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X2 + Y2 --> X2Y2 rate = k[X2]
A reaction and its experimentally determined rate law are represented above. A chemist proposes two different possible mechanisms for the reaction, which are given below.
X2 + Y2 --> X2Y2 rate = k[X2] A reaction and its experimentally determined rate law are represented above. A chemist proposes two different possible mechanisms for the reaction, which are given below.
- Only mechanism 2 is consistent with the rate law.
- Neither mechanism 1 nor mechanism 2 is consistent with the rate law.
- Only mechanism 1 is consistent with the rate law.
- Both mechanism 1 and mechanism 2 are consistent with the rate law. (correct)
4 FeS(s) + 7O2(g) --> 2Fe2O3(s) + 4SO2(g)
4 FeS(s) + 7O2(g) --> 2Fe2O3(s) + 4SO2(g)
- A diagram showing a small activation energy for the forward reaction, typical of a fast reaction.
- A diagram depicting a large activation energy for the forward reaction, indicating a slow reaction.
- A diagram with activation energy measured from the products' energy level to the peak of the energy curve, relevant for the reverse reaction.
- A diagram showing the activation energy extending from the reactants' energy level to the peak of the energy curve, accurately representing the energy barrier. (correct)
A student studying the kinetics of a reaction measures the concentration of reactant X over time. Which procedure will allow the student to determine the rate constant, k, for the reaction?
A student studying the kinetics of a reaction measures the concentration of reactant X over time. Which procedure will allow the student to determine the rate constant, k, for the reaction?
- Plot [X] versus time and determine the slope. (correct)
- Plot 1/[X] versus time and determine the slope.
- Plot [X]^2 versus time and determine the slope.
- Plot ln[X] versus time and determine the magnitude of the slope.
For the reaction $C_2H_4 + H_2 \rightarrow C_2H_6$, which action is most likely to increase the rate of reaction?
For the reaction $C_2H_4 + H_2 \rightarrow C_2H_6$, which action is most likely to increase the rate of reaction?
In a proposed two-step mechanism for ozone ($O_3$) destruction in the upper atmosphere, what is the expected concentration of intermediate species during the reaction?
In a proposed two-step mechanism for ozone ($O_3$) destruction in the upper atmosphere, what is the expected concentration of intermediate species during the reaction?
How does increasing $[H^+]$ at a constant temperature affect a reaction involving $H^+$ and $ClO^-$ ions?
How does increasing $[H^+]$ at a constant temperature affect a reaction involving $H^+$ and $ClO^-$ ions?
Which type of reaction energy profile best represents a two-step exothermic reaction with a fast second step?
Which type of reaction energy profile best represents a two-step exothermic reaction with a fast second step?
If two solid reactants with varying particle sizes are mixed, what condition results in the fastest reaction rate, assuming the same mass of each reactant in each case?
If two solid reactants with varying particle sizes are mixed, what condition results in the fastest reaction rate, assuming the same mass of each reactant in each case?
In an experiment, one student observes a significantly slower reaction rate compared to their peers. Which of the following variations in their procedure could explain this?
In an experiment, one student observes a significantly slower reaction rate compared to their peers. Which of the following variations in their procedure could explain this?
Two samples of $Mg(s)$ of equal mass react with $HCl(aq)$ in separate vessels. If the $Mg(s)$ is more finely divided in one vessel, how does this affect the reaction, and why?
Two samples of $Mg(s)$ of equal mass react with $HCl(aq)$ in separate vessels. If the $Mg(s)$ is more finely divided in one vessel, how does this affect the reaction, and why?
What is the primary effect of a solid nickel catalyst on a catalyzed reaction?
What is the primary effect of a solid nickel catalyst on a catalyzed reaction?
In a potential energy diagram for the uncatalyzed decomposition of $H_2O_2(l)$, arrows labeled A and B represent different amounts of energy. With a catalyst present, how do the magnitudes of A and B compare to the uncatalyzed reaction?
In a potential energy diagram for the uncatalyzed decomposition of $H_2O_2(l)$, arrows labeled A and B represent different amounts of energy. With a catalyst present, how do the magnitudes of A and B compare to the uncatalyzed reaction?
For the reaction $2N_2O(g) \rightarrow 2N_2(g) + O_2(g)$ with rate = $k[N_2O]$, what explains why increasing the initial concentration of $N_2O(g)$ increases the reaction rate?
For the reaction $2N_2O(g) \rightarrow 2N_2(g) + O_2(g)$ with rate = $k[N_2O]$, what explains why increasing the initial concentration of $N_2O(g)$ increases the reaction rate?
Which statement effectively explains why increasing temperature increases the rate of a chemical reaction?
Which statement effectively explains why increasing temperature increases the rate of a chemical reaction?
For the decomposition of $H_2O_2(aq)$, how can the order of the reaction be determined from experimental data of concentration vs. time?
For the decomposition of $H_2O_2(aq)$, how can the order of the reaction be determined from experimental data of concentration vs. time?
How does a solid catalyst increase the rate of a reaction where the slowest step requires a collision between two gaseous molecules?
How does a solid catalyst increase the rate of a reaction where the slowest step requires a collision between two gaseous molecules?
Two trials of a first-order catalytic decomposition of $H_2O_2(aq)$ are performed. If trial 2 has a lower initial rate than trial 1, which factor best explains this difference?
Two trials of a first-order catalytic decomposition of $H_2O_2(aq)$ are performed. If trial 2 has a lower initial rate than trial 1, which factor best explains this difference?
Given the reaction $2NO(g) + 2H_2(g) \rightarrow N_2(g) + 2H_2O(g)$ with the rate law rate = $k[NO]^2[H_2]$, what is the overall order of the reaction?
Given the reaction $2NO(g) + 2H_2(g) \rightarrow N_2(g) + 2H_2O(g)$ with the rate law rate = $k[NO]^2[H_2]$, what is the overall order of the reaction?
Which change is most likely to increase the rate of reaction between $Li(s)$ and water?
Which change is most likely to increase the rate of reaction between $Li(s)$ and water?
For the reaction $NO(g) + NO_3(g) \rightarrow 2NO_2(g)$, which orientation of collision is most likely to be effective?
For the reaction $NO(g) + NO_3(g) \rightarrow 2NO_2(g)$, which orientation of collision is most likely to be effective?
The rate of the reaction $N_2O(g) + CO(g) \rightarrow N_2(g) + CO_2(g)$ increases significantly in the presence of $Pd(s)$. Which best explains this observation?
The rate of the reaction $N_2O(g) + CO(g) \rightarrow N_2(g) + CO_2(g)$ increases significantly in the presence of $Pd(s)$. Which best explains this observation?
For the single-step reaction $NO(g) + NO_3(g) \rightarrow 2NO_2(g)$, a scientist calculates the rate of collisions but finds the observed rate much lower. Which best explains this discrepancy?
For the single-step reaction $NO(g) + NO_3(g) \rightarrow 2NO_2(g)$, a scientist calculates the rate of collisions but finds the observed rate much lower. Which best explains this discrepancy?
For the reaction $2NO(g) + Br_2(g) \rightarrow 2NOBr(g)$, two mechanisms are proposed. Which observation would support mechanism 1 but not mechanism 2?
For the reaction $2NO(g) + Br_2(g) \rightarrow 2NOBr(g)$, two mechanisms are proposed. Which observation would support mechanism 1 but not mechanism 2?
For a reaction with rate = $k[NO]^2[H_2]$, what happens to the rate if $[NO]$ is doubled and $[H_2]$ is halved?
For a reaction with rate = $k[NO]^2[H_2]$, what happens to the rate if $[NO]$ is doubled and $[H_2]$ is halved?
Based on the provided data for the decomposition of hydrogen peroxide, $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$, how does the rate of $O_2$ production change during the experiment?
Based on the provided data for the decomposition of hydrogen peroxide, $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$, how does the rate of $O_2$ production change during the experiment?
A student observes that without $MnO_2(s)$, the reaction is much slower. Which statement best explains this?
A student observes that without $MnO_2(s)$, the reaction is much slower. Which statement best explains this?
For the reaction $X(g) + 2Y(g) \rightarrow XY_2(g)$, which reactant is consumed more rapidly in trial 2, and why?
For the reaction $X(g) + 2Y(g) \rightarrow XY_2(g)$, which reactant is consumed more rapidly in trial 2, and why?
How can the difference between reaction path one and reaction path two for the decomposition of XY2 be best explained?
How can the difference between reaction path one and reaction path two for the decomposition of XY2 be best explained?
Consider the reaction $2X + 2Z \rightarrow X_2Z_2$. During a reaction with excess X, the plot of ln[Z] versus time is linear. What is the reaction order with respect to Z?
Consider the reaction $2X + 2Z \rightarrow X_2Z_2$. During a reaction with excess X, the plot of ln[Z] versus time is linear. What is the reaction order with respect to Z?
For the single-step reaction $NO_2Cl(g) + Cl(g) \rightarrow NO_2(g) + Cl_2(g)$, what factor will reduce the value of k, the rate constant?
For the single-step reaction $NO_2Cl(g) + Cl(g) \rightarrow NO_2(g) + Cl_2(g)$, what factor will reduce the value of k, the rate constant?
Flashcards
Rate-determining Step
Rate-determining Step
The slowest step in a reaction mechanism that determines the overall rate of the reaction.
Catalyst
Catalyst
A substance that increases the rate of a chemical reaction without being consumed in the process. Provides an alternate reaction pathway with a lower activation energy.
Activation Energy
Activation Energy
Energy required to start a chemical reaction. Catalysts lower this energy
Reaction Mechanism
Reaction Mechanism
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Rate Law
Rate Law
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Temperature Coefficient
Temperature Coefficient
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Elementary Reactions
Elementary Reactions
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Collision Orientation
Collision Orientation
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Exothermic Reaction
Exothermic Reaction
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Endothermic Reaction
Endothermic Reaction
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Heterogeneous Catalyst
Heterogeneous Catalyst
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Homogeneous Catalyst
Homogeneous Catalyst
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Reaction Order
Reaction Order
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Equilibrium
Equilibrium
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Entropy (S)
Entropy (S)
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Study Notes
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X2+Y2 --> X2Y2 has a rate of k[X2]
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Mechanisms 1 and 2 are both consistent with the rate law
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4FeS(s) + 7O2(g) --> 2Fe2O3(s) + 4SO2(g) : The combustion of some types of coal results in the formation of SO2(g) due to the presence of FeS(s).
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SO2(g) + O3(g) --> SO3(g) + O2(g): SO2(g) can react with O3(g) to form SO3(g)
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ΔH-298 = -242kJ/molrxn and ΔS-298 = -25J/(K·molrxn)
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A student studies the kinetics of X --> Products by measuring [X] over time
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Plotting ln[X] versus time and finding the magnitude of the slope allows for determination of the rate constant, k
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For C2H4 + H2 --> C2H6, adding a heterogenous catalyst to the reaction system is most likely to increase the reaction rate
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In a two-step mechanism for the destruction of ozone (O3) in the upper atmosphere the concentration is very low for the duration of the reaction
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Increasing [H+] at constant temperature increases the frequency of collisions between H+ ions and ClO- ions.
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A reaction energy profile with two steps and is exothermic best corresponds to the proposed mechanism
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For two solid reactants of varying particle size combined in a vessel a small black dots temperature of 950C will give the fastest reaction rate
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In one student's experiment the reaction proceeded at a much slower rate than it did in the other student's experiments because the student used a 1.5M solution of HNO3 instead of a 15.8 solution of HNO3.
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2 samples of Mg(s) of equal mass were placed in equal amounts of HCl(aq)
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More Mg atoms are exposed to HCl(aq) in Figure 2 than in Figure 1, therefore Figure 2's reaction will proceed faster
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With a solid nickel catalyst, the catalyst's presence results in a reaction pathway that needs lower activation energy
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2 N2O(g) --> 2 N2(g) + O2(g) with rate = k[N2O]. Increasing initial [N2O] at constant temperature increases reaction rate because molecular collisions become more frequent
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At higher temperatures, high-energy collisions happen more frequently, explaining why the rate of a chemical reaction increases with increasing temperature.
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H₂O₂(aq) --> 2H2O(l) + O2(g) with ΔH = -196 kJ/mol,m
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If the plot of ln [Hâ‚‚Oâ‚‚] versus time is a straight line, then the reaction is first order
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The collision must be successful to lead to increased reaction rate: the catalyst could adsorb one of the particles, making a successful (reaction-producing) collision with the other particle more likely
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Rate = k[H2O2]: catalyzed decomposition of H2O2 (aq)
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Trial 2's concentration of H2O2 (aq) was lower than that of Trial 1
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2 NO (g) + 2 H2 (g) --> N2 (g) + 2 H2O (g), with a rate law of k[NO]^2[H2]
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Overall order of reaction: third order
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If a 0.35 g sample of Li(s) is cut into small pieces, it will have the most likely increase the rate of reaction between Li(s) and water
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For the reaction between NO(g) + NO3(g) --> 2NO2(g),the diagram that correctly shows the orientation of collision between NO(g) and NO3(g) is most likely to be effective is (D)
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Nâ‚‚O(g) + CO(g) --> Nâ‚‚(g) + COâ‚‚(g): the rate increases significantly in the presence of Pd(s) because One of the reactants binds on the surface of Pd, which introduces an alternative reaction pathway with a lower activation energy
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For NO(g) + NO3(g) --> 2 NO2(g); rate = k[NO][NO3], the two reactant particles must collide with a particular orientation to react
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2 NO(g) + Br2(g) --> 2 NOBr(g): if the reaction rate is independent of [Br2] then that would support mechanism 1 but not mechanism 2
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2 NO(g) +2 H2(g) --> N2(g) + 2 H2O (g), Rate = k[NO]^2[H2]. If [NO] is doubled and [H2] is halved the rate doubles
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A student measures the volume of O2(g) produced by 2 H2O2(aq) -->2 H2O(l) + O2(g) at regular intervals while the temperature is held constant: the rate of O2 production decreases throughout the experiment.
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If a student observes that under the same conditions, H2O2(aq) decomposes in the presence of MnO2(s) occurs faster than without MnO2(s), it is because MnO2(s) provides a reaction pathway with a lower activation energy
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X(g) + 2Y(g) --> XY2(g): in trial 2 Y is consumed more rapidly because the rate of disappearance will be double that of X
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XY2 --> X + Y2: presence of a catalyst in path two would most likely account for the difference between reaction path one and path two.
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2 X + 2 Z --> X2Z2: With excess X, but Z is monitored over time: plot of the natural logarithm is shown, the reaction with respect to reactant Z is first order
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NOâ‚‚Cl (g) + Cl (g) --> NO2 (g) + Cl2 (g): Decreasing the temperature of the sealed container will reduce the value of k, the rate constant for the reaction.
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