLLISIONS 2. Concentration

  9. INCREASE pressure, INCREASESrxn rate (affects GASES ONLY!) Increasing pressure DECREASES VOLUME which DECREASES SPACE b/w particles  MORE COLLISIONS 3. Pressure 4. Temperature INCREASE temperature, increaserxn rate Greater SPEED MORE total COLLISIONS Greater AVERAGE Kinetic Energy (KE)  collisions take place with MORE energy

  10. INCREASE the surface area (by making PIECES SMALLER) INCREASES therxn rate Increasing surface area EXPOSES MORE REACTANT PARTICLES to possible collisions 5. Surface Area

  11. SPEEDS UP THE RXN WITHOUT CHANGING THE NATURE OF THE REACTANTS/PRODUCTS Provides a SHORTCUT or ALTERNATIVE PATHWAY for the mechanism Lowers the ACTIVATION ENERGY for the reaction 6. Catalyst Genie in a Bottle: https://www.youtube.com/watch?v=5q5bzHckSIM MnO2 H2O2 O2 + H2O

  12. Elephant Toothpaste Demo • https://www.youtube.com/watch?v=p1eG2y2mn54

  13. Unit 10: Kinetics, Thermodynamics, & Equilibrium Essential Question: Why do you have to strike a match just right to get it to light? Unit Exam is Friday May 8th

  14. Potential Energy Diagrams • Recall, we have talked about chemical bonds having stored energy (AKA potential energy). For that reason, chemists use diagrams called Potential Energy Diagrams to illustrate the potential (or stored) energy changes that occur during specific chemical reactions. • Recall: A reaction is the breaking and reforming of bonds • BREAK BONDS  FORM BONDS • A + B  C + D

  15. Heat of Reaction (ΔH) = the amount of HEATENERGY LOST or GAINED throughout a REACTION (ΔH = entHalpy) • PE OF THE PRODUCTS - PE OF THE REACTANTS • ΔH = HP – HR • Also recall, there are two (2) types of reactions: • 1. Reactions that release energy  EXOTHERMIC • ΔH = negative value (-) • energy is released (on right) • A + B  C + D + ENERGY • Example: Sodium in water – lots of heat (and fire!) produced as product; heat felt on a test tube during a reaction

  16. 2. Reactions that absorb/gain energyENDOTHERMIC • ΔH = positive value (+) • energy absorbed (on left) • A + B + ENERGY  C + D • Example: chemical cold packs (feels cold because it’s absorbing heat energy)

  17. Table I (of the Reference Tables) tells us if particular reactions are exothermic or endothermic based on the sign of the ΔH value. • 1. N2(g) + 2O2(g)2NO2(g) • +66.4 kJ • Endothermic • 2. N2(g) + 3H2(g)2NH3(g) • -91.8 kJ • Exothermic • 3. 2NH3(g)N2(g) + 3H2(g) • +91.8 kJ • Endothermic

  18. Forward Reaction = reading LEFT TO RIGHT in a reaction; reaction moves toward the right • A + B  C + D • Reverse Reaction = reading RIGHT TO LEFT in a reaction; reaction moves toward the left • A + B  C + D

  19. Activation Energy = amount of ENERGY needed to get a reaction STARTED or form the ACTIVATED COMPLEX of a reaction (you must get over the “bump” in order for a reaction to be successful) • How exactly does a catalyst shorten the reaction time needed for a reaction to complete? • The ACTIVATED COMPLEX is lowered which also lowers the ACTIVATION ENERGY needed for the reaction to get started (gets rxn started easier and makes the rxn complete itself faster)

  20. ENDOTHERMIC Potential Energy Diagrams  Positive ΔH • Product side (to the Right) is always HIGHER than the reactant side (to the Left) meaning that energy is absorbed. Label the following: A = PE of Reactants B = PE of Products C = PE of Activated Complex D = Activation E of Forward rxn E = Activation E of Reverse rxn F = Heat of Reaction ΔH Activated Complex = Highest energy point of the reaction; this is where full rearrangement of the reactants occurs.

  21. EXOTHERMIC Potential Energy Diagrams  NegativeΔH • *Product side (to the R) is always LOWER than the reactant side (to the L) meaning that energy is released. Label the following: A = PE of Reactants B = PE of Products C = PE of Activated Complex D = Activation E of Forward rxn E = Activation E of Reverse rxn F = Heat of Reaction (ΔH = Hp – Hr) • Question: If a catalyst were added to the above diagram, which letter quantities would change within the diagram? • Question: How does the addition of a catalyst change the heat of reaction (ΔH)? • Question: What are the benefits to adding a catalyst?

  22. Unit 10: Kinetics, Thermodynamics, & Equilibrium Essential Question: What is equilibrium? How can we use it to our advantage? Unit Exam is Friday May 8th

  23. A TES EquilibRium • Some physical and chemical reactions are capable of reaching equilibrium. • Equilibrium occurs WHEN THE RATE OF THE FORWARD REACTION EQUALS THE RATE OF THE REVERSE REACTION in a closed system. • When equilibrium is reached, IT DOES NOT MEAN that the reactants and products are of equal QUANTITIES. So…

  24. A TES EquilibRium • Equilibrium is represented by DOUBLE ARROWS instead of a single arrow. This allows us to illustrate that the reactions are proceeding in both directions (forward and reverse). • Equilibrium is DYNAMIC which means that it is constantly CHANGING or FLUCTUATING • Equilibrium means that reactant and product CONCENTRATIONS are CONSTANT. • *Equilibrium does NOT mean that reactant and product concentrations are equal. • *MEMORIZATION TRICK: “Con Con Requal” - Concentrations are constant, Rates are Equal

  25. Equilibrium Video • Chemical Equilibrium Lesson: https://youtu.be/yFqYrBxbURY • If Molecules were people: https://youtu.be/dUMmoPdwBy4

  26. Equilibrium • Define equilibrium in terms of reactant and product concentrations: • Define equilibrium in terms of forward and reverse reaction rates: Reactant and product concentrations are constant. Forward and reverse reaction rates are equal.

  27. TYPES OF EQUILIBRIUM (all occur in Closed Systems)*It’s all about the equal rates! • 1. Physical Equilibrium : Equilibrium that involves physical changes • a) Phase Equilibrium – occurs during a phase change • (s) (l) Rate of Melting = Rate of Freezing • (sealed container @ 0oC) • (l) (g) Rate of Evaporation = Rate of Condensation • (sealed container at 100oC)

  28. b) Solution Equilibrium – occurs at a solution’s SATURATION POINT • RATE of DISSOLVING = RATE of CRYSTALLIZATION • Example NaCl(s) NaCl(aq)

  29. 2. Chemical Equilibrium • Rate of FORWARD RXN = RATE of the REVERSE RXN • OR • Rate of BREAKING BONDS = Rate of FORMING BONDS

  30. Equilibrium Video • Dynamic Equilibrium: https://youtu.be/wlD_ImYQAgQ

  31. Video: https://youtu.be/dIDgPFEucFM LE CHATELIER’s PRINCIPLE • Le Chatelier’s principle explains HOW A SYSTEM at equilibrium WILL RESPOND TO stress. • STRESS = Any change in TEMPERATURE, CONCENTRATION, or PRESSURE put upon an system at equilibrium • When a STRESS is added to a system at equilibrium, the system will SHIFT in order to relieve that stress and reach a new equilibrium.

  32. SHIFT = an increase in the RATE of EITHER the forward OR the reverse rxn • SHIFT TO RIGHT (TOWARD PRODUCTS): • Rate of FORWARD reaction INCREASES (→) • Reactants Products • *Favors products • SHIFT TO LEFT (TOWARD REACTANTS): • Rate of REVERSE reaction INCREASES (←) • Reactants Products • *Favors reactants

  33. DIFFERENT TYPES OF STRESSES: • 1)Concentration as initial stress: Equilibrium changes (or shifts) when a reactant or product is added (introduced) or decreased (taken away) in a reaction that is at equilibrium • When the concentration of a reactant or product is INCREASED: the reaction will SHIFT AWAY from the increase (use up the excess)

  34. Example 1: 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) + HEAT • If we add H2O(g), the system would shift to the _______ and the [NH3] would _________________. • If we add O2(g), the system would shift to the _______ and the [NO] would _________________. • If we add H2O(g), the system would shift to the _______ and the [NO] would _________________. • If we added NO(g), which concentration(s) would decrease? _________________ Left increase Right increase Left decrease H2O & Heat

  35. Le Chatelier’s Principle • Part 1 (Conc. & Pressure) Video: https://www.youtube.com/watch?v=7zuUV455zFs

  36. Left • When the concentration of a reactant or product is DECREASED: the reaction will SHIFT TOWARD the side that has experienced the decrease in concentration (replaces what was taken) • Example 2: 2804 kJ + 6CO2(g) + 6H2O(l) C6H12O6(s) + 6O2(g) • 1. If we remove water, the system will shift to the _______ and the [C6H12O6] will _________________. • 2. If we remove O2(g), the system will shift to the _______ and the [C6H12O6] will _________________. • 3. If we remove glucose, which concentrations will decrease? • 4. If we remove CO2, which concentration(s) would increase? decrease Right increase CO2 and H2O CO2 and H2O *TRICK: “AA” – what YOU ADD, the SYSTEM shifts AWAY “TT” – what YOU TAKE, the SYSTEM shifts TOWARD

  37. Le Chatelier’s Principle • Part 2 (TEMP) Video: https://youtu.be/XhQ02egUs5Y

  38. 2) Temperature as initial stress: (involves increasing or decreasing the “HEAT” component of a reaction) • Note: HEAT/ENERGY/J/KJ will either be a reactant or a product • A + B C + D + HEAT • A + B + energy C + D • When temperature (or HEAT) is increased: the reaction will SHIFT AWAY from the rxn side containing “HEAT” (in the ENDOTHERMIC direction) • When temperature (or HEAT) is decreased: the reaction will SHIFT TOWARD the rxn side containing “HEAT” (in the EXOTHERMIC direction)

  39. NO and H2O • Example #1: 4NH3(g) + 5O2(g)4NO(g) + 6H2O(g) + HEAT • 1. If we added heat, which concentration(s) will decrease? • 2. If we added heat, which concentration(s) will increase? • Example #2: CO2(g) + H2O(l) + 890.4 kJ CH4(g) + 2O2(g) • 3. If we remove heat, which concentration(s) will decrease? • 4. If we remove heat, which concentration(s) will increase? NH3 and O2 CH4 and O2 CO2 and H2O

  40. 3) Pressure as initial stress: Recall, pressure affects GASES ONLY! So every other state (s, l, aq) in the reaction is UNAFFECTED for this type of stress • INCREASE PRESSURE: rxn shifts to side with LEAST # GAS MOLECULES (or least # moles of gas) • DECREASE PRESSURE: rxn shifts to side with GREATEST # GAS MOLECULES (or greatest # moles of gas) • NOTE: If the rxn contains NO GAS MOLECULES or if the rxnhas the SAME # GAS MOLECULES on each side, there is NO EFFECT and NO SHIFT results from an increase or decrease in pressure

  41. 1 gas molecule No gas molecules Example 1: CO2(g)CO2(aq) CO2(aq) • 1. If we increase the pressure, the concentrations of which species will increase? • 2. If we increase the pressure, the concentrations of which species will decrease? • 3. If we decrease the pressure, the concentrations of which species will increase? • 4. If we decrease the pressure, the concentrations of which species will decrease? CO2(g) CO2(g) CO2(aq)

  42. Example 2: N2(g) + 3H2(g)2NH3(g) • 1. If we increase the pressure, in which direction will the equilibrium shift? (Count moles of gases on each side 1st) • 2. If we increase the pressure, the concentration of which species will increase initially? • 3. If we decrease the pressure, the concentration of which species will decrease initially? • 4. If we decrease the pressure, the concentration of which species will increase initially? RIGHT NH3 N2 and H2 NH3

  43. States of Matter

  44. Gas Video • Invisible Properties of Gases - http://ed.ted.com/lessons/describing-the-invisible-properties-of-gas-brian-bennett

  45. Kinetic Molecular Theory (KMT): A theory used to explain the behavior of gases in terms of the MOTION of their particles. • Major Assumptions of KMT: • 1. Gas particles have NO volume (take up no space) • 2. The gas particles do NOT attract each other (NO intermolecular forces) • 3. The gas particles move in STRAIGHT, continuous motion. • 4. The gas particles are perfectly ELASTIC – lose NO speed after they COLLIDE w/one another or w/ the walls of the container.

  46. KMT is based on the CONCEPT or MODEL of an IDEAL GAS • An IDEAL GAS is THEORETICAL and is used to PREDICT the behavior of REAL GASES (O2, H2, He, etc.)

  47. *Trick: What is your ideal Boyfriend/Girlfriend? IdealBoyfriend/Girlfriend = Hot and No Pressure How to make a “real” BF/GF into an “ideal” BF/GF  Increase Temp; Lower Pressure REAL GASES: Deviate (are different) from ideal gas laws when molecules are closetogether. • Why? • Deviate from #1: Real gases contain particles that have volume. • Deviate from #2: Real gas particles are attracted to each other. • *These deviations occur especially under high PRESSURE and low TEMPERATURE. • QUESTION: When does a REAL GAS behave most nearly like an IDEAL GAS? • At high temp and low pressure

  48. Practice Regents • 1. According the kinetic molecular theory of gases, which assumption is correct? • Gas particles strongly attract each other. • Gas particles travel in curved paths • The volume of gas particles prevents random motion. • Energy may be transferred between colliding particles. • 2. A real gas would behave most like an ideal gas under conditions of • low pressure and low temperature • low pressure and high temperature • high pressure and low temperature • high pressure and high temperature * *

  49. Gas Video • ABC’s of Gases - http://ed.ted.com/lessons/1207-1-a-bennet-brianh264

  50. Combined Gas Law • Boyle’s Law: P1V1 = P2V2 • Charles’s Law: V1 = V2 T1 T2