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Chemistry Chapter 13

by: Devin Mart

Chemistry Chapter 13 CHM 170 002

Devin Mart
GPA 3.82

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Chemical Kinetics
General Chemistry 2
Dr. Richter
Chemistry 170 - General Chemistry 2
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This 4 page Bundle was uploaded by Devin Mart on Tuesday March 29, 2016. The Bundle belongs to CHM 170 002 at Missouri State University taught by Dr. Richter in Spring 2016. Since its upload, it has received 9 views. For similar materials see General Chemistry 2 in Chemistry at Missouri State University.


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Date Created: 03/29/16
  Mart 1  Chapter 13: Chemical Kinetics    ● Chemical kinetics studies how the molecular world changes with time.  ○ Reactions are predictions, they do not tell us what will happen, but what should or  might happen.  ● Chemical kinetics can be defined as the search for answers to the following questions:  1. What is the rate at which the reactants are converted into the products of a  reaction?  2. What factors influence the rate of reaction?  3. What is sequence of steps, or the mechanism, by which the reactants are  converted into products?  ● The term ​ rate is used to describe the change in a quantity that occurs per unit of time.  ○ Δ = change in  ○ [ x ] = any time brackets are used around a number the measurement is in  molarity.  ■ Speed = change in distance / change in time  ■ Weight loss = change in weight /  change in time  ○ The rate of a chemical reaction is a measure of how fast the reaction occurs.  ○ The rate of a chemical reaction is the change in the concentration of one of the  reactants or products that occurs during a given period of time, Δt.  ■ Rate = Δ[concentration] / Δt  ■ Rate = + (Δ[product] / Δt) = ­ (Δ[reactant] / Δt)  ○ The calculated rate is dependent on the time points taken.  ○ The rate is always positive, there can never be a negative rate.  ○ The ​ average reaction​ rate is the change in measured concentrations in any  particular time period.  ○ The ​ instantaneous rate​ is the change in concentration at any one particular time.  ● Reaction rates for at least one reactant or product must be measured ​ experimentally​.  ○ Spectroscopy can be used to measure concentration changes if color changes  occur.  ○ Pressure can be used to measure concentrations of the number of moles of gas  changes.  ● The rate (as we’ve seen) of a reaction often depends on the concentration of one or more  of the reactants.  ● The Rate Law is a mathematical relationship between the rate of the reaction and the  concentration of reactants.  ○ Rate = k[A]ⁿ  ○ The rate of a reaction is directly proportional to the concentration of each reactant  raised to a power.    Mart 2  ■ Rate = k[A]ⁿ[B]ⁿ  ● K is the rate constant while the A and B represent concentrations  of reactants.  ● The values of n and m determine how the rate depends on the concentration and the  reactant.  ○ If n (orm​) = 0, the reaction is zero order and the rate is independent of the  concentration of A (or B).  ○ If n (orm​) =1, the reaction is first order and the rate is directly proportional to the  concentration of A (or B).  ○ If n (orm​) =2, the reaction is second order and the rate is proportional to the  square of the concentration of A (or B).  ● The order of a reactant is not related to the stoichiometric coefficient of the reactant in the  balanced chemical equation.  ● Rate Laws:  ○ Rate laws are always determined experimentally, they cannot be determined  theoretically.  ○ Reaction order is always defined in terms of reactant (not product) concentrations.  ○ The order of a reactant is not relation to the stoichiometric coefficient of the  reactant in the balanced chemical equation.  ● One method for experimentally determining the order of a reaction is the method of  initial rates.  ○ The initial rate ­ the rate for a short period of time at the beginning of the reaction  ­ this is measured by running the reaction several times, each time varying the  concentration of only one reactant and measure its “initial” rate.  ○ The resulting change in rate indicates the order with respect to that reactant.  ○ If you are unsure about the reaction order using inspection, set up a ratio and  solve it mathematically.  ● When more than one reactant is present, the concentration of each is varied independently  of the other.  ○ Between different phases of the reaction only one of the reactants concentrations  will change at a time.  ● The units of the rate constant (k) vary depending on the order of the reaction.  ○ 0 reaction order = M*s (or mol L¹s⁻¹)  ○ 1st reaction order = 1/s (or¹)  ○ 2nd reaction order = M⁻¹*s⁻¹ (or L mol¹s⁻¹)  ○ 3rd reaction order = L²/mol²*s (or L²mo²s⁻¹)  ● The integrated rate law is a relationship between the concentrations of reactants and time.  ○ For a first order reaction, a plot of the natural log of the reactant concentration as  a function of time yields a straight line.    Mart 3  ○ For a second order reaction, a plot of the inverse of the reactant concentration as a  function of time yields a straight line.  ■ When a second order reaction is plotted using first­order parameters, a  straight line is not observed.  ○ For a zero order reaction, a plot of the reactant concentration as a function of time  yields a straight line.  ● The ​ half­lif (t.) of a reaction is the time required for the concentration of a reactant to  fall to one­half of its initial value.  ○ The half­life expression defines the dependence of half­life on the rate constant  and the initial concentration.  ○ Even though the concentration is changing as the reaction proceeds, the half­life  is constant.  ● The rates of chemical reaction are, in general, highly sensitive to temperature.  ○ The temperature dependence of the reaction rate is contained in the rate  “constant”.  ● The activation energy is the energy necessary to initiate a reaction.  ○ Not all chemical molecules in a container have the same kinetic energy.  ● The frequency factor is the how many times the reactants approach the activation barrier  per unit time.  ○ The frequency factor is the number of times the two molecules collide in the  correct orientation.  ○ The frequency factor and activation energy are determined experimentally using  an Arrhenius Plot.  ● The Activation Energy can also be determined if you know the rate constant at two  different temperatures.  ● Not all reactions occur in a single step.  ○ The second and third reactions are much faster than the first.  ● The rate laws for chemical reactions can be explained by the following general rules.  ○ 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 products.  ○ The overall rate law for a reaction is dominated by the rate law for the slowest  step in the reaction.  ● The rate limiting step in a reaction mechanism limits the overall rate of the reaction.  ○ Although the rate law for an overall chemical reaction cannot be determined from  the balanced equation, the rate law for an elementary step can be.  ● For a reaction mechanism to be valid, the following conditions must be met:  1. The elementary steps in the mechanism must sum to the overall reaction.    Mart 4  2. The rate law predicted by the mechanism must be consistent with the  experimentally determined rate law.  ● A catalyst increases the rate of a reaction by providing an alternative mechanism that has  a smaller activation energy.  ○ Catalysts lower the energy of activation, lowering the activation energy increases  the rate constant, k, and thereby increases the rate of the reaction.  ■ A catalyst increases the activation energy by providing a different  mechanism for the reaction through a new, lower energy pathway.  ○ A catalyst increases the rate of the forward and the reverse reactions.  ○ A catalyst is not consumed by the reaction, typically, small amounts of catalysts  affect the rate of reaction for a large amount of products.  ● Catalytic reactions also occur in the atmosphere.  ○ Significant ozone depletion has been observed around the globe, especially in the  Antarctic.  ● Catalytic converts also involve catalysis.  ○ Homogeneous catalysts exist in the same phase as the reactants, while  heterogeneous catalysts exist in a different phase than the reactants.     


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