UNIT 8 KINETICS & EQ: NOTE & PRACTICE PACKET

Transcription

1 UNIT 8 KINETICS & EQ: NOTE & PRACTICE PACKET 1

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3 Lesson 1: Kinetics = study of the RATE or SPEED at which REACTIONS occur A REACTION is the Reaction Mechanism = STEP BY STEP PROCESS needed to make a product; how you get from a to b (like a recipe) Just like when we bake a cake we must follow directions o CAN T OMIT any STEPS! o CAN T CHANGE THE ORDER of the steps! o CAN T OMIT any REACTANTS (ingredients) Determine whether each of the following chemical reactions is an example of a slow 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 WHAT DETERMINES THE RATE OF A REACTION? 1. NUMBER OF STEPS = more steps can mean a slower reaction 2. RATE DETERMINING STEP = the 3

4 Collision Theory: In order for a reaction to occur, * Lesson 2: Factors Affecting Rate of Reaction Factor How Rate Affected Why does it increase the rate? Ionic = smaller 1. Nature of Reactants NaCl (aq) + AgNO 3(aq) NaNO (aq) + AgCl (s) Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3 - (aq) Na + (aq) + NO 3 - (aq) AgCl (s) (1 step) Covalent = larger CH 4( g) + O 2( g) CO 2( g) + 2H 2 O ( l) (break 4 C-H bonds, 1 O-O bond, form 2 C-O bonds, and 4 O-H bonds) 4

5 2. Concentration INCREASE concentration, 3. Pressure INCREASE pressure, Increasing pressure 4. Temperature INCREASE temperature, Greater SPEED à Greater AVERAGE KE à INCREASE the surface area Increasing surface area 5. Surface Area (How many surfaces are there?) 5

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7 Classwork 8-1 & 8-2: 1. In order for a reaction to occur the particles must with proper and. Therefore, the more collisions the reactant particle have, the faster the rate. 2. Recall 5 ways to increase the rate of reaction. Be specific Matches have the potential to burn on fire. But they will not without sufficient activation energy. Explain what activation energy means and what type of activation energy the matches need. 4. Which event must always occur for a chemical reaction to take place? A) formation of a precipitate B) formation of a gas C) effective collisions between reacting particles D) addition of a catalyst to the reaction system 5. Increasing the temperature increases the rate of a reaction by A) lowering the activation energy B) increasing the activation energy C) lowering the frequency of effective collisions between reacting molecules D) increasing the frequency of effective collisions between reacting molecules 6. After being ignited in a Bunsen burner flame, a piece of magnesium ribbon burns brightly, giving off heat and light. In this situation, the Bunsen burner flame provides A) ionization energy B) activation energy C) heat of reaction D) heat of vaporization 7. As the number of effective collisions between reacting particles increases, the rate of reaction A) Decreases B) increases C) remains the same 8. In most aqueous reactions as temperature increases, the effectiveness of collisions between reacting particles A) Decreases B) increases C) remains the same 9. Given the reaction: Mg + 2 H 2 O Mg(OH) 2 + H 2 At which temperature will the reaction occur at the greatest rate? A) 25ºC B) 50ºC C) 75ºC D) 100ºC 7

8 10. A 5.0-gram sample of zinc and a 50.-milliliter sample of hydrochloric acid are used in a chemical reaction. Which combination of these samples has the fastest reaction rate? A) a zinc strip and 1.0 M HCl(aq) B) a zinc strip and 3.0 M HCl(aq) C) zinc powder and 1.0 M HCl(aq) D) zinc powder and 3.0 M HCl(aq) 11. A 1.0-gram piece of zinc reacts with 5 milliliters of HCl(aq). Which of these conditions of concentration and temperature would produce the greatest rate of reaction? A) 1.0 M HCl(aq) at 20. C B) 1.0 M HCl(aq) at 40. C C) 2.0 M HCl(aq) at 20. D) 2.0 M HCl(aq) at 40. C 12. At STP, which 4.0-gram zinc sample will react fastest with dilute hydrochloric acid? A) lump C) bar B) powdered D) sheet metal 13. Given the reaction: Fe(s) + 2 HCl(aq) FeCl2(aq) + H2(g) In this reaction, 5 grams of powdered iron will react faster than a 1-gram piece of solid iron because the powdered iron A) has less surface area B) has more surface area C) is less dense D) D) is more dense 14. Which statement best explains the role of a catalyst in a chemical reaction? A) A catalyst is added as an additional reactant and is consumed but not regenerated. B) A catalyst limits the amount of reactants used. C) A catalyst changes the kinds of products produced. D) A catalyst provides an alternate reaction pathway that requires less activation energy. 15. Which change would most likely increase the rate of a chemical reaction? A) decreasing a reactant's concentration B) decreasing a reactant's surface area C) cooling the reaction mixture D) adding a catalyst to the reaction mixture 8

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10 Lesson 3: Heat of Reaction,ΔH 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 à FORM BONDS A + B à C + D Heat of Reaction (ΔH) = PE OF THE PRODUCTS - PE OF THE REACTANTS Also recall, there are two (2) types of reactions: 1. Reactions that release energy à A + B à C + D + ENERGY Ex: Sodium in water heat (and fire!) as product 2. Reactions that absorb/gain energy à A + B + ENERGY à C + D Ex: baking (need oven to supply heat) 10

11 Whatever you do to a chemical equation, you also must do to the ΔH. Ex 1: If you REVERSE a reaction (flip the products and reactants), then reverse you must the sign for ΔH. Ex 2: If you double the equation (or the coefficients), then you must double the ΔH. Table I (of the Reference Tables) tells us if particular reactions are exothermic or endothermic based on sign of the Δ H value. Δ H (kj/mol) Endothermic/Exothermic 1. N 2 (g) + 2O 2 (g) 2NO 2 (g) 2. N 2 (g) + 3H 2 (g) 2NH 3 (g) 3. 2NH 3 (g) N 2 (g) + 3H 2 (g) 4. 2C(s) + 2O 2 (g) 2CO 2 (g) 5. CO 2 (g) CO(g) + ½ O 2 (g) 8 11

12 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 Activation Energy = HOW EXACTLY DOES A CATALYST SHORTEN THE REACTION TIME NEEDED FOR A REACTION TO COMPLETE? The ACTIVATED COMPLEX is lowered OR The ACTIVATION ENERGY is decreased OR The REACTION PATHWAY is shortened 12

13 Classwork 8-3: Use Table I to complete the chart below: 13

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15 Lesson 4A:ENDOTHERMIC Potential Energy Diagrams à * Product Label the following: A = Potential Energy of the Reactants B = Potential Energy of the Products C = Potential Energy of the Activated Complex D = Activation energy of the Forward rxn E = Activation Energy of the Reverse rxn F = Heat of the Reaction (ΔH = H p H r ) ACTIVATED COMPLEX = 15

16 Lesson4B: EXOTHERMIC Potential Energy Diagrams à * Product Label the following: A = Potential Energy of the Reactants B = Potential Energy of the Products C = Potential Energy of the Activated Complex D = Activation energy of the Forward rxn E = Activation Energy of the Reverse rxn F = Heat of the Reaction (ΔH = H p H r ) Question: If a catalyst were added to the above diagram, which letter quantities would change within the diagram? Answer: Question: How does the addition of a catalyst change the heat of reaction (ΔH)? (Increase, decrease, or remains the same) Answer: Question: What are the benefits to adding a catalyst? Answer: 16

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22 Lesson 5: EQUILIBRIUM Some physical and chemical reactions are capable of reaching equilibrium. When equilibrium is reached, IT DOES NOT MEAN that the reactants and products are of equal QUANTITIES. So o Equilibrium is o Equilibrium is o Equilibrium means that Define equilibrium in terms of reactant and product concentrations: Define equilibrium in terms of forward and reverse reaction rates: 22

23 TYPES OF EQUILIBRIUM (all occur in ) *IT S ALL ABOUT THE EQUAL RATES! 1. Physical Equilibrium: Equilibrium that involves physical changes a) Phase Equilibrium Examples: (sealed 0ºC) (sealed 100ºC) b) Solution Equilibrium occurs at a solution s example: 2. Chemical Equilibrium: O R RATE of BREAKING BONDS = RATE of FORMING BONDS 23

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27 Lesson 6A:LE CHATELIER s PRINCIPLE chemistry/essentialchemistry/ flash/lechv17.swf Le Chatelier s principle explains HOW A SYSTEM WILL RESPOND TO. STRESS = 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. SHIFT = 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 27

28 Lesson 6B: TYPES OF STRESSES-CONCENTRATION 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 Example 1: 4NH 3 (g) + 5O 2 (g) 4NO(g) + 6H 2 O(g) + HEAT 1. If we add H 2 O(g), the system would shift to the and the [NH 3 ] would. 2. If we add O 2 (g), the system would shift to the and the [NO] would. 3. If we add H 2 O(g), the system would shift to the and the [NO] would. 4. If we added NO(g), which concentration(s) would decrease? Example 1: 4NH 3 (g) + 5O 2 (g) 4NO(g) + 6H 2 O(g) + HEAT 1. If we remove oxygen, the system will shift to the and the [NH 3 ] will. 2. If we remove water, the system will shift to the and the [NO] will. 3. If we remove ammonia, which concentration(s) will decrease? 4. If we remove NO(g), which concentration(s) would increase? TRICK AA what YOU ADD, the SYSTEM shifts AWAY from TT what YOU TAKE, the SYSTEM shifts TOWARDS 28

29 Lesson 6C: TYPES OF STRESSES-TEMPERATURE 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: When temperature (or HEAT) is decreased: Example #1: 4NH 3 (g) + 5O 2 (g) 4NO(g) + 6H 2 O(g) + HEAT 1. If we added heat, which concentration(s) will decrease? 2. If we added heat, which concentration(s) will increase? Example #2: CO 2 (g) + H 2 O(l) kj CH 4 (g) + 2O 2 (g) 3. If we remove heat, which concentration(s) will decrease? 4. If we remove heat, which concentration(s) will increase? 29

30 Lesson 6D: TYPES OF STRESSES-PRESSURE Pressure as initial stress: Recall, INCREASE PRESSURE: DECREASE PRESSURE: NOTE: Example 1: CO 2 (g) CO 2 (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? 30

31 Example 2: N 2 (g) + 3H 2 (g) 2NH 3 (g) 1. If we increase the pressure, in which direction will the equilibrium shift? (Count moles of gases on each side 1 st ) 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? AND LASTLY WHY DO CHEMICAL AND PHYSICAL CHANGES OCCUR? Turn the page please 31

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39 Lesson 7: Entropy ENTROPY (ΔS): The MORE ORDER you have, the LESS ENTROPY in your system. The LESS ORDER you have, the MORE ENTROPY in your system. is the most significant factor in determining : Changing from (s) à (l) à (aq) à (g) = Draw particle diagrams to illustrate each of the following phases: s l aq g *Entropy. *Entropy NOTE: If there is no phase change, count up the # molecules on each side (RULE: # moles â = =, # moles á = = ) For the following determine if there is an increase, decrease, or no change in entropy: 1. 2KClO 3(s) 2KCl (s) + 3O 2(g) 9. H 2(g) + Cl 2(g) 2HCl (g) 2. H 2 O (l) H 2 O (s) 10. Ag + (aq) + Cl - (aq) AgCl (s) 3. N 2(g) + 3H 2(g) 2NH 3(g) 11. 2N 2 O 5(g) 2NO 2(g) + O 2(g) 4. NaCl (s) Na + (aq) + Cl - (aq) 12. 2Al (s) + 2I 2(s) 2AlI 3(s) 5. KCl (s) KCl (l) 13. H + (aq) + OH - (aq) H 2 O (l) 6. CO 2(s) CO 2(g) 14. 2NO (g) N 2(g) + O 2(g) 7. H + (aq) + C 2 H 3 O 2 - (aq) HC 2 H 3 O 3(l) 15. H 2 O (g) H 2 O (s) 8. C (s) + O 2(g) CO 2(g) 39

40 Classwork 8-7: