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3/23/11
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3/9 Objective: SWBAT determine the driving force of a
chemical reaction. Do Now: 1) Hand in work completed yesterday. 2) Using the models at your pod, construct the
reactants of the following equation: H2 + O2 à
HW – Unit 9 pg 23 Q 1
Making things Happen
Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 1: Kinetics
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What causes a chemical reac9on? Kinetics:
The study of chemical reactions Mechanism:
How a reaction happens Think of reactant molecules…why do they react? 1. They must collide with each other 2. The collisions must have enough energy 3. The molecules must have the proper orientation Effective Collisions:
Collisions of reactant that cause them to react. 3
Reac9on Mechanisms Can be very complex. Here’s an easy one: Theoretical Reaction:
2A + B à A2B Reaction Mechanism:
A + B à AB + A àA2B
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Here’s a Real One: Reaction:
H2 + I2 à 2HI Step 1:
H2 à 2H –Absorbs energy Step 2:
I2 à 2I –Absorbs energy Step 3:
2H + 2I à 2HI—Releases energy Net Equation:
H2 + I2 à 2HI 5
All Road Lead To Rome Reactions can have different pathways depending on local conditions. Experimentation is needed to determine the reaction pathway. Rate-determining Step:
The slowest step in the reaction pathway. It determines the overall rate of the reaction.
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Factors that Affect Reac9on Rate Reaction Rate:
How many reactions happen/Unit of time
NOT HOW FAST/SLOW THE REACTION IS!!! Reaction Rate can be affected in 2 ways: 1. Changing the reaction mechanism. 2. Changing the number of collisions between
reactants
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Changing the Reac9on Rate
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1. Use of a catalyst Catalyst:
1. Changes one or more steps in a reaction (usually lowers the activation energy). 2. Shortens the reaction mechanism. 3. Is not consumed during the reaction.
Examples: – Biological enzymes – Catalytic converter in your car – Dissolving substances in water.
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2. Use of an Inhibitor Inhibitor:
1. Changes one or more steps in a reaction (usually increases the activation energy).
2. Lengthens the reaction mechanism. 3. Is not consumed in the reaction
Examples: – Patina on the surface of a metal
preventing oxidation of the rest of the metal
– Painting your car 10
Changing the Number of Collisions
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1. Nature of Reactants Different substances react at different rates. Has to do with how much energy those substances need to react. Ionic reactions happen much faster than covalent reactions. Why?
In solution, ionic bonds are broken, covalent bonds aren’t.
It takes more energy to break the covalent bonds.
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2. Temperature Why? Hotter the substance = higher the average KE. Higher the average KE = higher the rate of effective collisions. Examples: – Putting batteries in the fridge. – Heating up your food to cook it.
Which has a higher rate of collisions? A car going 30mph or a car going 100mph?
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3. Concentra9on Why? Higher concentration = more of a substance per unit of area. More substance present = higher rate of effective collisions Increasing pressure on gases is the same thing as increasing concentration. Examples: – Dilute vs. Concentrated Acids.
Which has a higher rate of collisions? A parking lot with 10 cars or 100 cars?
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4. Surface area Why? Higher surface area = more surface to collide with other reactants = higher rate of effective collisions. Only effects solids. Examples: – Dissolving a sugar cube vs. dissolving powdered
sugar. – Using a shotgun vs. using a bb gun. – Crushed garlic flavors more than a whole clove.
Which has a higher rate of collisions? A parking lot with 10 cars or 100 cars?
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To Review: Reactions depend upon effective collisions. Effective collisions require the proper orientation and energy of reactants. Reactions happen according to reaction pathways. Anything that affects the number of effective collisions or the pathway of the reaction will effect the rate of the reaction.
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Take out a sheet of paper, calc.., and ref tables.
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CO + 2H2 à CH3OH
1.2 grams of H2 are made to react with 7.45 grams of CO. What is the limiting reactant? How much CH3OH should be produced? If we actually recover 7.52g of CH3OH, what is our % yield?
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%yield = moles actual/moles theore9cal If we actually recover 7.52g of CH3OH, what is our
% yield?
19 20
What now?
Any Ques)ons?
3/10 Objective: SWBAT calculate the effect of
temperature of reaction rate. Do Now: Describe ways to speed up the following
reaction: N2 + H2 à NH3
HW: Study for quiz on %yield and Collision Theory
Quiz
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23 24
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3/23/11
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25
Alka Seltzer Lab
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3/11 Objective: SWBAT construct a potential energy diagram given a chemical reaction. Do Now: Quiz. HW: Chapter 9 pgs. 26-27.
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Fire and Ice
Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 2: Potential Energy Diagrams
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What Happens During a Reac9on? We know it involves reactants à products. But what about energy? A reaction can do one of two things with energy: 1. Absorb it (endothermic) 2. Release it (exothermic)
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All About H H = enthalpy:
A measurement of the amount of energy stored in a substance (units kJ/mole). For endothermic reactions ΔH is positive. For exothermic reactions ΔH is negative. ΔH values of many reactions are listed on Reference Table I.
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Endothermic Reac9ons General Formula:
A + B + Energy à C + D a. The reactants absorb energy (heat) b. This causes the temperature of the
surroundings to decrease. c. The products have more energy than the
reactants (stored in their bonds). d. The products are less stable than the reactants. This is how explosives are made!
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A General Example For the reaction:
A + B à C HA = 40kJ HB = 20kJ Hc = 110kJ
How much energy is absorbed during this reaction? What is ΔH? Rewrite the equation to show the Conservation of Energy:
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A Table I Example N2 (g) + O2 (g) à 2NO (g) ΔH = +182.6 kJ/mole
(+) ΔH = Endothermic Reaction. This means that Nitrogen and Oxygen need to absorb 182.6 kJ of energy per mole of NO that will be formed. NO is an unstable compound.
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Exothermic Reac9ons General Formula:
A + B à C + D + energy a. The reactants release energy (heat) b. This causes the temperature of the surroundings
to increase. c. The products have less energy than the reactants
(stored in their bonds). d. The products are more stable than the reactants. This is what is left behind after an explosion!
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A General Example For the reaction:
A + B à C HA = 60kJ HB = 40kJ Hc = 30kJ
How much energy is released during this reaction? What is ΔH? Rewrite the equation to show the Conservation of Energy:
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A Table I Example C (g) + O2 (g) à CO2 (g) ΔH = -393.5 kJ/mole
(-) ΔH = Exdothermic Reaction. This means that Carbon and Oxygen release 393.5 kJ of energy per mole of CO2 that will be formed. CO2 is a stable compound.
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Poten9al Energy Diagrams A graph that shows what happens to potential energy as a reaction occurs. What has to happen for a reaction to occur? 1. The reactants have some amount of energy
stored in their bonds. 2. To get them to react, some energy is put in to
the reactants (the “activation energy”) to get them to collide effectively. PE goes up until the reactants become the “activated complex”.
3. Once the activated complex reacts, the PE decreases to the energy of the products
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What goes in a PE diagram 1. Y-axis: PE (kJ) 2. X-axis: Reaction coordinate (unmeasured time). 3. H of reactants 4. H of activated complex 5. H of products. 6. The Activation energy (arrow from H of
reactants to H of activated complex) 7. ΔH: Arrow from (H of reactants to H of
products)
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An Endothermic PE Diagram
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Note: 1. PE increases. Temperature of surroundings will decrease.
2. The Products have more energy stored in their bonds than the reactants
An Exothermic PE Diagram
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Note: 1. PE decreases. Temperature of surroundings will increase.
2. The Products have less energy stored in their bonds than the reactants
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The Effect of A Catalyst Catalysts increase the rate of a reaction by lowering the activation energy. Less energy to get to the activated complex = higher frequency of effective collisions.
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Pgs. 24-‐25 PRACTICE
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43 44
45
Collision Theory POGIL
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What now?
Any Ques)ons?
3/14 Objective: SWBAT describe the properties of
chemical equilibrium. Do Now: 1) Hand in your PE Graph from HW.
2) Write your name on a separate sheet of paper, take out your reference tables, and a calculator. Close all other books.
HW: Ch 9 pgs 28-29 Questions.
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N2 (g) + 3 H2 (g) à 2 NH3
4.0 L of each reactant are used to produced 1.5g of NH3. 1. Which is the limiting reagent?
2. Calculate the percent yield.
49 50
Pgs. 26-‐27 Review
51 52
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⇋ Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 3: Equilibrium Systems 54
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Equilibrium Strikes Again Equilibrium:
A state of rate balance between two opposing changes. The rate of the forward change is equal to the rate of the reverse change. Most reactions are reversible:
A + B ↔ C + D + energy The double arrow means that both reactions are occurring simultaneously
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Forward and Reverse A Forward reaction:
N2 (g) + 3H2 (g) à 2NH3 (g) + 92kJ
A Reverse reaction: 2NH3 (g) + 92kJ à N2 (g) + 3H2 (g)
An Equilibrium:
N2 (g) + 3H2 (g) ↔ 2NH3 (g) + 92kJ
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Proper9es of Equilibrium 1. Equilibrium is dynamic. The forward and reverse
reactions are occurring simultaneously. Since the rates of both are equal, it looks like nothing is happening, but don’t be fooled!
2. Equilibrium can only happen in a closed system, isolated from its surroundings.
3. Changing the conditions of the equilibrium system will change the equilibrium.
4. Don’t fool yourself into thinking that just because the rates of the reactions are the same, that the concentrations of the reactants and products also have to be the same. They just have to be constant.
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Different Equilibrium Points
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Types of Equilibrium 1. Chemical Equilibrium:
The rates of a forward and reverse reaction are equal.
2. Solution Equilibrium: The rate of a solvent dissolving is equal to the rate of a dissolved solvent precipitating out of solution.
3. Physical Equilibrium (aka “Phase Equilibrium”) The rate of a forward phase change is equal to the rate of a reverse phase change.
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Chemical Equilibrium Keq= the “equilibrium constant” The ratio of the amount of product to reactant at equilibrium Keq = [products]/[reactants]
Square brackets mean “concentration”
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The Mass Ac9on Expression The ratio of the concentrations of the products to the concentrations of the reactants. Tells us how far a reaction proceeded before reaching equilibrium. For a system:
aA + bB ↔ cC + dD
Keq = [C]c [D]d
[A]a [B]b
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Set up the Mass Ac9on Expression!
2SO2 (g) + O2 ↔ 2SO3 (g)
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Set up the Mass Ac9on Expression!
CaCl2 (s) ↔ Ca2+ (aq) + 2Cl- (aq)
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Set up the Mass Ac9on Expression!
2O3 (g) ↔ 3O2 (g)
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Keq Keq is what we get when we solve the mass action expression. It has no units. It’s just a ratio. For a system:
A + B ↔ 2C + Energy The equilibrium concentrations are [A]=2M, [B]= 5M, and [C]=10M. What is Keq?
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What does Keq mean Keq tells us how much product was formed before the system reached equilibrium. There are 5 major points where Keq is informative
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Value of Keq [Products] : [Reactants] Meaning 0 <1 1 >1 ∞
[Products]<<[Reactants] [Products] < [Reactants] [Products] = [Reactants] [Products] > [Reactants] [Products]>>[Reactants]
Reac9on didn’t start Reac9on went a bit
Reac9on went halfway Reac9on went mostway Reac9on to comple9on
Bo-om line: Bigger the Keq, the closer the reac9on gets to comple9on before reaching equilibrium
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3/15 Objective: SWBAT calculate how chemical equilibrium can
be disturbed. Do Now: Find the Keq for the following reaction:
S2(g) + 2H2(g) 2H2S H2 = 2.16 M S2 = 0.30 M H2S = 0.50 M HW: Chapter 9 pg. 30 Questions.
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⇋
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Solu9on Equilibrium There is only a certain amount of space for solute particles in a solution. Once that limit gets reached, the solution is “saturated” (more on that next unit) At the saturation point, every solvent molecule that enters solution means one has to precipitate out.
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Keq works for solu9ons, too! It works exactly the same way that it does for chemical changes. Bottom line: The higher the Keq for a solution, the more solvent is dissolved.
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Ksp Ksp- The solubility product constant Used for substances with very low solubility. We don’t divide by the amount of undissolved solute (it would make things ridiculously small). For a system
MA2 (s) ↔ M+ (aq) + 2A- (aq) Ksp = [M+] [A-]2
Bottom Line: The higher the Ksp, the more soluble the substance is.
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Honors Reference Table A A table of solubility product constants for nearly insoluble salts at 1 atm and 298 K
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Physical Equilibrium When two phase changes are in equilibrium. Happens at particular temperatures: Ex:
H2O (s) ↔ H2O at 273K
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Pgs. 28-‐29 Review
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75 76
77
Equilibrium Demo
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Bean Simula9on Lab
79 80
What now?
Any Ques)ons?
⇋ 3/16
Objective: SWBAT describe changes in entropy as favored or unfavored. Do Now: Did part I reach equilibrium? How do you know? HW: 1) Bring in design draft tomorrow. 2) Ch. 9 pg 31 Questions.
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Ready, Set, Go!
Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 4: Will a Reaction Happen On Its Own…
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Spontaneity Spontaneous Chemical Reactions:
Reactions that proceed on their own once initiated. The spontaneity of a reaction depends on two things: 1. Enthalpy 2. Entropy
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Enthalpy (H) Enthalpy:
The amount of heat energy stored in a system. Nature favors reactions that undergo a DECREASE in enthalpy (that release energy). Spontaneous reactions tend to release energy. Most exothermic reactions are spontaneous. Most endothermic reactions are non-spontaneous. Does a match keep burning once struck? Why?
84
3/23/11
15
Entropy (S) Entropy:
The randomness (disorder) of a system. Nature favors reactions that undergo an INCREASE in entropy. As temperature increases, entropy increases. Entropy increases from solid àliquid àgas. Why?
85
4 possible situa9ons 1. Exothermic reactions that increase entropy:
ALWAYS SPONTANEOUS 2. Endothermic reactions that decrease entropy:
ALWAYS NONSPONTANEOUS 3. Exothermic reactions that decrease entropy:
SPONTANEOUS AT LOW TEMPERATURES 4. Endothermic reactions that increase entropy:
SPONTANEOUS AT HIGH TEMPERATURES
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Gibbs Free Energy (G) A measurement of the amount of energy available in a system to do work. Nature favors a decrease in free energy. If the change in free energy (ΔG) is negative in a system, the reaction will be spontaneous.
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Josiah Willard Gibbs (1839 – 1903)
ΔG = ΔH – (T ΔS) The Gibbs Free Energy Equation!
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ΔH ΔS ΔG Spontaneous?
-‐ + Always nega9ve
Always
-‐ -‐ Nega9ve if T is low
At Low Temp
+ -‐ Nega9ve if T is high
At High Temp
+ + Never nega9ve
Never
Fun with Gibbs 1. Calculating ΔG: For the fictitious reaction A + B à C + 30.0 kJ at 25oC (298 K): If ∆H for this reaction is –30.0 kJ and ∆S is –0.010 kJ/K, calculate ∆G and determine if this reaction is spontaneous or nonspontaneous.
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3/16 Objective: SWBAT apply the Gibbs free energy equation. Do Now: Is this reaction spontaneous? How do you know? Ammonia is synthesized from nitrogen and hydrogen gases at
a temperature of 475 degrees Celsius.
N2(g)+3H2(g) --->2NH3(g)
If delta H = -92.2kJ and delta S= -0.1987kJ/K HW: Science Fair Design due tomorrow. Quiz tomorrow. Test next week.
90
3/23/11
16
Using Honors Reference Table B Table B- Energies of Formation of Compounds at 1 atm and 298K. Tells us about the spontaneity and ΔH associated with making compounds at specific conditions. For synthesis. Reverse values for decomposition.
91
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!
Determining ΔS Rearrange the equation to solve for ΔS:
ΔS = (ΔH-ΔG)/T Calculate ΔS for CO (use Table B values) Is it spontaneous always? Sometimes? Never?
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Calculate the ΔS for C2H4 Is it spontaneous always? Sometimes? Never?
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Calculate ΔS for ICl Is it spontaneous always? Never? Sometimes?
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Determine the Temperature when something becomes spontaneous
The temperature at which a reaction becomes spontaneous is the equilibrium temperature. ΔG is 0 at this temperature.
ΔG = ΔH – (TΔS) 0 = ΔH – (TΔS)
TΔS =ΔH T=ΔH/ΔS
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Determine the temperature at which ICl reaches equilibrium
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3/23/11
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3/18 Objective: SWBAT predict the direction an equilibrium will shift when stress is applied. Do Now: Quiz HW: Edited design due MondayUnit 9 Test next week.
97 98
What now?
Any Ques)ons?
3/21 Objective: SWBAT predict the concentrations of reactants and products at equilibrium. Do Now: 1) Calculate the formula mass of AlBr3. Show your work. 2) Hand in Science Fair Designs. HW: pg 32 Questions. Chapter 9 Test Friday
99
Like a SeeSaw
Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 5: Changing Equilibrium 100
Equilibrium can Shiq LeChatelier’s Principle: If a system at equilibrium is subjected to a stress, the system will shift in a way that relieves the stress. This will cause a change in concentration of reactants and products until equilibrium is re-established.
101
Henry Louis Le Chatelier (1850 – 1936)
Shiqing Equilibrium
“Forward Shift”: Shifting to favor the forward reaction (aka “shifting to the right”) “Reverse Shift”: Shifting to favor the reverse reaction (aka “shifting to the left”)
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Reactants ↔ Products Forward Shiq à
ß Reverse Shiq
3/23/11
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A Pile of Shiqs You Must Learn These!
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Real Big Shiqs
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N2 (g) + 3H2(g) ↔ 2NH3 (g) + heat
KNO3 (s) + 34.89kJ ↔ K+1 (aq) + NO3-‐1(aq)
What happens to [K+1] when the temperature is increased? What happens to [NO3
-1] when the temperature is decreased?
105
Shiq!
2CO (g) + O2 (g) ↔ 2CO2 (g) + 566kJ
What happens to [CO2] when [CO] is added? What happens to [O2] when [CO2] is removed? What happens to [CO] when pressure is increased?
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Shiq!
N2 (s) + 2O2 + 66.4kJ ↔ 2NO(g) List five things that can be done that will increase [NO]
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Shiqs!
NaCl (s) + 3.88kJ ↔ Na+(aq) + Cl-‐(aq)
List four things that can be done that will increase [NaCl]
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3/23/11
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LeChatlier Principle POGIL
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3/22 Objective: SWBAT predict how a given shift will affect dynamic equilibrium. Do Now: If the beaker started with 100 g of water, how would you know that the following reaction is at equilibrium?
H2O(l) ↔ H2O(g)
HW: Chap 9 pg 33 questions
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Pg.30 Review
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112
Pg 31 Q Review
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3/23/11
20
115
LeChatlier Principle POGIL
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Systems Thinking Ac9vity • Make sure safety rules are on write-up
117
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What now?
Any Ques)ons?
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