chemical kinetics
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Chemical Kinetics. SCH4U: Grade 12 Chemistry. Unit Mind Map. Chemical Kinetics. Chemical kinetics? Why do we study this? Who makes use of chemical kinetics?. Chemical Kinetics. Chemical kinetics studying how we can make reactions go faster or slower Why do we study this? - PowerPoint PPT PresentationTRANSCRIPT
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Chemical KineticsSCH4U: Grade 12 Chemistry
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Unit Mind Map
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Chemical Kinetics• Chemical kinetics?
• Why do we study this?
• Who makes use of chemical kinetics?
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Chemical Kinetics• Chemical kinetics studying how we
can make reactions go faster or slower• Why do we study this?
• Who makes use of chemical kinetics?
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Chemical Kinetics• Chemical kinetics studying how we
can make reactions go faster or slower• Why should we study this?
• Economics effects Materials are expensive
• Developing pharmaceutical drugs
• Who makes use of chemical kinetics?
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Chemical Kinetics• Chemical kinetics studying how we can
make reactions go faster or slower• Why should we study this?
• Economics effects Materials are expensive• Developing pharmaceutical drugs
• Who makes use of chemical kinetics?• Biologists metabolic rxns, food digestion, bone
regeneration• Automobile engineers rate of rusting,
decrease pollutants• Agriculture slow down food ripening
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Reaction Rates• What is a chemical reaction rate?
• Possible formula for measuring rate?
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Reaction Rates• What is a chemical reaction rate?
• measure of how fast reactants are used up or how fast products are produced
• Possible formula for measuring rate?
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Reaction Rates• What is the reaction rate?
• measure of how fast reactants are used up or how fast products are produced
• Possible formula for measuring rate?• Rate = Δ Concentration/Δ Time
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Sample problemN2 (g) + 3H2 (g) 2NH3(g)
• What is the average rate of production of ammonia for the system if the concentration is 3.5mol/L after 1.0min and 6.2mol/L after 4.0 min?
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Sample problemN2 (g) + 3H2 (g) 2NH3(g)
• What is the average rate of production of ammonia for the system if the concentration is 3.5mol/L after 1.0min and 6.2mol/L after 4.0 min?
• ANSWER: 0.9mol/(L*min)
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Measuring reaction rates Graphically
• Average rate of reaction take the slope of the secant of the line
• Instantaneous rate of reaction take the slope of the tangent of the line
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Measuring reaction rates Graphically
• Average rate of reaction take the slope of the secant of the line
1. Draw a secant line between two points
2. Calculate the slope
Slope = Rise/Run= Δconcentration/Δ time
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Measuring reaction rates Graphically
• Instantaneous rate of reaction take the slope of the tangent of the line
1. Draw a tangent line to the graph
2. Calculate the slope of the tangent line
Slope = Rise/Run= Δconcentration/Δ time
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Measuring reaction rates
• What are some factors that we can measure experimentally to determine the rate of reaction ?
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Measuring reaction rates
• Production of a gas
• Production of ions
• Changes in colour
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Reaction rates and stoichiometry
CO(g) + NO2(g) CO2(g) + NO(g)
In this reaction, the ratio of CO to NO is 1:1
Therefore, the disappearance of CO is the same as the production of NO Rate = - Δ [CO]/Δt = + Δ[NO]/Δt
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Reaction rates and stoichiometry
Suppose the ratio is NOT 1:1?Example, H2 (g) + I2 (g) 2HI (g)
2 mols of HI are produced for every 1 mol H2 usedRate =
The rate at which H2 is used up is only half of which HI is produced
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Sample Problem• IO3
- (aq) + 5I- (aq) + 6H+ (aq) 3I2 (aq) + 3H2O (l)
The rate of consumption of iodate ions (IO3-)
is determined experimentally to be 3.0 x 10-5 mol/(L*s). What are the rates of reaction or the other reactants and products in this reaction? • Complete p.364 #3, 4• Homework: practice problem package
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Collision Theory• Why do reactions occur the way that they
do?
VS
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Collision Theory• Collision theory reactions can only
occur if:1. There is a collision between molecules
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Collision Theory• Collision theory
reactions can only occur if:
1. There is a collision between molecules 2. the molecules are oriented in the correct way
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Collision Theory• Collision theory reactions
can only occur if:1. There is a collision between molecules2. the molecules are oriented in the correct way 3. enough energy is provided to break the chemical bonds that hold molecules together (Activation energy)
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Activation Energy• Activation Energy minimum potential
energy the system needs to overcome for the molecules to react
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Factors That Affect Rate of Reaction
• What are some of the factors that may speed up or slow down a chemical reaction? P.392
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Factors That Affect Rate of Reaction
• What are some of the factors that may speed up or slow down a chemical reaction?• Chemical nature of the reactant• Concentration of reactants• Surface area of reactants• Temperature• Catalysts
• Get into groups, brainstorm and then explain these factors to the rest of class in a creative way, making use of collision theory
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Lesson 2• Review collision theory
• Rate Laws
• Group Practice problems
• Individual practice problems
• homework
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Unit Mind Map
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Collision theory
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Collision Theory• Collision theory reactions
can only occur if:1. There is a collision between molecules2. the molecules are oriented in the correct way 3. enough energy is provided to break the chemical bonds that hold molecules together (Activation energy)
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Factors That Affect Rate of Reaction
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Factors That Affect Rate of Reaction
• Chemical nature of the reactant
• Concentration of reactants• Surface area of reactants• Temperature• Catalysts
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The Rate Law• Mathematical relationship between reaction
rate and factors that affect it• Determined empirically (experimentally)• Rate = k[X]m[Y]n
• e.g. 2NO2 + F2 2NO2FThe above reaction is 1st order with respect to NO2 and 2nd order with respect to F2. What is the rate law equation?
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The Rate LawAnswer:• e.g. 2NO2 + F2 2NO2F
rate = k[NO2]1[F2]2
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Steps to solve rate law problems
eg. 2BrO3-
(aq) + 5HSO3- (aq) Br2 (g) + 5SO4
2- (aq) + H2O (l) + 3H+ (aq)
Rate = k [BrO3-]m [HSO3
-]n
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Steps to solve rate law problems
eg. 2BrO3-
(aq) + 5HSO3- (aq) Br2 (g) + 5SO4
2- (aq) + H2O (l) + 3H+ (aq)
1. Write out rate equation: rate = k [BrO3-]m [HSO3
-]n
2. pick a trial where [BrO3-] changes but [HSO3
-] stays constant= trial 1 and trial 2
3. Using a ratio, determine the relationship between change in concentration and change in rate = (trial 1/trial2) = (4.0/2.0) = (1.6/0.8)
= 2 = 24. Using the following chart, determine the rate order of [BrO3
-]
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Steps to solve rate law problems
5. If the concentration is doubled (2.0 to 4.0), the rate also doubles (0.80 to 1.60), therefore the rate order is 1m=1
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Steps to solve rate law problems
6. To find the rate order of the other reactant, follow the same steps but pick a trial where [HSO3
-] changes but [BrO3-] stays
constant= trial 2 and trial 3= (6.0/3.0) = (0.8/0.2)= 2 = 47. Use the following table to determine the rate order
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Steps to solve rate law problems
8. As the concentration doubles the rate was multiplied by 4, therefore rate order of HSO3
- = 2 (n=2)9. Plug values into rate equation rate = k [BrO3
-]m [HSO3-]n and
solve for k.
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Example 1
From the data collected above, determine the rate law of the following equation:aA + bB products• Rate = k[A]m[B]n
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Example 1
• aA + bB products• Rate = k[A]m[B]n
• m = 0; n = 1; k = 0.6s-1
• Rate = 0.6s-1 [B]1
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Example 2
From the data collected above, determine the rate law of the following equation:• 2NO2 + F2 2NO2F• Rate = k[A]m[B]n
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Example 2
• 2NO2 + F2 2NO2F• Rate = k[A]m[B]n
• m = 2, n = 0, k = 4 x 10-3M-1s-1
• Rate =4 x 10-3M-1s-1 [NO2]2
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Team problem• Get into groups of 3 and obtain problem clues
a) Determine the order with respect to each reactant b) Determine the overall order of reactionc) Write the rate expression for the reaction.d) Find the value of the rate constant, k.
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Practice Problem
• From the data collected above, determine the rate law for the following equation:
IO3- (aq) + 5I- (aq) + 6H+ (aq) 3I2 (aq) + 3H2O (l)
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Practice Problem
• From the data collected above, determine the rate law for the following equation:
IO3- (aq) + 5I- (aq) + 6H+ (aq) 3I2 (aq) + 3H2O (l)
• m = 1; n = 1; o = 2; k = 5M-3s-1
• Rate = 5M-3s-1 [IO3-]1 [I-]1 [H+]2
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Homework• P. 377 practice # 1, 2, 3, 6,
• P. 415 # 15 a-e
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Lesson 3• Reaction mechanisms:
• Ping pong activity• Assembly line
• Elephant toothpaste demonstration• Practice problems
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Unit Mind Map
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Reaction mechanisms
• Rate Laws determined experimentally• Equation: Reactants products• This actually occurs in a series of steps
called elementary steps• Analogy: Cooking takes place in several steps
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Reaction mechanisms
• What are the chances of the following reaction occurring in one step?
Hint: collision theory• 4HBr (g) + O2 (g) 2H2O (g) + 2Br2 (g) actually
occurs in 3 separate steps
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Reaction mechanisms
4HBr (g) + O2 (g) 2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps:HBr (g) + O2 (g) HOOBr (g) (slow)HOOBr (g) + HBr (g) 2HOBr (g) (fast)2HOBr (g) + 2HBr (g) 2H2O (g) + 2Br2 (g) (fast)4HBr (g) + O2 (g) 2H2O (g) + 2Br2
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Reaction mechanisms
4HBr (g) + O2 (g) 2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps:HBr (g) + O2 (g) HOOBr (g) (slow)HOOBr (g) + HBr (g) 2HOBr (g) (fast)2HOBr (g) + 2HBr (g) 2H2O (g) + 2Br2 (g) (fast)4HBr (g) + O2 (g) 2H2O (g) + 2Br2 Rate law: rate = k[HBr2][O2]
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Reaction mechanisms:
working backwardsOverall equation: 4HBr (g) + O2 (g) 2H2O (g) + 2Br2 (g)
Rate law: rate = k[HBr]1[O2]1
Devise a proposed mechanism for this reaction
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Reaction mechanisms:
working backwardsOverall equation: 4HBr (g) + O2 (g) 2H2O (g) + 2Br2 (g) Rate law: rate = k[HBr][O2]Devise a proposed mechanisms for this reactionStep 1. using the rate law, write out the rate-determining (slow) stepStep 2. use the overall equation to determine what needs to be added to achieve the overall equationStep 3. cross out intermediates and add up what is left to produce the overall equation
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Reaction mechanisms:
working backwards4HBr (g) + O2 (g) 2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps:HBr (g) + O2 (g) HOOBr (g) (slow)HOOBr (g) + HBr (g) 2HOBr (g) (fast)2HOBr (g) + 2HBr (g) 2H2O (g) + 2Br2 (g) (fast)4HBr (g) + O2 (g) 2H2O (g) + 2Br2
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Potential Energy Diagram
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Elephant toothpaste demonstration
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Elephant toothpaste demonstration
Elephant toothpaste demonstration reaction:2H2O2 O2 + 2H2ORate = k[H2O2]1[I-]1
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Elephant toothpaste demonstration
Elephant toothpaste demonstration reaction:2H2O2 O2 + 2H2ORate = k[H2O2]1[I-]1
Proposed Reaction Mechanism:H2O2 + I- IO- + H2O (Slow)H2O2 + IO- I- + H2O + O2 (Fast)2H2O2 O2 + 2H2O
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Elephant toothpaste demonstration
Possible energy potential diagram?
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Practice ProblemPropose a possible mechanism for the following reaction:2N2O5 (g) 2N2O4 (g) + O2 (g)
r = k[N2O5]1
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Practice ProblemPropose a possible mechanism for the following reaction:2N2O5 (g) 2N2O4 (g) + O2 (g)
r = k[N2O5]1
Possible mechanism:N2O5 N2O4 + O (Slow)O + N2O5 N2O4 + O2 (Fast)2N2O5 2N2O4 + O2
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Practice ProblemPropose a possible mechanism for the following reaction:2NO2 + F2 2NO2F
r = k[NO2]1[F2]1
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Practice ProblemPropose a possible mechanism for the following reaction:2NO2 + F2 2NO2Fr = k[NO2]1[F2]1
Possible mechanism:NO2 + F2 NO2F + F (Slow)F + NO2 NO2F (Fast)2NO2 + F2 2NO2F
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Homeworkp. 390 #2p. 391 # 1, 2, 3Also, complete worksheet.