Unit 8 Kinetics & Equilibrium Notes & CW

Transcription

1 Unit 8 Kinetics & Equilibrium Notes & CW 2

2 The bold, underlined words are important vocabulary words that you should be able to define and use properly in explanations. This is a study guide for what you will be tested on throughout the year. The objectives are divided into categories of Knowledge (what you have to know) and Application (what you have to be able to do). o o o o o o o o o o o o KINETICS AND EQUILIBRIUM Knowledge Application The Collision Theory states that a chemical reaction is most likely to occur if reactant particles collide with the proper energy and orientation. o Use the Collision Theory to explain how factors such as temperature, surface area, and concentration influence the rate of reaction Ex: Increasing the temperature, surface area, or concentration all lead to an increase in the rate of a reaction because they all increase the number of effective collisions between reactant particles. The rate (speed) of a chemical reaction depends on several factors: temperature, concentration, nature of reactants, surface area, and the presence of a catalyst. Ionic compounds generally react faster than covalent (molecular) compounds A catalyst provides an alternate reaction pathway, which has lower activation energy than an uncatalyzed reaction. Energy released or absorbed during a chemical reaction can be represented by a potential energy diagram. The difference in PE of the products and reactants is called the heat of reaction ( H) H = PE products PE reactants H values for many chemical reactions are listed in Table I Chemical and physical changes can reach equilibrium Saturated solutions are examples of systems in physical equilibria (aq s) At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction and the measurable quantities of reactants and products remain constant at equilibrium*care* LeChatelier s principle can be used to predict the effect of a stress (such as a change in pressure, volume, concentration, or temperature) on a system at equilibrium. According to LeChatelier s principle, a system at equilibrium will shift to reduce the effects of a stress placed on the system. It will shift AWAY from an INCREASE and will shift toward a decrease in concentration or temperature ( shift means that either 2 o o o o o o o o Explain, in terms of the number of bonds broken, why ionic compounds generally react faster than covalent compounds Explain how a catalyst speeds up a reaction Read and interpret a potential energy diagram Draw and label the following parts of a potential energy diagram for both an endothermic and exothermic reaction PE of reactants and PE of products heat of reaction ( H) activation energy (for both the forward and reverse reactions) activation energy with a catalyst present Distinguish between examples of physical equilibria and chemical equilibria Describe what is happening to the concentrations or amounts of reactants and products in a system at equilibrium Describe the rates of opposing reactions in a system at equilibrium Describe, in terms of LeChatelier s principle, the effects of stress on a given system at equilibrium, including: Changing the temperature/heating/cooling Changing the concentration of a reactant or product Changing the pressure or volume (this affects systems involving gases)

3 o o the forward or the reverse reaction will be favored (go faster) until the rates are again equal and equilibrium is re-established). Changing the pressure or volume only affects systems that contain gases o Also be able to explain why any shifting occurs in terms of Collision Theory Systems in nature tend to undergo changes toward lower energy and higher entropy. 3

4 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) 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 4

5 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) 5

6 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?) 6

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10 Lab Activity: Popping Film Canisters and Rate of Reaction Wear Goggles! Purpose Research which variable affects the rate that Alka Seltzer tablets dissolve in water. Introduction - Alka Seltzer is an over the counter indigestion medicine. A fun way to study dissolving rates is to place Alka Seltzer in water in a film canister and time how long it takes until the lid pops off. Each time you place the Alka Seltzer in the film canister, you will first change one variable to see how that affects the time until the top pops off. Background The active ingredients of Alka Seltzer tablets are sodium bicarbonate (baking soda), citric acid from fruit juice and acetylsalicylic acid (aspirin). Bubbles of CO 2 gas are produced whenever an acid is added to baking soda. In the present case, there are two acids citric acid and aspirin. So there are two reactions producing carbon dioxide. The first reaction is between the baking soda and citric acid. The second reaction is between the baking soda and the aspirin. Alka Seltzer is an anti-acid because both the sodium citrate and the sodium acetylsalicylate neutralize the excess stomach acid. The carbonation produced by dissolving carbon dioxide gas in water also settles the stomach, as anyone knows who drinks a carbonated soft drink to treat a mild upset stomach. Procedure Your group will rotate through 4 labs stations (about 5 mins/ station). At each lab station, you will test one variable to see how it affects the time it takes to pop the top of the film canister. At Each Lab Table, Read ALL directions before beginning. Be ready to complete steps quickly and to ACCURATELY record the time it takes the top to pop off. Safety Point canister tops AWAY from all people. 1. Station 1- temperature a) Fill film canister ½ full with cold water. (ice water) b) Place ¼ tablet into film canister and cover with lid. c) How long in seconds does it take to pop off the top? Use a timer from time you insert Alka Seltzer until the top pops off. d) Record your results in the data table. e) Clean out film canister in sink, clean up any spills or splashes. f) Repeat Steps #1-4, except this time, use warm water. (tap water) g) Repeat Steps #1-4, except this time, use hot water. (on hot plate) 10

11 2. Station 2- Surface Area a) Fill film canister ¼ full with room temperature water. b) Place ¼ tablet into film canister and cover with lid. c) How long in seconds does it take to pop off the top? Use a timer from time you insert Alka Seltzer until the top pops off. Record your results in the data table. d) Clean out film canister in sink, clean up any spills or splashes e) Repeat steps # 1-4 except this time, break the ¼ piece tablet piece into 2 smaller pieces. (put both pieces in) f) Repeat steps #1-4 except this time, break the ¼ piece of tablet piece into 3-4 smaller pieces. (put all the pieces in) g) Repeat steps #1-4 except this time, crush the ¼ piece of tablet piece into a powder. (put powder on lid, flip & close) 3. Station 3- Amount of solvent(water) a) Fill film canister ¼ full with room temperature water. b) Place ¼ tablet into film canister and cover with lid. c) How long in seconds does it take to pop off the top? Use a timer from time you insert Alka Seltzer until the top pops off. Record your results in the data table. d) Clean out film canister in sink, clean up any spills or splashes e) Repeat Steps #1-4, except this time, fill the canister to ½ full. f) Repeat steps # 1-4, except this time, fill the canister to ¾ full. 4. Station 4- Agitation a) Fill film canister ¼ full with room temperature water. b) Place ¼ tablet into film canister and cover with lid. c) How long in seconds does it take to pop off the top? Use a timer from time you insert Alka Seltzer until the top pops off. Record your results in the data table. d) Clean out film canister in sink, clean up any spills or splashes e) Repeat Steps #1-4, except this time, shake canister for 1 second after placing lid on canister. f) Repeat steps # 1-4, except this time, shake the canister for 2 seconds after placing lid on canister. 11

12 Data/Observations: Station 1 Temperature Time in seconds cold water warm water Hot water Station 2 Surface Area Time (sec.) 1 piece (1/4 tablet) 2 pieces (1/2 the ¼) 3 small pieces Crushed pieces Station 3 (Amount of Solvent- Water)Concentration Time in seconds ¼ full ½ full ¾ full Station 4 Agitation Time in seconds No shake 1 sec shake 2 sec shake Analysis Q s 1. Describe what causes the lid to pop off the canister no matter what variable you are testing. 2. Identify the solute and the solvent used in this activity. 3. For EACH lab table/variable tested, state the variable you changed AND how it affected the variable you measured. (ex: For the Temperature station we changed & it caused to happen). 4. Why would you expect a packet of sugar to dissolve faster in hot tea than in iced tea? 12

13 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 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

14 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 14

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16 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) 16

17 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 17

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 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 18

19 Lab Activity: Endothermic/Exothermic PRE-LAB DISCUSSION All chemical reactions release or absorb energy. Chemical reactions that release energy in the form of heat are called exothermic reactions. Some chemical reactions absorb energy and are called endothermic reactions. PURPOSE After examining each reaction in the laboratory, you should be able to classify each reaction as exothermic or endothermic and determine if the heat of reaction( ) is negative or positive. SAFETY The chemicals can be toxic, please treat with respect and care. Rinse skin immediately upon contact. You MUST wear your goggles at all times. Be sure to rinse and dry your thermometer after each use. PROCEDURE In Part I, you will study the reaction between acetic acid and sodium bicarbonate (baking soda). An equation for the reaction is (Is it balanced? If not, try to balance the equation.): HC 2 H 3 O 2 + NaHCO 3 NaC 2 H 3 O 2 + H 2 O + CO 2 acetic acid + sodium bicarbonate sodium acetate + water + carbon dioxide gas Step 1: Add approximately 5 ml of acetic acid to a large test tube. Step 2: Record the temperature of the acetic acid. Step 3: Add a small scoop of sodium bicarbonate to the acetic acid. Step 4: Gently stir (with the thermometer) until all sodium bicarbonate has been dissolved. Observe the temperature, record any temperature change and other observations. Step 5: Discard the solution into the correct area and clean your equipment. In Part 2, you will study the reaction between hydrochloric acid and zinc metal. An equation for the reaction is (Is it balanced? If not, try to balance the equation.): HCl + Zn ZnCl 2 + H 2 hydrochloric acid + zinc metal zinc chloride + hydrogen gas Step 1: Add approximately 5 ml of hydrochloric acid to a large test tube. Step 2: Record the temperature of the hydrochloric acid. Step 3: Add a small piece of zinc to the hydrochloric acid. Step 4: Gently stir (with the thermometer) until all zinc has been dissolved. Observe the temperature, record any temperature change and other observations. Step 5: Discard the solution into the correct area and clean your equipment. 19

20 In Part 3, you will study the reaction between magnesium sulfate (Epsom salt) and water. An equation for the reaction is (Is it balanced? If not, try to balance the equation.): MgSO 4 + H 2 O MgSO 4 + H 2 + O 2 magnesium sulfate + water hydrogen gas + oxygen gas Step 1: Add approximately 10 ml of water to a large test tube. Step 2: Record the temperature of the water. Step 3: Add a small scoop of magnesium sulfate to the water. Step 4: Gently stir (with the thermometer) until all magnesium sulfate has been dissolved. Observe the temperature, record any temperature change and other observations. Step 5: Discard the solution into the correct area and clean your equipment. In Part 4, you will study the reaction between calcium chloride and water. An equation for the reaction is (Is it balanced? If not, try to balance the equation.): CaCl 2 + H 2 O CaCl 2 + H 2 + O 2 calcium chloride + water calcium chloride + oxygen gas Step 1: Add approximately 10 ml of water to a large test tube. Step 2: Record the temperature of the water Step 3: Add a small piece(s) of calcium chloride to the water. Step 4: Gently stir (with the thermometer) until all calcium chloride has been dissolved. Observe the temperature, record any temperature change and other observations. Step 5: Discard the solution into the correct area and clean your equipment. 20

21 Endothermic/Exothermic Lab OBSERVATIONS Part 1: Initial temperature of acetic acid Temperature after adding sodium bicarbonate to the acetic acid Did you feel a temperature change? What did you observe during the reaction? Part 2: Initial temperature of hydrochloric acid Temperature after adding zinc metal to the hydrochloric acid Did you feel a temperature change? What did you observe during the reaction? Part 3: Initial temperature of water Temperature after adding magnesium sulfate to the water Did you feel a temperature change? What did you observe during the reaction? Part 4: Initial temperature of water Temperature after adding calcium chloride to the water Did you feel a temperature change? What did you observe during the reaction? ANALYSIS Part 1: The reaction of sodium bicarbonate and acetic acid is exothermic or endothermic? What evidence do you have? Would it have a negative or positive value? 21

22 Part 2: The reaction of zinc metal and hydrochloric acid is exothermic or endothermic? What evidence do you have? Would it have a negative or positive value? Part 3: The reaction of magnesium sulfate and water is exothermic or endothermic? What evidence do you have? Would it have a negative or positive value? Part 4: The reaction of calcium chloride and water is exothermic or endothermic? What evidence do you have? Would it have a negative or positive value? Calcium chloride is as an ice-melting compound on sidewalks and city streets. Explain what is happening (use one of the terms-exothermic or endothermic in your explanation). Does the energy go from the surrounding to the chemicals or from the chemicals to the surroundings in an exothermic reaction? GOING FURTHER Define endothermic and exothermic in your own words. Classify each of the following as an exothermic or endothermic process. Melting ice cubes Burning a candle Evaporation of water Baking Bread Splitting a gas molecule apart Formation of snow in clouds Part 1: Balanced equation HC 2 H 3 O 2 + NaHCO 3 NaC 2 H 3 O 2 + H 2 O + CO 2 Part 2: Balanced equation HCl + Zn ZnCl 2 + H 2 Part 3: Balanced equation MgSO 4 + H 2 O MgSO 4 + H 2 + O 2 Part 4: Balanced equation CaCl 2 + H 2 O CaCl 2 + H 2 + O 2 22

23 Classwork 8-3: Use Table I to complete the chart below: 23

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25 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 = 25

26 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: 26

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32 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: 32

33 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 33

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37 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 37

38 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 38

39 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? 39

40 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? 40

41 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 41

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44 Complete the following table using the simulator. Activity - Reversible Reactions Simulation Determine whether you want to test an exo- or endothermic reaction. Change the E a and T, one variable at a time, keeping the other variables constant. Experiment Number Initial T Final T Type of reaction (endo/exo) E a Initial A Initial B Final A Final B K (equilibrium constant) 1. How does changing E a affect the reaction? Explain. 2. How do changes in T affect the value of K? 3. Is the reaction A B endothermic or exothermic? Explain. 4. How do you know when equilibrium is established? 5. Describe the conditions of a reversible reaction you conduct using the simulator. How do you know it is reversible? 44

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50 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) 50

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52 Unit Review: Kinetics and Equilibrium Place a checkmark next to each item that you can do! If a sample problem is given, complete it as evidence. 1. I can still do everything from Unit I can still do everything from Unit I can still do everything from Unit I can still do everything from Unit I can still do everything from Unit I can still do everything from Unit I can still do everything from Unit 7. Definition: effective collision 9. I can define effective collision and collision theory collision theory As the temperature, the reaction rate for most chemical reactions because there are effective collisions between particles. 10. I can state and apply the relationship between temperature and reaction rate in terms of collision theory. Given the reaction: 2Mg(s) + O 2 (g) -----> 2MgO(s) At which temperature would the reaction occur at the greatest rate? A) 0 o C B) 15 o C C) 95 o C D) 273K 52

53 As the surface area, the reaction rate 11. I can state and apply the relationship between surface area and reaction rate in terms of collision theory. because there are effective collisions between particles. At STP, which 4.0 g sample of Zn(s) will react most quickly with dilute hydrochloric acid? A) lump B) bar C) powdered D) sheet metal As the concentration, the reaction rate 12. I can state and apply the relationship between concentration and reaction rate in terms of collision theory. because there are effective collisions between particles. At 20 o C, a reaction between powdered Zn(s) and hydrochloric acid will occur most quickly if the concentration of the HCl is A) 1.0 M B) 1.5 M C) 2.5 M D) 2.8 M 13. I can state the unit used to measure energy. 14. Based on the location of the energy term, I can determine if the reaction is exothermic or endothermic. Energy is measured in. Given the following balanced equation: I + I -----> I kj Is this reaction exothermic or endothermic? Justify your answer. 15. I can use Table I to determine if a reaction is exothermic or endothermic. 53

54 16. I can define potential energy diagram, reaction coordinate, PE reactant, PE product, heat of reaction ( H), activation energy, catalyst. Definitions: potential energy diagram reaction coordinate PE reactant PE product heat of reaction ( H) activation energy catalyst entropy 17. Given a potential energy diagram, I can determine if the reaction is exothermic or endothermic. Give the potential energy diagram below, determine if the reaction is exothermic or endothermic. Justify your answer. 54

55 18. Given a potential energy diagram, I can determine the PE reactant, PE product, H, and activation energy. Given the potential energy diagram below, determine the PE reactant, PE product, H, and the activation energy. PE reactant = PE product = H = activation energy = 19. Given a potential energy diagram for an uncatalyzed reaction diagram, I can how the diagram will change when a catalyst is been added. Draw a dotted line on the potential energy diagram shown below to indicate how it will change if a catalyst is added. 20. I can rank the three phases of matter from least entropy to most entropy. 21. I can state the trends in nature for entropy and energy. Least entropy Most entropy < < In nature most systems in nature tend to undergo reactions that have a(n) in entropy and a(n) in energy. As Mrs. S says, nature is like a teenager ---- lazy and messy! 55

56 22. Given a balanced equation, I can determine if the reaction results in an overall increase or decrease in entropy. 23. I can define forward reaction, reverse reaction, reversible reaction, and closed system Definitions: forward reaction reverse reaction reversible reaction closed system 24. I can state the three types of equilibrium. The three types of equilibrium are: equilibrium equilibrium and equilibrium 56

57 25. I can state two conditions that apply to all systems at equilibrium. 26. Given a list of reactions, I can identify reactions that show equilibrium (chemical, phase, or solution). In a system at equilibrium the of the forward and reverse reaction must be and the of the reactants and products must be. Which balanced equation represents phase equilibrium? A) H 2 (g) + I 2 (s) <-----> 2HI(g) B) I 2 (s) <-----> I 2 (g) H 2 O C) KCl(s) <-----> KCl(aq) D) 2KCl(s) + 3O 2 (g) -----> 2KClO 3 Which balanced equation represents solution equilibrium? A) H 2 (g) + I 2 (s) <-----> 2HI(g) B) I 2 (s) <-----> I 2 (g) H 2 O C) KCl(s) <-----> KCl(aq) D) 2KCl(s) + 3O 2 (g) -----> 2KClO 3 Which balanced equation represents chemical equilibrium? A) H 2 (g) + I 2 (s) <-----> 2HI(g) B) I 2 (s) <-----> I 2 (g) H 2 O C) KCl(s) <-----> KCl(aq) D) 2KCl(s) + 3O 2 (g) -----> 2KClO In terms of saturation, I can describe a solution that is at equilibrium. 28. I can state LeChatelier s Principle. In terms of saturation, a solution that is at equilibrium must be. LeChatelier s Principle states 57

58 29. Given a balanced equation at equilibrium, I can predict the direction of shift in the equilibrium when the temperature, concentration, or pressure is changed or if a catalyst is added. Given the reaction at equilibrium: 2SO 2 (g) + O 2 (g) <-----> 2SO 3 (g) + 392kJ Predict the direction of shift in the equilibrium (right, left, no shift) when the following changes are made to the system. Change Direction of Shift Increase concentration of SO 2 Increase concentration of SO 3 Increase temperature Increase pressure Add a catalyst 58

59 Practice: Kinetics and Equilibrium 1. Given the potential energy diagram for a reaction: Which interval on this diagram represents the difference between the potential energy of the products and the potential energy of the reactants? A) 1 B) 2 C) 3 D) 4 2. Given the potential energy diagram for a chemical reaction: 4. Which information about a chemical reaction is provided by a potential energy diagram? A) the oxidation states of the reactants and products B) the average kinetic energy of the reactants and products C) the change in solubility of the reacting substances D) the energy released or absorbed during the reaction 5. Given the potential energy diagram of a chemical reaction: Which statement correctly describes the energy changes that occur in the forward reaction? A) The activation energy is 10. kj and the reaction is endothermic. B) The activation energy is 10. kj and the reaction is exothermic. C) The activation energy is 50. kj and the reaction is endothermic. D) The activation energy is 50. kj and the reaction is exothermic. 3. A solution that is at equilibrium must be A) concentrated B) dilute C) saturated D) unsaturated 59 Which arrow represents the potential energy of the reactants? A) A B) B C) C D) D 6. The activation energy required for a chemical reaction can be decreased by A) increasing the surface area of the reactant B) increasing the temperature of the reactant C) adding a catalyst to the reaction D) adding more reactant 7. Given the reaction: 2 H2(g) + O2(g) 2 H2O( ) kj What is the approximate H for the formation of 1 mole of H2O( )? A) kj B) kj C) kj D) kj 8. According to Reference Table I, which statement best describes the formation of HI(g)? A) It is exothermic, and heat is released. B) It is exothermic, and heat is absorbed. C) It is endothermic, and heat is released. D) It is endothermic, and heat is absorbed.

60 Kinetics and Equilibrium Review 9. The diagram below represents the energy changes that occur during the formation of a certain compound under standard conditions. 13.A potential energy diagram is shown below. According to Reference Table I, the compound could be A) C2H6(g) B) CO2(g) C) HI(g) D) NH3(g) 10.The potential energy diagram of a chemical reaction is shown below. Which reaction would have the lowest activation energy? A) the forward catalyzed reaction B) the forward uncatalyzed reaction C) the reverse catalyzed reaction D) the reverse uncatalyzed reaction 14.Given the potential energy diagram: Which arrow represents the part of the reaction most likely to be affected by the addition of a catalyst? A) A B) B C) C D) D 11. Which statement best describes a chemical reaction in which energy is released? A) It is exothermic and has a negative H. B) It is exothermic and has a positive H. C) It is endothermic and has a negative H. D) It is endothermic and has a positive H. 12. Based on Reference Table I, which reaction is endothermic? A) NaOH(s) Na + (aq) + OH (aq) B) NH4Cl(s) NH4 + (aq) + Cl (aq) C) CO(g) + O2(g) CO2(g) D) CH4(g) + 2 O2(g) CO2(g) + 2 H2O( ) 60 With reference to energy, the reaction A + B AB can best be described as A) endothermic, having a + H B) endothermic, having a H C) exothermic, having a + H D) exothermic, having a H 15. The heat energy absorbed or released during the formation of products is equal to A) B) C) D)

61 Kinetics and Equilibrium Review 16. Base your answer on the potential energy diagram below. The potential energy of the activated complex is equal to the sum of A) X + Y B) X + W C) X + Y + W D) X + W + Z 17. According to Table I, which compound has a higher potential energy than the elements from which it is formed? A) aluminum oxide (s) B) hydrogen oxide ( ) C) carbon dioxide (g) D) nitrogen (II) oxide (g) 18. Given the reaction: What is the heat of formation of nitrogen (II) oxide in kj/mole? A) B) C) D) 19. The reaction A(g) + B(g) C(g) + D(g) + 30 kj has a forward activation energy of 20 kj. What is the activation energy for the reverse reaction? A) 10 kj B) 20 kj C) 30 kj D) 50 kj 20. Assume that the potential energy of the products in a chemical reaction is 60 kilojoules. This reaction would be exothermic if the potential energy of the reactants were A) 50 kj B) 20 kj C) 30 kj D) 80 kj Which of the following best describes exothermic chemical reactions? A) They never release heat. B) They always release heat. C) They never occur spontaneously. D) They always occur spontaneously. 22. For a given reaction, adding a catalyst increases the rate of the reaction by A) providing an alternate reaction pathway that has a higher activation energy B) providing an alternate reaction pathway that has a lower activation energy C) using the same reaction pathway and increasing the activation energy D) using the same reaction pathway and decreasing the activation energy 23. Given the reaction: A + B C + D The reaction will most likely occur at the greatest rate if A and B represent A) nonpolar molecular compounds in the solid phase B) ionic compounds in the solid phase C) solutions of nonpolar molecular compounds D) solutions of ionic compounds 24. Beaker A contains a 1 gram piece of zinc and beaker B contains 1 gram of powdered zinc. If 100 milliliters of 0.1 M HCl is added to each of the beakers, how does the rate of reaction in beaker A compare to the rate of reaction in beaker B? A) The rate in A is greater due to the smaller surface area of the zinc. B) The rate in A is greater due to the larger surface area of the zinc. C) The rate in B is greater due to the smaller surface area of the zinc. D) The rate in B is greater due to the larger surface area of the zinc. 25. If the pressure on gaseous reactants is increased, the rate of reaction is increased because there is an increase in the A) activation energy B) volume C) concentration D) heat of reaction

62 Kinetics and Equilibrium Review 26. Base your answer to the following question on the table below, which represents the production of 50 milliliters of CO2 in the reaction of HCl with NaHCO3. Five trials were performed under different conditions as shown. (The same mass of NaHCO3 was used in each trial.) 31. A piece of Mg(s) ribbon is held in a Bunsen burner flame and begins to burn according to the equation: 2Mg(s) + O2 (g) 2MgO(s). Which trial would produce the fastest reaction? A) trial A B) trial B C) trial C D) trial D 27. Which statement explains why the speed of some chemical reactions is increased when the surface area of the reactant is increased? A) This change increases the density of the reactant particles. B) This change increases the concentration of the reactant. C) This change exposes more reactant particles to a possible collision. D) This change alters the electrical conductivity of the reactant particles. 28. As the number of moles per liter of a reactant in a chemical reaction increases, the number of collisions between the reacting particles A) decreases B) increases C) remains the same 29. Which conditions will increase the rate of a chemical reaction? A) decreased temperature and decreased concentration of reactants B) decreased temperature and increased concentration of reactants C) increased temperature and decreased concentration of reactants D) increased temperature and increased concentration of reactants 30. An increase in the temperature of a system at equilibrium favors the A) endothermic reaction and decreases its rate B) endothermic reaction and increases its rate C) exothermic reaction and decreases its rate D) exothermic reaction and increases its rate 62 The reaction begins because the reactants A) are activated by heat from the Bunsen burner flame B) are activated by heat from the burning magnesium C) underwent an increase in entropy D) underwent a decrease in entropy 32. Two reactant particles collide with proper orientation. The collision will be effective if the particles have A) high activation energy B) high ionization energy C) sufficient kinetic energy D) sufficient potential energy 33. Given the equation representing a phase change at equilibrium: C2H5OH( ) «C2H5OH(g) Which statement is true? A) The forward process proceeds faster than the reverse process. B) The reverse process proceeds faster than the forward process. C) The forward and reverse processes proceed at the same rate. D) The forward and reverse processes both stop. 34. Which statement must be true when solution equilibrium occurs? A) The solution is at STP. B) The solution is supersaturated. C) The concentration of the solution remains constant. D) The masses of the dissolved solute and the undissolved solute are equal.

63 Kinetics and Equilibrium Review 35. When AgNO3(aq) is mixed with NaCl(aq), a reaction occurs which tends to go to completion and not reach equilibrium because A) a gas is formed B) water is formed C) a weak acid is formed D) a precipitate is formed 36. The vapor pressure of a liquid at a given temperature is measured when the rate of evaporation of the liquid becomes A) less than the rate of condensation B) greater than the rate of condensation C) equal to the rate of condensation D) equal to a zero rate of condensation 37. A liquid in a stoppered flask is allowed to stand at constant temperature until the liquid level in the flask remains constant. Which condition then exists in the flask? A) Only liquid is evaporating. B) Only vapor is condensing. C) The rate of condensation is greater than the rate of evaporation. D) The rate of condensation is equal to the rate of evaporation. 38. The diagram below shows a bottle containing NH 3(g) dissolved in water. 40. Given the reaction: AgI(s) «Ag + (aq) + I (aq) Solution equilibrium is reached in the system when A) dissolving stops occurring B) crystallization stops occurring C) both dissolving and crystallization stops occurring D) dissolving occurs at the same rate that crystallization occurs 41. Ammonia is produced commercially by the Haber reaction: N2(g) + 3 H2(g) «2 NH3(g) + heat The formation of ammonia is favored by A) an increase in pressure B) a decrease in pressure C) removal of N2(g) D) removal of H2(g) 42. Given the reaction at equilibrium: 4 HCl(g) + O2(g) «2 Cl2(g) + 2 H2O(g) If the pressure on the system is increased, the concentration of Cl2(g) will A) decrease B) increase C) remain the same 43. Given the reaction at equilibrium: 2 SO2(g) + O2(g) «2 SO3(g) How can the equilibrium, NH3(g) NH3(aq), be reached? A) Add more water. B) Add more NH3(g). C) Cool the contents. D) Stopper the bottle. 39. In terms of energy and entropy, systems in nature tend to undergo changes toward A) higher energy and higher entropy B) higher energy and lower entropy C) lower energy and higher entropy D) lower energy and lower entropy 63 If the temperature remains constant, an increase in pressure will A) have no effect on the equilibrium B) shift the equilibrium to the right C) shift the equilibrium to the left D) change the value of the equilibrium constant 44. Which of these changes produces the greatest increase in entropy? A) CaCO3(s) CaO(s) + CO2(g) B) 2 Mg(s) + O2(g) 2 MgO(s) C) H2O(g) H2O( ) D) CO2(g) CO2(s)

64 Kinetics and Equilibrium Review 45. Given the reaction at equilibrium: 50. Given the system at equilibrium: N2(g) + O2(g) = 2 NO(g) If the temperature remains constant and the pressure increases, the number of moles of NO(g) will A) decrease B) increase C) remain the same 46. Given the reaction at equilibrium: A(g) + B(g) C(g) + D(g) The addition of a catalyst will A) shift the equilibrium to the right B) shift the equilibrium to the left C) increase the rate of forward and reverse reactions equally D) have no effect on the forward or reverse reactions 47. Given the reaction at equilibrium: 2 A(g) + 3 B(g) «A2B3(g) + heat Which change will not affect the equilibrium concentrations of A(g), B(g), and A2B3(g)? A) adding more A(g) B) adding a catalyst C) increasing the temperature D) increasing the pressure 48. Given the equation representing a reaction at equilibrium: Which change causes the equilibrium to shift to the right? A) decreasing the concentration of B) decreasing the pressure C) increasing the concentration of D) increasing the temperature 49. Given the reaction at equilibrium: N2(g) + O2(g) + energy 2 NO(g) Which change will result in a decrease in the amount of NO(g) formed? A) decreasing the pressure B) decreasing the concentration of N2(g) C) increasing the concentration of O2(g) D) increasing the temperature 64 Which changes occur when this system? is added to A) The equilibrium shifts to the right and the concentration of increases. B) The equilibrium shifts to the right and the concentration of decreases. C) The equilibrium shifts to the left and the concentration of increases. D) The equilibrium shifts to the left and the concentration of decreases. 51. Given the system at equilibrium: N2O4(g) kj 2 NO2(g) What will be the result of an increase in temperature at constant pressure? A) The equilibrium will shift to the left, and the concentration of NO2(g) will decrease. B) The equilibrium will shift to the left, and the concentration of NO2(g) will increase. C) The equilibrium will shift to the right, and the concentration of NO2(g) will decrease. D) The equilibrium will shift to the right, and the concentration of NO2(g) will increase. 52. Given the reaction for the Haber process: N2 + 3 H2 «2 NH3 + heat The temperature of the reaction is raised in order to A) increase the percent yield of nitrogen B) increase the rate of formation of ammonia C) affect the forward reaction rate most D) affect the reverse reaction rate least 53. Given the equilibrium system: 2 A(g) + B(g) + 10 kcal «C(g) Which conditions would yield the most product? A) low temperature and high pressure B) low temperature and low pressure C) high temperature and high pressure D) high temperature and low pressure

65 Kinetics and Equilibrium Review 54. Which phrase best describes the reaction below? A) exothermic with an increase in entropy B) exothermic with a decrease in entropy C) endothermic with an increase in entropy D) endothermic with a decrease in entropy 55. Which potential energy diagram indicates a reaction can occur spontaneously? A) B) C) D) 65