Unit 9b: Equilibrium, Enthalpy, and Entropy

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Unit 9b: Equilibrium, Enthalpy, and Entropy

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Unit 9b Vocabulary: 1. Activated Complex: The species that are formed and decomposed during the mechanism, and is also called the intermediate. 2. Activation Energy: The energy that must be added to allow the reactants to complete the reaction and form the activated complex. 3. Catalyst: A chemical that is added to a reaction to eliminate steps in the mechanism and increase the reaction rate and decrease the activation energy without itself being consumed by the reaction. 4. Effective Collision: A collision between reactant particles that results in a chemical reaction taking place. 5. Enthalpy: The total amount of potential energy stored in a substance. 6. Endothermic: A reaction that absorbs and stores energy from the surrounding environment. 7. Entropy: A system’s state of disorder. Entropy increases as temperature increases. Entropy increases as a substance goes from solid to liquid to gas. 8. Equilibrium: A system where the rate of forward change is equal to the rate of reverse change. At equilibrium there is no net change. 9. Exothermic: A reaction that releases stored energy into the surrounding environment. 10. Favored: A change in a thermodynamic property that contributes towards the reaction being spontaneous. 11. Free Energy: The total amount of energy available in a system to do work. Free Energy is a combination of both enthalpy and entropy. 12. Heat of Reaction: The net gain or loss of potential energy during a chemical reaction. 13. Inhibitor: A chemical that is added to a reaction to add steps to the mechanism to decrease the reaction rate and increase the activation energy without itself being consumed by the reaction. 14. Kinetics: The study of reaction mechanisms and reaction rates. 15. Nonspontaneous: A reaction that requires a constant input of energy to occur, or the reaction will reverse or stop. 16. Reaction Rate: The amount of reactant consumed in a given unit of time. Website upload-Unit 9b: Equilibrium Systems

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

17. Spontaneous: A reaction that continues independently once started. 18. Thermodynamics: The study of heat flow during physical and chemical changes. 19. Unfavored: A change in a thermodynamic property that contributes towards the reaction being nonspontaneous.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Unit 9b Homework Assignments: Assignment:

Date:

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium

Objective: What is the role equilibrium has in chemistry? Equilibrium: Equilibrium is a continuous state of the rate of balance between two opposing changes. In a state of equilibrium the rate of the forward change is equal to the rate of the reverse change. Most chemical reactions are reversible: A + B  C + D + energy C + D + energy  A + B A + B  (±energy)  C + D

= forward reaction When the rate of the = reverse reaction forward reaction equals Double arrows () the rate of the reverse indicate that BOTH reaction, a state of reactions are occurring equilibrium is reached. at the same time.

If you ride up a moving escalator, you are moving at the rate that the escalator is moving upwards. However, if you turn around and start to walk DOWN the up escalator, and you match the escalator’s rate (up) but in the opposite direction (down), to someone watching you it looks as if you are not moving. However, you are still expending energy trying to go to the bottom, and the escalator is expending energy trying to carry you uphill. If anything was to upset the process (power failure to the escalator; you trip and fall, etc.), the equilibrium would be upset and you would either make it to the bottom or ride to the top. Website upload-Unit 9b: Equilibrium Systems

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium examples: Haber Process for ammonia gas: Forward reaction:

N2(g) + 3 H2(g)  2 NH3(g) + 92 kJ (exo)

Reverse reaction:

2 NH3(g) + 92 kJ  N2(g) + 3 H2(g) (endo)

Equilibrium:

N2(g) + 3 H2(g)  2 NH3(g) + 92 kJ

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Properties

Objective: What properties of equilibrium will systems have?

Properties of Systems at Equilibrium: 1. Equilibrium is a dynamic state; think of equilibrium as a continuous pathway, never achieving a ‘set’ endpoint. Particles of reactants are reacting and forming products at the same rate that products are decomposing back into the reactants they came from. Remember that the system is in continuous motion, though it may look like the reaction is stagnant. 2. Equilibrium can only be maintained in a closed system. A closed system neither gains nor loses anything. This includes energy (loss or gain), adding reactants, or the removal of products. 3. As long as the system is closed, a system at equilibrium will remain that way forever. Changing ANY condition of equilibrium will alter the balance of the entire equilibrium (see pgs. 33-40). 4. Equilibrium occurs at different concentrations of product and reactant. Depending on the nature of the species involved, assuming we start with the forward reaction, the rate of the reverse reaction will increase as the product is formed during the forward reaction. When the forward AND reverse reaction rates are equal, equilibrium is achieved. This may occur at different concentrations of product and reactant. Website upload-Unit 9b: Equilibrium Systems

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Diagram:  Equilibrium may be reached ANYWHERE along a line that starts at 0% and ends at 100%.  At any point along the line the percentage of the reaction going forward (reactants) ADDED to the percentage of the reaction going backward (products) equals 100%. (% forward reaction) + (% reverse reaction) = 100%

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Three Types of Equilibrium Objective: What forms of equilibrium are possible in chemistry? 1. Chemical Equilibrium: i. If the rate of the forward reaction is equal to the rate of the reverse reaction the reaction has achieved chemical equilibrium. You have seen the Haber Process for the production of ammonia:

N2(g) + 3 H2(g)  2 NH3(g) + 92 kJ This reaction produces ammonia (and heat), but some of the ammonia, NH3(g), produced will decompose during the reaction back into reactants, N2(g) and 3 H2(g). ii. When the rate of synthesis (forward reaction) equals the rate of decomposition (reverse reaction), and no other changes occur, this system will be at equilibrium. iii. As stated before, changing ANY component of the system will change the equilibrium of the entire system.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

2. Solution Equilibrium: i. If a solution becomes saturated, the rate of dissolving equals the rate of precipitation, and the reaction has achieved solution equilibrium.

NaCl(s)  Na+1(aq) + Cl-1(aq) ii. When sodium chloride is first placed into pure water, the solid ionic crystals dissolve. As the concentration of the dissolved ions increases, some of those dissolved Na+1(aq) and Cl-1(aq) ions will temporarily rejoin to form a soluble precipitate which almost immediately dissolves again. Eventually all the ions will be held apart by the polar water molecules, and no more solid may enter the solution until some ions come out of solution as precipitate. At this point the rate of dissolving equals the rate of precipitation, and you have a SATURATED solution. Additional added solid would not dissolve, or only as a temporary supersaturated solution. Solution Equilibriums

Unsaturated - solute almost all Close to Saturation - solute undissolved; reaction almost all almost all dissolved; reaction forward mostly forward, some reverse Website upload-Unit 9b: Equilibrium Systems

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Saturated-dissolving rate is equal to precipitate formation rate; no net change Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

3. Physical Equilibrium: i. If the rate of a forward phase change is equal to the rate of a reverse phase change, then the system is in Physical (or Phase) Equilibrium. ii. Physical equilibrium occurs AT the phase change temperature. Remember that during a phase change, all energy input is going towards increasing the potential energy of the substance, as there is no increase in average kinetic energy (temperature) at the phase change temperature. For water, the boiling (vaporization point) at 1 atm is 373 K. This means if water is maintained in a sealed container at 1 atm and 373 K, for each water molecule that changes from liquid to gaseous, another water molecule will change from gaseous to liquid.

Liquid-Gaseous Equilibrium for Water at 1 atm and 373 K

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Triple Point Temperature: An interesting phenomenon of Phase Equilibrium is the concept of the Triple Point temperature. At a substance’s Triple Point, that substance can exist in THREE phases simultaneously. Note that the Triple Point is also a function of pressure.

For the diagram above, note that water is different from most materials in that as you increase the pressure at the melting-freezing point, the freezing point of water decreases with increasing pressure. This is one of the reasons why a sealed can or bottle of a carbonated beverage can have the liquid contents BELOW normal freezing temperature of water, but as soon as the container is opened to the atmosphere (pressure drops rapidly), the liquid may suddenly partially freeze, due to the fact that the pressure drops much more rapidly than the temperature can increase.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Systems Practice Regents Problems: (ungraded) 1. Which statement about a system at equilibrium is true? a) The forward reaction rate is less than the reverse reaction rate. b) The forward reaction rate is greater than the reverse reaction rate. c) The forward reaction rate is equal to the reverse reaction rate. d) The forward reactions stop and the reverse reactions continue. 2. Given the reaction in water of AgCl(s)  Ag+1(aq) + Cl-1(aq), once equilibrium is reached, which statement is accurate? a) The AgCl(s) will be completely consumed. b) The rates of the forward and reverse reactions are equal. c) The entropy of the forward reaction will continue to decrease. d) The concentration of Ag+1(aq) is greater than the concentration of Cl-1(aq). 3. Which type(s) of change, if any, may reach equilibrium? a) A physical change, only. b) A chemical change, only. c) Neither a chemical nor a physical change. d) Both a chemical and a physical change. 4. In a reversible reaction, a chemical equilibrium is attained when the a) Concentration of the reactants reaches zero. b) Concentration of the products remains constant. ONLY if ONE constant! c) Rate of the forward reaction is greater than the rate of the reverse reaction. d) Rate of the reverse reaction is greater than the rate of the forward reaction. 5. Given the reaction of H2O(s)  H2O(l), at which temperature will equilibrium exist when the atmospheric pressure is equal to 101.3 kPa? a) 0 K c) 273 K b) 100 K d) 373 K 6. The temperature at which solid and liquid phases of the same type of matter exist in equilibrium is called its a) Boiling point c) Melting point b) Heat of fusion d) Heat of vaporization Website upload-Unit 9b: Equilibrium Systems

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Student name: _________Key__________

Class Period: _3, 5, & 10_

Please carefully remove this page from your packet to hand in. Equilibrium Systems homework 1. Which of the following need to be equal at equilibrium? a) The masses of both products and reactants. b) The volumes of both products and reactants. c) The concentrations of both the products and reactants. d) The rates of formation of both the products and the reactants. 2. A stoppered (sealed) flask contains 20.0 grams of liquid water and 20.0 grams of water vapor. Does a state of equilibrium of water exist in the flask? a) Yes, the bottle is stoppered. b) Yes, the amount of each component is equal. c) Yes, but only if the rates of evaporation and precipitation are equal. d) Yes, but only if a) and b) are both true. e) Yes, but only if a) and c) are both true. f) Yes, but only if a), b), and c) are all true. 3. A mixture of 50.0 g of water ice and 100.0 g of liquid water is massed and then kept at a steady 0.0˚C in a closed container. After one hour, you mass the contents of the container again. What would you predict the resulting masses to be? a) The mass of the ice and the mass of the liquid will remain constant. b) The mass of the ice decreased and the mass of the water increased as the ice melted. c) The mass of the ice increased and the mass of the water decreased as the water froze. 4. The same mixture in question #3 above of 50.0 g of water ice and 100.0 g of liquid water is massed and then heated to 3.98˚C, the point of maximum density for water. After one hour, you mass the contents of the container again. What would you predict the resulting masses to be? a) The mass of the ice and the mass of the liquid will remain constant. b) The mass of the ice decreased and the mass of the water increased as the ice melted. c) The mass of the ice increased and the mass of the water decreased as the water froze. Website upload-Unit 9b: Equilibrium Systems

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Use the information below to answer questions #5 and #6. When sodium chloride is first added to distilled water, it looks like the solid crystals disappear into the water. Continue to add crystals, and the rate that they disappear slows down until eventually you achieve a condition where any added solid sinks to the bottom of the container. 5. When the added sodium chloride crystals no longer dissolves in the water and sinks to the bottom of the container, what type of a solution do you have? a) Moist b) Saturated c) Unsaturated d) Supersaturated 6. In the same conditions as listed above, while the solid crystals sit on the bottom of the container, what will happen to the crystals in the water? a) They’ll stay exactly the same. b) They will change size and only get larger over time. c) They will change size and only get smaller over time. d) They will change size, but the total mass of the solids will be constant. 7. The vapor pressure of a liquid at a given average kinetic energy in a sealed system is measured when the rate of evaporation of the liquid is a) Less than the rate of condensation. b) Equal to the rate of condensation. c) Equal to a zero rate of condensation. d) Greater than the rate of condensation. 8. A system is said to be in a state of dynamic equilibrium when the a) Concentration of products is the same as the concentration of reactants. b) Concentration of products is greater than the concentration of reactants. c) Rate at which products are formed is the same as the rate at which reactants are formed. d) Rate at which products are formed is greater than the rate at which reactants are formed. 9. A solution that is at equilibrium must be a) Dilute b) Saturated

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c) Unsaturated d) Concentrated

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Reaction Progress

Objective: How do we know if a reaction will continue when started? Spontaneous Chemical Reactions:  A reaction that continues without additional input once it has initiated is called a Spontaneous Reaction. 1. Spontaneous reactions are very useful in the industrial world, and also in chemistry. 2. Once you start a spontaneous reaction it will go to completion until either the reactants are consumed or it enters a state of equilibrium if the products are not removed. 3. Spontaneous reactions require a balance between two factors: i. Enthalpy- the available (potential) energy in a substance ii. Entropy-the amount of randomness (disorder) in a system Expanding Universe!

Falling is easier!

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Enthalpy

Objective: How do we express the potential energy in a system? Enthalpy: i. Enthalpy is the heat content (potential energy) of a system. ii. Nature favors reactions that undergo a decrease in enthalpy. iii. Let go of a ball in your hand; it falls. iv. Falling is a spontaneous (no additional energy) decrease in enthalpy (potential energy). The ball can’t fall again from its starting height.  A decrease in potential energy is favored in nature, so exothermic reactions are the most favored (and most common - see Table I). v. Most exothermic (decreasing enthalpy) reactions are spontaneous, and complete once started. (Think of a bonfire; it burns as long as it has fuel.) vi. Conversely, most endothermic reactions are nonspontaneous, and require constant input of energy to keep going. (Think of ice; if you keep your ice in the freezer, you prevent outside energy from getting to it.)

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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(-) sign is exothermic; gives off heat; Red = fire (feels hot!)

(+) sign is endothermic; absorbs heat; Blue = ice (feels cold!)

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Enthalpy

Objective: How do we express the potential energy in a system? Enthalpy in common situations:  Imagine you are on a bike at the top of a hill. If you push off and pick your feet up, and you could coast down the hill without any additional energy (well, you should steer!) After the initial ‘push’ (EA), the reaction (you and the bike) are on a spontaneous exothermic change.  Now, you are on the bottom of the hill and need to ride back to the top. You have the same initial ‘push’ (EA), but you need to expend almost continuous energy to climb the hill. If you stop pedaling, the reaction (you and the bike) will stop, and probably reverse itself (roll down the hill.) This is a nonspontaneous endothermic change.

Watch Crash Course Chemistry Enthalpy video - 11:23

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Entropy

Objective: How do we express the amount of order in a system? Entropy: The randomness (disorder) of a system is called Entropy. Nature favors reactions that increase entropy.  As a substance increases in temperature, the substance undergoes an increase in entropy as well. As each subsequent phase change occurs, the randomness (disorder) of the particles of that substance increases.  In order of LEAST to MOST entropy, the phases are: solid  liquid gas. Solids are locked in a lattice, and gases have very random movement controlled only by the confines of their container. Liquids fall in between.  As nature favors entropy, nature favors increases in phase.

Kitchen

Freezer

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Reactants are a solid(s) AND a gas(g); product is a solid(s) – Entropy DECREASED, and is UNFAVORED

Reactants are a solid(s) AND a gas(g); product is a gas(g) – Entropy INCREASED, and is FAVORED

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Regents Chemistry ’14-‘15 Mr. Murdoch

Spontaneous Reactions

Objective: What factors will drive a reaction to complete on its own? Spontaneous Reactions:  For a reaction to occur spontaneously, both Enthalpy AND Entropy will be considered. i. Nature favors DECREASING enthalpy (exothermic processes); ii. Nature favors INCREASING entropy (phase change to less order)  If both enthalpy and entropy are favored, then the reaction will be spontaneous.

1. Favored Reactions: A favored reaction will be spontaneous at all temperatures.

i. This reaction has a ∆H of -84.0 kJ/mole of C2H6(g) produced, enthalpy decreases, is exothermic, and is favored. ii. This reaction starts with a solid and a gas, and ends with only gas. Entropy increases, and is favored. iii. Both enthalpy and entropy are favored, and this reaction will ALWAYS be spontaneous at any temperature. iv. No additional energy after EA will be needed to complete the reaction. Website upload-Unit 9b: Equilibrium Systems

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Nonspontaneous Reactions Objective: What factors will drive a reaction to stop on its own? 2. Unfavored Reactions: An unfavored reaction will be nonspontaneous at all temperatures. An unfavored reaction will require constant energy input to complete.

i. This reaction has a ∆H of +33.2 kJ/mole of NO2(g) produced, enthalpy increases, is endothermic, and is unfavored. ii. This reaction starts with a total of three moles of gaseous reactants, and ends with only two moles of gaseous product. While mass is conserved, the number of particles has decreased, or entropy decreased, which is against nature, and therefore unfavored. iii. Both enthalpy and entropy are unfavored, and this reaction will ALWAYS be nonspontaneous at any temperature. iv. After EA is input, continuous energy will be required to maintain this reaction.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

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Topic: Partially Spontaneous Rx’s Objective: What combinations of factors will drive a reaction? 3. Lower-Temperature favored Reactions: If enthalpy is favored, but entropy is unfavored, the reaction will be spontaneous at lower temperatures.

CO2(g) has a sublimation/deposition temperature of near 195 K, meaning below that it is in the solid phase. i. This reaction has a ∆H of -283.0 kJ/mole of CO2(g) produced, so this reaction is exothermic, which is favored. ii. This reaction starts with a total of three moles of gas, and ends with only two moles of gas. Entropy decreases, and is unfavored. Below 195 K entropy decreases even more (becomes solid), further unfavoring the reaction. iii. This reaction will be spontaneous only at temperatures when the product is in the gaseous phase. Below 195 K the entropy decreases more, and the reaction will be nonspontaneous. Additional energy after EA will be needed to complete the reaction.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Partially Spontaneous Rx’s Objective: What combinations of factors will drive a reaction? 4. Higher-Temperature favored Reactions: If enthalpy is unfavored, but entropy is favored, the reaction will be spontaneous at higher temperatures. H2O

i. This reaction has a ∆H of +25.69 kJ/mole of NH4NO3(s) when decomposed in water, so this reaction is endothermic, which is unfavored. ii. This reaction starts out as a solid, but dissolves in aqueous solution. In an aqueous solution, ions may move freely, and entropy increases, which is favored. iii. This reaction will be spontaneous only at temperatures when the product is in the aqueous phase. Below about 273 K the entropy decreases more (becomes solid), and the reaction will be nonspontaneous. Additional energy after EA will be needed to complete the reaction, mostly to keep the water from freezing.

Watch Bozeman Science Spontaneous Processes video - 7:42

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Spontaneity of Water

Objective: Why does water ice melt above 273 K? 5. Spontaneity of the melting of water ice: The reaction for the melting of water ice to liquid water is: H2O(s)  H2O(l)

+ 6.01 kJ (334 J/g x 18.0 g/mole)

i. This process has a +∆H (the Heat of Fusion of water), is endothermic, and enthalpy increases, which is unfavored. ii. This process starts with a solid, and ends with a liquid, which is a phase change to less order, so entropy increases, which is favored. iii. The combination of increasing enthalpy and increasing entropy makes for a process that is spontaneous at higher temperatures. Below 273 K the process is nonspontaneous, but above 273 K the process will maintain without any additional energy. Once any ice melts, the kinetic energy in the water is greater than the kinetic energy in ice, and melting will continue. *Note: This is DIFFERENT than water right at the freezing point of 273 K/ 0.0°C. Water AT a temperature of 273 K/ 0.0°C would equally freeze/melt if all other factors are kept the same.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Enthalpy & Entropy Practice Regents Problems: (ungraded) 1. According to Reference Table I, which reaction below spontaneously forms a compound from its reactants? a) H2(g) + I2(g)  2 HI(g) c) 2 H2(g) + O2(g)  2 H2O(g) b) N2(g) + O2(g)  2 NO(g) d) N2(g) + 2 O2(g)  2 NO2(g) 2. Which change below is exothermic? a) Melting of iron b) Freezing of water

c) Sublimation of iodine d) Vaporization of ethanol

3. Which reaction below has the greatest increase in entropy? a) H2O(g)  H2O(l) c) 2 H2O(g)  2 H2(g) + O2(g) b) H2O(l)  H2O(s) d) 2 H2O(l)  2 H2(g) + O2(g) 4. According to Reference Table I, which compound decreases in enthalpy as it dissolves? a) NaCl c) KNO3 d) NH4NO3 b) LiBr

Given the reaction: 2 Na(s) + Cl2(g)  2 NaCl(s) 5. As the reactants form products, the entropy of the chemical system will a) Increase c) Remain the same b) Decrease 6. Which chemical reaction will always be spontaneous? a) An exothermic reaction in which entropy increases. b) An exothermic reaction in which entropy decreases. c) An endothermic reaction in which entropy increases. d) And endothermic reaction in which entropy decreases.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Student name: ________Key________ Class Period: _3, 5, & 10_ Please carefully remove this page from your packet to hand in. Enthalpy, Entropy, and Reaction Spontaneity Homework: For the given reactions below, state if entropy increases or decreases, and if the change is favored or unfavored. 1 pt. ea. Reaction

Entropy: Inc or Dec? Change: Fav or Unfav?

CO2(s)  CO2(g)

I

F

I2(g)  I2(s)

D

U

1 C(s) +1 O2(g)  1 CO2(g)

D

U

4 Al(s) + 3 O2(g)  2 Al2O3(s)

D

U

2 CO(g) + O2(g)  2 CO2(g)

D

U

2 H2(g) + O2(g)  2 H2O(l)

D

U

For the given reactions below, state if the reactions are spontaneous, nonspontaneous, or at equilibrium, and also state whether enthalpy and entropy are favored or unfavored. 1 pt. ea.

Spont, Nonspont, Enthalpy & entropy Fav or Unfav? or at Equil

Reaction 4 Al(s) + 3 O2(g)  2 Al2O3(s) + 3351 kJ

E

Enthalpy: __F__ Entropy: __U___

2 CO(g) + O2(g)  2 CO2(g) + 566 kJ

E

Enthalpy: __F__ Entropy: __U__

NaOH(s)  Na+1(aq) + OH-1(aq) + 44.51 kJ

S

Enthalpy: _F__ Entropy: __F__

2 C(s) + 2 H2(g) + 52.4 kJ  C2H4(g)

N

Enthalpy: __U_ Entropy: __U__ Cont’d next page

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Answer the following questions about water freezing. 1 pt. ea. Pure water freezes at Standard Pressure at temperatures of 273 K/0.0°C or below, as shown by the reaction: H2O(l)  H2O(s) + 6.01 kJ 1. Is this process an increase or decrease in entropy? ___Decrease___ 2. Explain your answer for question #1 above. Reactants went from liquid (med entropy) to solid (low entropy)

3. Is the change in entropy favored or unfavored? ___Unfavored___ 4. Explain your answer for question #3 above. Entropy wants to go towards more disarray: S  L G

5. When water freezes, is it exothermic or endothermic? _____Exo______ 6. Explain your answer for question #5 above. 6.01 kJ is released (a product); entropy goes L  S

7. Is the change in enthalpy favored or unfavored? ___Favored___ 8. Explain your answer for question #7 above. Reaction loses heat, and loss of heat is favored

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Topic:

Regents Chemistry ’14-‘15 Mr. Murdoch

Changing Equilibrium

Objective: How can we change system equilibrium for our benefit? Equilibrium Changes:  Equilibrium systems are dynamic; this means that they are continuously in some form of change. However, the reaction MUST be in a closed system, or equilibrium cannot be maintained.  What if you want to change the equilibrium in a system? A system at equilibrium would be forming products at the same rate as the products would be decomposing back into the starting reactants.  We can change ONE aspect of a system at equilibrium at a time and force the system to do what WE want; we can control the reaction.

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Topic:

Regents Chemistry ’14-‘15 Mr. Murdoch

Le Chatelier’s Principle

Objective: What does changing an equilibrium do to a system? Le Chatelier’s Principle:

“If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change to reestablish an equilibrium.”  Le Chatelier’s Principle may be paraphrased to say this: o If a system at equilibrium has some a stressor (any factor that changes reaction rate), the equilibrium for that system will shift in a way that lessens the added stress, in the direction of whichever reaction rate was increased by the stressor. o Stressors in chemistry include: i. Temperature; ii. Concentration; iii. Pressure (only for gasses); iv. Murdoch  Any stressor introduced may cause a change in the concentrations of both the reactants AND the product until equilibrium is restored at a new point.

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Stressor

Regents Chemistry ’14-‘15 Mr. Murdoch

Shift Increases the number of collisions between reactant particles, driving the forward reaction faster

Change on Concentration

Removing reactant

Decreases the number of collisions between reactant particles, driving the reverse reaction faster

Reactants: Increases

Adding product

Increases the number of collisions between product particles, driving the reverse reaction faster

Reactants: Increase

Removing product

Decreases the number of collisions between product particles, driving the forward reaction faster

Reactants: Decrease

Increasing temperature

Favors the endothermic reaction, shifting the equilibrium away from the heat to absorb the excess heat energy

Depends on Direction of shift:

Adding reactant

Reactants: Decrease Products: Increase

Products: Decrease

Products: Decrease

Products: Increase

Favors the exothermic reaction, shifting If the shift is towards the products, then products the equilibrium towards the heat to will increase and reactants release energy will decrease Increasing System shifts towards side with fewer pressure If the shift is towards the moles of gas to reduce pressure (gases only) reactants, then reactants will increase and products Decreasing System shifts toward side with more will decrease pressure moles of gas to increase pressure (gases only) Decreasing temperature

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Shift

Topic:

Objective: How do we determine the possibility of equilibrium shift?  Some factors that may affect the rate of a reaction have NO effect on systems at equilibrium. These factors include: i. Catalysts ii. Inhibitors iii. Surface area  The above factors allow a system to achieve equilibrium faster, but once equilibrium is established, these factors affect both reactions equally. The equilibrium would not change then.

Water levels at equilibrium

Add water to left container; raises left level

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Water levels find new equilibrium; higher levels

Remove added water from left container; lowers left water level

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Water levels find new equilibrium

Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Determining Shift Direction Objective: How do we determine the direction of equilibrium shift? Application of Determining Direction of Equilibrium Shift: For the equilibrium: N2(g) + 3 H2(g)  2 NH3(g) + heat Stressor

Shift

Add N2(g)

Forwards-add a reactant; shifts to products

Remove N2(g)

Reverse-remove a reactant; shifts to reactants

Add H2(g)

Forwards-add a reactant; shifts to products

Remove H2(g)

Reverse-remove a reactant; shifts to reactants

Add NH3(g)

Reverse-add a product; shift to reactants

Remove NH3(g)

Forwards-remove a product; shift to products

Increase Temp

Reverse-increase Temp; shift away from heat

Decrease Temp

Forwards-decrease Temp; shift towards heat

Increase Press

Forwards-increase Press; shifts to side with fewer moles of gas

Decrease Press

Reverse-decrease Press, shifts to side with more moles of gas

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Change in Concentration N2(g): decreases H2(g): decreases NH3(g): increases N2(g): increases H2(g): increases NH3(g): decreases N2(g): decreases H2(g): decreases NH3(g): increases N2(g): increases H2(g): increases NH3(g): decreases N2(g): increases H2(g): increases NH3(g): decreases N2(g): decreases H2(g): decreases NH3(g): increases N2(g): increases H2(g): increases NH3(g): decreases N2(g): decreases H2(g): decreases NH3(g): increases N2(g): decreases H2(g): decreases NH3(g): increases N2(g): increases H2(g): increases NH3(g): decreases Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Topic:

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Shift

Objective: How do we determine the direction of equilibrium shift? For the equilibrium: KNO3(s) + 34.89 kJ  K

+1

(aq) +

NO3-1(aq)

1. What happens to the concentration of K+1(aq) when temperature is increased? i. Stressor: increased temperature ii. Shift: away from heat input (forward) iii. Change in Concentration: Since the shift is towards K+1, the concentration of K+1(aq) increases (along with [NO3-1(aq)] 2. What happens to the concentration of NO3-1(aq) when the temperature is decreased? i. Stressor: decreasing temperature ii. Shift: towards heat input (reverse) iii. Change in Concentration: Since the shift is away from NO3-1(aq), the concentration of NO3-1(aq) decreases (along with [K+1(aq)]

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Topic:

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Shift

Objective: How do we determine the direction of equilibrium shift? For the equilibrium: 2 CO(g) + O2(g)  2 CO2(g) + 566 kJ 1. What happens to the concentration of CO2(g) when CO(g) is added to the equilibrium system? i. Stressor: increasing concentration of a reactant ii. Shift: away from reactant (forward) iii. Change in Concentration: Since the shift is towards CO2(g), the concentration of CO2(g) increases 2. What happens to the concentration of O2(g) when CO2(g) is removed from the equilibrium system? i. Stressor: decreasing concentration of a product ii. Shift: towards product (forward) iii. Change in Concentration: Since the shift is away from O2(g), the concentration of O2(g) decreases 3. What happens to the concentration of CO(g) when pressure is increased? i. Stressor: increasing pressure ii. Shift: towards side with fewer moles of gas (forwards) iii. Change in Concentration: Since the shift is away from CO(g), the concentration of CO(g) decreases

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Topic:

Regents Chemistry ’14-‘15 Mr. Murdoch

Equilibrium Shift

Objective: How do we determine the direction of equilibrium shift? For the equilibrium: N2(g) + 2 O2(g) + 66.4 kJ  2 NO(g) 1. State five (5) things that can be done to the equilibrium that will result in an increase of the concentration of NO(g). Desired shift: to make more NO(g), you must shift towards NO(g), so shift equilibrium forwards  How can we drive the equilibrium forwards? a. Add N2(g): (adding reactant drives the reaction forward) b. Add O2(g): (adding reactant drives the reaction forward) c. Remove NO(g): (removing product drives the reaction forwards) d. Increase Temperature: (adding heat makes the reaction shift away from heat) e. Increase Pressure: (adding pressure shifts the equilibrium towards the side with fewer moles of gas) For the equilibrium: NaCl(s) + 3.88 kJ  Na+1(aq) + Cl-1(aq) 2. State four (4) things that can be done to increase the concentration of NaCl(s). Desired shift: to make more NaCl(s), you must shift towards NaCl(s), so shift equilibrium reverse  How can we drive the equilibrium in reverse? a. Remove NaCl: (removing a reactant makes the reaction shift in reverse) b. Remove heat: (removing heat makes the reaction shift towards the heat) c. Increase Na+1(aq): (adding product with Na+1 ions drives the reaction in reverse) d. Increase Cl-1(aq): (adding a product with Cl-1 ions drive the reaction in reverse) Website upload-Unit 9b: Equilibrium Systems

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Changing Equilibrium Practice Regents Problems: (ungraded) 1. Given the reaction at equilibrium: 2 SO2(g) + O2(g)  2 SO3(g) + heat Which change will shift the equilibrium to the right? a) b) c) d)

Increasing the pressure Increasing the temperature Decreasing the amount of O2(g) Increasing the amount of SO3(g)

2. Given the reaction at equilibrium: N2(g) + O2(g) + 182.6 kJ  2 NO(g) Which change would cause an immediate increase in the rate of the forward reaction? a) b) c) d)

Decreasing the reaction pressure Decreasing the reaction temperature Increasing the concentration of N2(g) Increasing the concentration of NO(g)

3. Given the equilibrium reaction: X + Y  2 Z + heat The concentration of the product may be increased by a) b) c) d)

Adding a catalyst Adding more heat to the system Decreasing the concentration of X Increasing the concentration of Y

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Notes page:

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Student name: _________________________

Class Period: _______

Please carefully remove this page from your packet to hand in. Changing Equilibrium Homework: For each of the following systems at equilibrium, predict the effect of a given change on the concentration of each of the specific substances. Write I if the concentration increases, D if the concentration decreases, and R if the concentration remains the same. 1. 2 NH3(g) + heat  N2(g) + 3 H2(g) Stressor #1: increase in [N2(g)]

Direction of shift: __  Rev __

What is the resulting effect on the concentration of: [NH3(g)]: _______I_______ Stressor #2: increase in temperature

[H2(g)]: _______D_______ Direction of shift: __Fwd ___

What is the resulting effect on the concentration of: [N2(g)]: _______I_______ Stressor #3: increase in pressure

[NH3(g)]: _______D_______ Direction of shift: ____ Rev____

What is the resulting effect on the number of moles of N2(g): _______D_______ What is the resulting effect on the number of moles of NH3(g): _______I_______ 2. 2 NO(g)  N2(g) + O2(g) + heat Stressor #1: decrease in [O2(g)]

Direction of shift: _____Fwd _____

What is the resulting effect on the concentration of: [N2(g)]: _______I_______ Stressor #2: decrease in temperature

[NO(g)]: _______D_______ Direction of shift: ____Fwd _____

What is the resulting effect on the concentration of: [O2(g)]: _______I_______

[NO(g)]: _______D_______ Direction of shift: __none __

There are the SAME # of moles on either side; What is the resulting effect on the number of moles of O2(g): _______R_______ pressure does NOT What is the resulting effect on the number of moles of NO(g): _______R_______affect equilibrium

Stressor #3: increase in pressure

Cont’d next page

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Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Complete the following questions by circling the correct answer. Given the equilibrium: N2(g) + 3 H2(g)  2 NH3(g) + heat 1. If N2(g) is added to the system at equilibrium, in which direction will the equilibrium shift? Forward

Reverse

2. If H2(g) is removed from the system at equilibrium, in which direction will the equilibrium shift? Forward

Reverse

3. If NH3(g) is added to the system at equilibrium, in which direction will the equilibrium shift? Forward

Reverse

4. If the temperature is decreased in the system at equilibrium, in which direction will the equilibrium shift? Forward

Reverse

5. If the pressure is increased in the system at equilibrium, in which direction will the equilibrium shift? Forward

Reverse

6. If H2(g) is removed from the system at equilibrium, what will happen to the concentrations of: N2(g):

Increase

Decrease

Remain the same

NH3(g):

Increase

Decrease

Remain the same

7. If NH3(g) is removed from the system at equilibrium, what will happen to the concentrations of: N2(g):

Increase

Decrease

Remain the same

H2(g):

Increase

Decrease

Remain the same

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Key

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key

Regents Chemistry ’14-‘15 Mr. Murdoch

Notes page:

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Unit 9b: Equilibrium, Enthalpy, and Entropy

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Key Regents Chemistry ’14-‘15 Mr. Murdoch Unit 9b: Equilibrium, Enthalpy, and Entropy 1. Student N...

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