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**Chapter 12: Chemical Kinetics**

4/11/2017 Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Rates Chemical Kinetics: The area of chemistry concerned with reaction rates and the sequence of steps by which reactions occur. Reaction Rate: Either the increase in the concentration of a product per unit time or the decrease in the concentration of a reactant per unit time. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Rates 2N2O5(g) 4NO2(g) + O2(g) N2O5 and O2 are colorless while NO2 is brown. The reaction progress may be monitored by keeping track of the reaction mixture color. decrease increase Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Rates 2N2O5(g) 4NO2(g) + O2(g) The rate of change for both NO2 and O2 is positive (increasing) and the rate of change for N2O5 is negative (decreasing). The calculated rate is dependent on the time points taken. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Rates 2N2O5(g) 4NO2(g) + O2(g) Rate of decomposition of N2O5: Dt D[N2O5] -( M M) (400 s s) = This is for a specific time period. The reaction rate is defined to be a positive quantity. s M = 1.9 x 10-5 Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Rates 2N2O5(g) 4NO2(g) + O2(g) General rate of reaction: D[N2O5] Dt D[NO2] Dt D[O2] Dt 1 2 1 4 rate = - = = The rate of reaction should be independent of whether you monitor the reactants or the products. The negative sign in front of the reactants is needed since their rate decreases. a A + b B d D + e E D[A] Dt D[B] Dt D[D] Dt D[E] Dt 1 a 1 b 1 d 1 e rate = - = - = = Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Rates 2N2O5(g) 4NO2(g) + O2(g) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws and Reaction Order**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws and Reaction Order Rate Law: An equation that shows the dependence of the reaction rate on the concentration of each reactant. aA + bB products rate [A]m[B]n a rate = k[A]m[B]n k is the rate constant Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws and Reaction Order**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws and Reaction Order The values of the exponents in the rate law must be determined by experiment; they cannot be deduced from the stoichiometry of the reaction. Rate laws are determined only from experiment and may not be written directly from the coefficients. k is the rate constant. The reaction order with respect to A is m, with respect to B is n, and the overall reaction order is m+n. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Experimental Determination of a Rate Law**

Chapter 12: Chemical Kinetics 4/11/2017 Experimental Determination of a Rate Law 2NO(g) + O2(g) 2NO2(g) [O2]n rate = k[NO]m Compare the initial rates to the changes in initial concentrations. One reactant concentration changes while the others remain constant. You can also use logarithm method for more complicated problems. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Experimental Determination of a Rate Law**

Chapter 12: Chemical Kinetics 4/11/2017 Experimental Determination of a Rate Law 2NO(g) + O2(g) 2NO2(g) [O2]n rate = k[NO]2 The concentration of NO doubles, the concentration of O2 remains constant, and the rate quadruples. 2m = 4 m = 2 One reactant concentration changes while the others remain constant. You can also use the logarithm method for more complicated problems. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Experimental Determination of a Rate Law**

Chapter 12: Chemical Kinetics 4/11/2017 Experimental Determination of a Rate Law 2NO(g) + O2(g) 2NO2(g) [O2] rate = k[NO]2 The concentration of O2 doubles, the concentration of NO remains constant, and the rate doubles. 2n = 2 n = 1 Copyright © 2008 Pearson Prentice Hall, Inc.

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**Experimental Determination of a Rate Law**

Chapter 12: Chemical Kinetics 4/11/2017 Experimental Determination of a Rate Law 2NO(g) + O2(g) 2NO2(g) [O2] rate = k[NO]2 Reaction Order With Respect to a Reactant NO: second-order O2: first-order Overall Reaction Order 2 + 1 = 3 (third-order) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Experimental Determination of a Rate Law**

Chapter 12: Chemical Kinetics 4/11/2017 Experimental Determination of a Rate Law 2NO(g) + O2(g) 2NO2(g) [O2] rate = k[NO]2 Units for this third-order reaction: M s rate [NO]2 [O2] 1 M2 s k = = = (M2) (M) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Experimental Determination of a Rate Law**

Chapter 12: Chemical Kinetics 4/11/2017 Experimental Determination of a Rate Law 2NO(g) + O2(g) 2NO2(g) [O2] rate = k[NO]2 Overall order of this reaction is third. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Integrated Rate Law for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Integrated Rate Law for a First-Order Reaction A product(s) D[A] Dt - = k[A] rate = k[A] Calculus can be used to derive an integrated rate law. [A]t concentration of A at time t [A]0 initial concentration of A [A]t [A]0 ln = -kt x y ln = ln(x) - ln(y) Using: ln[A]t = -kt + ln[A]0 y = mx + b Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Integrated Rate Law for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Integrated Rate Law for a First-Order Reaction ln[A]t = -kt + ln[A]0 y = mx + b A plot of ln[A] versus time gives a straight-line fit and the slope will be -k. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Integrated Rate Law for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Integrated Rate Law for a First-Order Reaction ln[A]t = -kt + ln[A]0 This is a plot of [A] versus time. The best-fit is a curve and not a line. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Integrated Rate Law for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Integrated Rate Law for a First-Order Reaction ln[A]t = -kt + ln[A]0 The plot on the left is a plot of concentration versus time which does not yield a straight-line fit for a first-order reaction (it would for a zeroth-order reaction). The one on the right does and the slope is equal to -k. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Integrated Rate Law for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Integrated Rate Law for a First-Order Reaction 2N2O5(g) 4NO2(g) + O2(g) rate = k[N2O5] Slope = -k Copyright © 2008 Pearson Prentice Hall, Inc.

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**Integrated Rate Law for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Integrated Rate Law for a First-Order Reaction 2N2O5(g) 4NO2(g) + O2(g) rate = k[N2O5] Calculate the slope: Slope = -k ( ) s (-3.912) s 1 = s 1 k = Copyright © 2008 Pearson Prentice Hall, Inc.

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**Half-Life for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Half-Life for a First-Order Reaction Half-Life: The time required for the reactant concentration to drop to one-half of its initial value. A product(s) rate = k[A] t = t1/2 [A]t [A]0 ln = -kt = t1/2 [A] 2 [A]0 = -kt1/2 1 2 ln t1/2 = k 0.693 or Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Half-Life for a First-Order Reaction**

Chapter 12: Chemical Kinetics 4/11/2017 Half-Life for a First-Order Reaction t1/2 = k 0.693 For a first-order reaction, the half-life is independent of the initial concentration. Each successive half-life is an equal period of time. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions A product(s) D[A] Dt - = k[A]2 rate = k[A]2 Calculus can be used to derive an integrated rate law. [A]t concentration of A at time t [A]0 initial concentration of A = kt + [A]0 1 [A]t Simplest case of one reactant. y = mx + b Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions 2NO2(g) 2NO(g) + O2(g) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions 2NO2(g) 2NO(g) + O2(g) Plotting ln[NO2] versus time gives a curve and not a straight-line fit. Therefore, this is not a first-order reaction. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions 2NO2(g) 2NO(g) + O2(g) [NO2] 1 Plotting versus time gives a straight-line fit. Therefore, this is a second-order reaction. Slope = k Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions 2NO2(g) 2NO(g) + O2(g) Calculate the slope: Slope = k M 1 ( ) M s 1 = 0.540 ( ) s M s 1 k = 0.540 Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions Half-life for a second-order reaction A product(s) rate = k[A]2 t = t1/2 = kt + [A]0 1 [A]t = t1/2 [A] 2 [A]0 [A]0 1 = kt1/2 + 2 = t1/2 k[A]0 1 Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Second-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Second-Order Reactions = t1/2 k[A]0 1 For a second-order reaction, the half-life is dependent on the initial concentration. Each successive half-life is twice as long as the preceding one. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Copyright © 2008 Pearson Prentice Hall, Inc.

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**Zeroth-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Zeroth-Order Reactions For a zeroth-order reaction, the rate is independent of the concentration of the reactant. A product(s) D[A] Dt - = k rate = k[A]0 = k Calculus can be used to derive an integrated rate law. [A]t concentration of A at time t [A]0 initial concentration of A [A]t = -kt + [A]0 y = mx + b Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Zeroth-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Zeroth-Order Reactions A plot of [A] versus time gives a straight-line fit and the slope will be -k. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Zeroth-Order Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Zeroth-Order Reactions rate = k[NH3]0 = k Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Mechanisms Reaction Mechanism: A sequence of reaction steps that describes the pathway from reactants to products. Elementary Reaction (step): A single step in a reaction mechanism. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Mechanisms Experimental evidence suggests that the reaction between NO2 and CO takes place by a two-step mechanism: NO2(g) + NO2(g) NO(g) + NO3(g) elementary reaction NO3(g) + CO(g) NO2(g) + CO2(g) elementary reaction NO2(g) + CO(g) NO(g) + CO2(g) overall reaction An elementary reaction describes an individual molecular event. The overall reaction describes the reaction stoichiometry and is a summation of the elementary reactions. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Mechanisms NO2(g) + NO2(g) NO(g) + NO3(g) NO3(g) + CO(g) NO2(g) + CO2(g) Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Reaction Mechanisms Experimental evidence suggests that the reaction between NO2 and CO takes place by a two-step mechanism: NO2(g) + NO2(g) NO(g) + NO3(g) elementary reaction NO3(g) + CO(g) NO2(g) + CO2(g) elementary reaction NO2(g) + CO(g) NO(g) + CO2(g) overall reaction A reactive intermediate is formed in one step and consumed in a subsequent step. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Reaction Mechanisms-Molecularity**

Chapter 12: Chemical Kinetics 4/11/2017 Reaction Mechanisms-Molecularity Molecularity: A classification of an elementary reaction based on the number of molecules (or atoms) on the reactant side of the chemical equation. unimolecular reaction: O3*(g) O2(g) + O(g) bimolecular reaction: O3(g) + O(g) 2 O2(g) termolecular reaction: O(g) + O(g) + M(g) O2(g) + M(g) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Elementary Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Elementary Reactions The rate law for an elementary reaction follows directly from its molecularity because an elementary reaction is an individual molecular event. unimolecular reaction: O3*(g) O2(g) + O(g) rate = k[O3] bimolecular reaction: O3(g) + O(g) 2 O2(g) rate = k[O3][O2] termolecular reaction: O(g) + O(g) + M(g) O2(g) + M(g) rate = k[O]2[M] Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Elementary Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Elementary Reactions Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Overall Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Overall Reactions Rate-Determining Step: The slow step in a reaction mechanism since it acts as a bottleneck and limits the rate at which reactants can be converted to products. Typical problems are either determining the rate law from a given reaction mechanism and verifying that a mechanism is plausible. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Overall Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Overall Reactions Initial Slow Step k1 NO2(g) + NO2(g) NO(g) + NO3(g) slow step k2 NO3(g) + CO(g) NO2(g) + CO2(g) fast step NO2(g) + CO(g) NO(g) + CO2(g) overall reaction Based on the slow step: rate = k1[NO2]2 Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Overall Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Overall Reactions Initial Fast Step k1 2NO(g) N2O2(g) fast step, reversible k-1 k2 N2O2(g) + H2(g) N2O(g) + H2O(g) slow step k3 N2O(g) + H2(g) N2(g) + H2O(g) fast step Recall that the rate law is written in terms of reactants. Reactive intermediates are not included in the rate law since they have such short lifetimes. N2O2 is an intermediate and we must rewrite the rate law without the intermediate. 2NO(g) + 2H2(g) N2(g) + 2H2O(g) overall reaction Based on the slow step: rate = k2[N2O2][H2] Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Overall Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Overall Reactions rate = k2[N2O2][H2] intermediate First step: Rateforward = k1[NO]2 Ratereverse = k-1[N2O2] k1[NO]2 = k-1[N2O2] The first step is reversible and thus the forward and reverse rates are equal. [NO]2 [N2O2] = k-1 k1 k-1 k1 Slow step: rate = k2[N2O2][H2] rate = k2 [NO]2[H2] Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Rate Laws for Overall Reactions**

Chapter 12: Chemical Kinetics 4/11/2017 Rate Laws for Overall Reactions Procedure for Studying Reaction Mechanisms Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**The Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 The Arrhenius Equation Typically, as the temperature increases, the rate of reaction increases. 2N2O5(g) 4NO2(g) + O2(g) rate = k[N2O5] The rate constant is dependent on temperature. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**The Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 The Arrhenius Equation Collision Theory: As the average kinetic energy increases, the average molecular speed increases, and thus the collision rate increases. Copyright © 2008 Pearson Prentice Hall, Inc.

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**The Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 The Arrhenius Equation Activation Energy (Ea): The minimum energy needed for reaction. As the temperature increases, the fraction of collisions with sufficient energy to react increases. Copyright © 2008 Pearson Prentice Hall, Inc.

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**The Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 The Arrhenius Equation Transition State: The configuration of atoms at the maximum in the potential energy profile. This is also called the activated complex. The activation energy is the height of the barrier. Copyright © 2008 Pearson Prentice Hall, Inc.

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**The Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 The Arrhenius Equation k = Ae-E /RT a k rate constant A collision frequency factor Ea activation energy R gas constant T temperature (K) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Using the Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 Using the Arrhenius Equation ln(k) = ln(A) + ln(e-E /RT) a + ln(A) T 1 R -Ea ln(k) = RT Ea ln(k) = ln(A) - rearrange the equation y = mx b Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Using the Arrhenius Equation**

Chapter 12: Chemical Kinetics 4/11/2017 Using the Arrhenius Equation + ln(A) T 1 R -Ea ln(k) = R -Ea Slope = T 1 Watch the units for the activation energy and for R. Plot ln(k) versus Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Catalysis Catalyst: A substance that increases the rate of a reaction without itself being consumed in the reaction. A catalyst is used in one step and regenerated in a later step. rate-determining step H2O2(aq) + I1-(aq) H2O(l) + IO1-(aq) H2O2(aq) + IO1-(aq) H2O(l) + O2(g) + I1-(aq) fast step 2H2O2(aq) 2H2O(l) + O2(g) overall reaction Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Catalysis Since the catalyst is involved in the rate determining step, it often appears in the rate law. rate = k[H2O2][I1-] rate-determining step H2O2(aq) + I1-(aq) H2O(l) + IO1-(aq) H2O2(aq) + IO1-(aq) H2O(l) + O2(g) + I1-(aq) fast step 2H2O2(aq) 2H2O(l) + O2(g) overall reaction Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Chapter 12: Chemical Kinetics**

4/11/2017 Catalysis Note that the presence of a catalyst does not affect the energy difference between the reactants and the products Copyright © 2008 Pearson Prentice Hall, Inc.

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**Homogeneous and Heterogeneous Catalysts**

Chapter 12: Chemical Kinetics 4/11/2017 Homogeneous and Heterogeneous Catalysts Homogeneous Catalyst: A catalyst that exists in the same phase as the reactants. Heterogeneous Catalyst: A catalyst that exists in a different phase from that of the reactants. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Homogeneous and Heterogeneous Catalysts**

Chapter 12: Chemical Kinetics 4/11/2017 Homogeneous and Heterogeneous Catalysts Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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**Homogeneous and Heterogeneous Catalysts**

Chapter 12: Chemical Kinetics 4/11/2017 Homogeneous and Heterogeneous Catalysts Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

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Chapter 12 Chemical Kinetics.

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