CHM 132 Chapter Notes - Chapter 15: Rate Equation, Rate-Determining Step, Reversible Reaction

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silverbee937

10 Aug 2017

School

University of Rochester

Department

Chemistry

Course

CHM 132

Professor

Kara Bren

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Chemical kinetics the area of chemistry that concerns reaction rates. Rate the cha(cid:374)ge i(cid:374) a gi(cid:448)e(cid:374) (cid:395)ua(cid:374)tity o(cid:448)e(cid:396) a specific ti(cid:373)e (cid:894) []/ t(cid:895) Reaction rate the change in concentration of a reactant or product per unit of time (tends to be a positive number therefore for reactants the equation is actually - []/ t(cid:895) Rates might change but one can always find an average rate. In the above equation, one must be careful when choosing whether to find the rate of a decomposition, b production or c production. They will not all yield the same rate. Types of rate laws: differential rate law the more common one, a rate law that expresses how the rate depends on concentration. Integrated rate law expresses how the concentrations depend on time: once you know one, you know the other one. 15. 3 determining the form of the rate laws. First order is when n = 1.

Related Questions

Please answer both questions. They are part of one question in a pre lab. I have attached the introductory statement referred to in the first part. Thanks!

Did the reaction actually stop when the blue-black color appeared? Explain briefly

EXPERIMENT 16 A Kinetic Study of an Iodine Clock Reaction OBJECTIVES: To measure experimental reaction rates under different conditions. To use the method of initial rates to determine the order of reaction. To determine the rate constant of a reaction at different temperatures. To determine the activation energy of a reaction using graphical techniques. rate law. To make deductions regarding a reaction mechanism, based on the INTRODUCTION: According to chemical kinetics, it is possible to write down a Rate Law' for any chemical equation: Rate k[A]m [B] This rate law can be used to calculate the rate at which that particular reaction will proceed when a given concentration' of [A] and [B] are present in the system. However, to do this we must know the value of (it is temperature dependent). k rate constant m order with respect to concentration of A. n order with respect to concentration of B. It is important to realize that the values of k, m and n cannot be determined by looking at the chemical equation. They must be deduced experimentally someone must actually run and time the reaction under varying conditions. In this lab, we will studythe Rate Law forthe following reaction between the persulfate ion,s2os and the iodide ion. General Rate Lane. Rate -k ]m [s20g2-1n This reaction is called a clock reaction b of the thiosulfatestarch indicator system we use to time the reaction. Every you run this reaction, your flask will contain an indicator system consisting of time 5.0 x 105 moles of thiosulfate' ions, S2032, and a little starch. When the two reactants, Ose are first mixed start the reaction, every molecule of I2 that is produced reacts immediately with any nt according to the following chemical equation. S2032-preset 21- S40 12 2 S203 This removes the first 2,5 x 10 s moles I2 (look at stoichiometry of reactions above) from the system but then the thiosulfate is gone. This point is marked by the almost instantaneous appearance of an intense blue-black Also called Rate Equation and Rate Law Expression square brackets around anything, [X], means moles per liter of Xor molarin of x. You pipet 100 ml of 0.0050 M Naasio, into each flask.

A student studies the kinetics of reaction 1 below using a clock reaction technique, reaction 2, and the procedures given in this experiment. For reaction mixture 1, he measures a time to color change of 121 seconds and for reaction mixture 2, he measures a time of 58 seconds. The temperatures of the reaction mixtures were the same for both experiments.

6Iâ (aq)+ BrO3â (aq) + 6H+(aq) --> 3I2 (aq) + Brâ (aq) + 3H2O(l) (Reaction1,slow)

I2 (aq) + 2 S2O32â (aq) --> 2 Iâ (aq) + S4O62â (aq) (Reaction 2, fast)

a. Determine the change in concentration of S2O32â that occurred during the time period measured. This should be the same for both the reaction mixtures. First, use M1V1=M2V2 to determine the initial concentration of the S2O32â after mixing the contents of the two flasks but before the reaction begins. Next, find the final concentration of S2O32â by remembering that all the S2O32â is consumed when the color change occurs. You can then find the change in the concentration of S2O32â.

[S2O32â]0 =

â[S2O32â] =

b. Use the stoichiometry of Reactions 1 and 2 to determine the change in concentration of BrO3â during

the same time period. â[BrO3â] =

c. Calculate the initial rate for each reaction mixture. Recall that Rate = â â[BrOâ ] / ât

Rate for mixture 1:

Rate for mixture 2:

d. Method of Initial Rates

Complete the table below using M1V1=M2V2 to find the initial concentration of each reactant after mixing the contents of flasks I and II for both reaction mixtures. Enter the initial rates computed above.

Rxn Mixture [Iâ]0(M) [BrO3â]0(M) [H+]0 (M) Rate (M/s)

Use the method of initial rates and the relevant data above to determine the order of reaction 1 with respect to the concentration of iodide ions.

Order with respect to [Iâ]:_________________________

e. It is critical for the success of this experiment that the concentrations of the reactants Iâ, BrO3â, and H+ do not change significantly over the time period measured.

Why is this important?

f. By what percent did the bromate ion concentration change during the time period considered for mixture 1 in this experiment?

PercentChange = (|â[BrO3â ]| / [BrO3â ]0 ) x 100%

CHM 132 Chapter Notes - Chapter 15: Rate Equation, Rate-Determining Step, Reversible Reaction

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