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Chemical Kinetics
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Close-up of a crashing wave.
Source: Getty Images
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Figure 15.1: Starting with pure nitrogen dioxide at 300°C
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Rate Law
instantaneous rate of reaction.
Rate = k(phenolphthalein)
the concentration of phenolphthalein in a solution that was initially 0.005 M in phenolphthalein and 0.61 M in OH- ion.
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Examples of Rate Law
2 HI(g) H2(g) + I2(g)
CH3Br(aq) + OH-(aq) CH3OH(aq) + Br-(aq)
Rate = k(CH3Br)(OH-)
(CH3)3CBr(aq) + OH-(aq) (CH3)3COH(aq) + Br-
But Rate = k((CH3)3CBr)
Rate law is not related to stochiometry!
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Decomposition of N2O5
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Figure 15.2: Plot of the concentration of N2O5
Rate = k(N2O5)
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Figure 15.3: A plot of In[N205] versus time.
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Figure 15.4: Plot of [N2O5] versus time for the decomposition reaction of N2O5.
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Rate = k(X)We then test this assumption by checking concentration versus time data
for the reaction to see whether they fit the first-order rate law.
ln (X) - ln (X)0 = - kt y = mx + b
ln (X) = -kt + ln (X)0
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2nd Order Reaction
Rate = k(X)2
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Solution containing BrO3-
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Butadiene and its dimer
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Figure 15.5: (a) A plot of ln[C4H6] versus t. (b) A plot of 1/[C4H6] versus t.
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Figure 15.6: Plot of [A] versus t for a zero-order reaction
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Figure 15.7: Decomposition reaction takes place on a platinum surface
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Figure 15.8: Molecular representation of the elementary steps in the reaction of NO2 and CO.
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Single Step Reaction
CH3Br(aq) + OH-(aq) CH3OH(aq) + Br-(aq)
Rate = k(CH3Br)(OH-)
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Multiple Step Reaction
(CH3)3CBr(aq) + OH-(aq) (CH3)3COH(aq) + Br-(aq)
(CH3)3CBr (CH3)3C+ + Br- Slow step
(CH3)3C+ + H2O (CH3)3COH2
+ Fast step
(CH3)3COH2+ + OH- (CH3)3COH + H2O Fast step
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General Rules for Rate Law
•The rate of any step in a reaction is directly proportional to the concentrations of the reagents consumed in that step.•The overall rate law for a reaction is determined by the sequence of steps, or the mechanism, by which the reactants are converted into the products of the reaction.•The overall rate law for a reaction is dominated by the rate law for the slowest step in the reaction.
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The rate a which this colored solution enters the flask is determined by the size of the funnel stem, not how fast the solution is poured.
Source: American Color
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Slow and Fast reactions shown as molecular models
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Figure 15.10: A plot showing the exponential dependence of the rate constant on the absolute temperature
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Figure 15.11: The change in potential energy as a function of reaction progress for the reaction
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Figure 15.12: Plot showing the number of collisions with a particular energy
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Figure 15.13: Several possible orientations for a collision between two BrNO molecules.
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Figure 15.14: Plot of In(k) versus 1/T for the reaction
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Figure 15.15: Energy plots for catalyzed and uncatalyzed pathways for a given reaction
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Homogeneous Catalysis
H2O2(aq) + I-(aq) H2O(aq) + OI-(aq)
In the second step, the OI- ion is reduced to I- by H2O2.
OI-(aq) + H2O2(aq) H2O(aq) + O2(g) + I-(aq)
the first step in this reaction is the rate-limiting step,
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Figure 15.16: Effect of a catalyst on the number of reaction-producing collisions.
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Figure 15.17: Heterogeneous catalysis of the hydrogenation of ethylene.
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Figure 15.18: The exhaust gases from an automobile engine are passed through a catalytic converter to
minimize environmental damage.
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General structure of protein
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Figure 15.19: Removal of the end amino acid from a protein by reaction with a molecule of water.
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Figure 15.20: Structure of the enzymecarboxypeptidase-A, which contains 307 amino acids.
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Figure 15.21: Protein-substrate interaction
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Cutaway model of a catalytic converter used in automotive exhaust systems.
Source: Delphi Automotive Systems
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Collect samples of extremophiles from
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A micrograph of the extremophile Archaeoglobus fulgidis, and organis that lives in the hot sediments near
submarine hydrothermal vents.
Source: Photo Researchers, Inc.
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Reaction coordinate
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Reaction coordinate
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Rate
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Concentration of reactant
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Time (s); Time (s); Time (s)
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Laser spectroscopy
Source: California Institute of Technology
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Figure 15.9: The STM images of the reaction of CO and O2
Source: University of California, Irvine