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CHM 11600 Lecture 2: Methods of Rates

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cerisegiraffe538

25 Mar 2015

School

Purdue University

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Course

Professor

Nash/ Wasserman

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Chm 11600 - general chemistry - lecture 2_methods of rates. Method of instantaneous rates: general form of rate law. Evaluate relationships between rate/concentration and to determine order. If [h2o2] is doubled and rate is doubled, k[h2o2]1. If [h2o2] is quadrupled and rate is quadrupled, k[h2o2]2. If [h2o2] is increased by a factor of 8 and rate is increased by a factor of 8, k[h2o2]3. 2. 31*10-5 m/s = k[0. 500m] k = 4. 62*10-5s-1 = kavg. Relationship between rate of disappearance of reactants and rates of appearance of products: [h2o2] / t = [h2o] / t = 2(+ [o2] / t ) Perform several experiments with different concentrations of reactants. Advantage: method works for reversible chemical reactions. 0. 050 aa + bb dd + ee. Rate = -(1/a) [a] / t = -(1/b) [b] / t =(1/c) [c] / t =(1/e) [e] / t. Rate law (only includes reactants) ie. a b + c.

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Related Questions

Urea (NH₂CONH₂) is the end product in protein metabolism in animals. The decomposition of urea in 0.1 M HCl occurs according to the reaction

NH₂CONH₂(aq) + H⁺(aq) + 2 H₂O(l) 2 NH₄⁺(aq) + HCO₃^-(aq)

The reaction is first order in urea and first order overall. When [NH₂CONH₂] = 0.200 M, the rate at 61.05 °C is 8.56 × 10^-5 M/s. (a) What is the rate constant, k? (b) What is the concentration of urea in this solution after 4.00 × 10³ s if the starting concentration is 0.500 M? (c) What is the halflife for this reaction at 61.05 °C?

b) The integrated rate law for the first order reaction

can be used for calculation of the amount of substrate remained after any time period.

In this case,

c) The half-life for the first order reaction is

periwinklegerbil478

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.

For a reaction aA rightarrow Products, [A]_o = 13 M, and the first two half-lives are 35 and 70 minutes, respectively, a) Write the rate law for this reaction. b) Calculate k. c) Calculate [A] at t = 335 minutes The Model of Chemical Kinetics represents the reaction rate constant as: k = (2p) times exp(-Ea/RT). What does each variable in the equation mean- that is, in regard to their contribution towards the reaction rate constant? In other words, what does each variable actually mean/refer to? Chemists commonly use a rule of thumb that an increase of 10 K in temperature doubles the rate of a reaction. What must the activation energy be for this statement to be true for a temperature increase from 63 degree C to 73 degree C? ln(k_2/k_1) = E_a/R(1/T_1 - 1/T_2) (i) What is a catalyst? (ii) What are the two types of catalysts? Explain specifically why they are different? (iii) What is one example of each type of catalyst? (i) What does it mean when a reaction is in chemical equilibrium? Specifically, within the context of reactants? products, and the forward and reverse reaction rates, what does chemical equilibrium mean? (ii) What is Le Chateleir's Principle? Within the context of reactant/product concentration, temperature, and pressure, how does Le Chatelier's Principle predict a reaction will respond to variations in reactant, product concentration, temperature, and pressure. (i) For the given reactions: N_2(g) + 3H_2(g) harr 2NH_3(g), H_2O(g) + CO(g) harr H_2(g) + CO_2(g), HCN(aq) harr H^+(aq) + CN^-(aq), 2KClO3(s) harr 2KCl(s) + 3O_2(g) and 2H_2 O(l) harr 2H_2(g) + O_2(g). Write the equation that governs each of their equilibrium rate constant K (ii) If N_2(g) + 3H_2(g) harr 2NH_3(g)'s (at 25 degree C) has K = 25, what would K be (at 25 degree C) for 2NH_3(g) harr N_2(g) + 3H_2(g)? (iii) What would K be if [5 times (N_2(g)+3H_2(g) harr 2NH_3(g)]? (i) For N_2(g) + 3H_2(g) harr 2NH_3(g): P_NH3: 1.9 times 10^-1 atm, P_N2 = 8.9 times 10^-2 atm, P_H2 = 1.9 times 10^-2 atm, what is K_p? (ii) Using K_p, calculated in the previous step, what is K at 45 degree C, where K_p = K(RT)^Delta n? (R = 0.08206 L atm/mol K). What is the difference between homogeneous and heterogeneous? (ii) If K is larger than 1, does the equilibrium life farther to the right (favoring the production of products) or to the left (favoring reactant formation)? (iii) If K is less than 1, does the equilibrium life farther to the right (favoring the production of products) or to the left (favoring the production of reactants)?