Chapter 16

 The time taken for the disappearance of the reactant or the appearance of the product. Rate is a ratio as the amount of reactant disappeared divided by the time.  Average rate: The change in the concentration divided by the total time elapsed.  Rate = amount reacted or produced/ time interval  units: g/s, mol/s, or %/s

 Instantaneous rate: rate measured between very short interval Initial rate: instantaneous rate at the beginning of an experiment

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 Rate depends on the concentration:  In some reactions doubling the concentration doubles the rate of reaction. In some doubling the reaction increases the reaction four folds.This happens in the decomposition of HI to form H 2 and I 2.

 An expression for the rate of a reaction as a function of the concentration of one or more of the reactants.  Rate=k [A] n  This equation is the general rate law. The exponent n, is called the order with respect to substance A and must be determined from experimental data.

 Order of a chemical reaction can be said as the exponent on the concentration for a specified reactant in a rate law expression.

 Determine the rate law equation for the following reaction, given the experimental data  3A  C  Concentration of A Reaction rate  0.2M1.0M/s  0.4M4.0M/s

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 The rate law equations you have looked so far have been for reactions involving only one reactant.  If more than one reactant is found to contribute to the rate of the reaction, then all contributing reactants must appear in the rate law.  The rate law equation for this will be  rate=k [A] n [B] m

 The value of n is the order with respect to reactant A. The value of m is the order with respect to reactant B. The overall reaction order will be the sum of n and m.  From the above equation if you double the concentration of A and the rate doubles then the reaction is first order with respect to A. If you double the conc. of B (keeping the conc of A constant, and the rate quadruples the the rate of the reaction is second order with respect to B.

 For the reaction A and B for this example the rate law would be rate=k[A][B] 2  Rate laws cannot be derived from a chemical equation.  2N 2 O₅↔4NO 2 +O 2  Keq=[NO 2 ] 4 [O 2 ]/[N 2 O₅] 2  rate=k[N 2 O₅] 1

 The slowest step in a mechanism, the step that determines the overall rate of reaction is the rate determining step.  Mechanism is a proposed sequence of steps that describes how reactants are changed into products.  Each step in the mechanism is called as elementary step.

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 Temperature: An increase in temperature is accompanied by an increase in the reaction rate. Temperature is a measure of the kinetic energy of a system, so higher temperature means higher average kinetic energy of molecules and more collisions per unit time.  For most chemical reactions the rate at which the reaction proceeds will approximately double for each 10°C increase in temperature. Once the temperature reaches a certain point, some of the chemical species may be altered (e.g., denaturing of proteins) and the chemical reaction will slow or stop.

 Concentration: A higher concentration of reactants leads to more effective collisions per unit time, which leads to an increasing reaction rate (except for zero order reactions). Similarly, a higher concentration of products tends to be associated with a lower reaction rate.

 Medium: The rate of a chemical reaction depends on the medium in which the reaction occurs. It sometimes could make a difference whether a medium is aqueous or organic; polar or nonpolar; or liquid, solid, or gaseous.

 Surface area: It is easier to dissolve sugar if it is crushed. Crushing the sugar increases its surface tension.The larger surface area allows more sugar molecules to contact the solution.

 Catalyst: A catalyst is a substance that alters the rate of a chemical reaction without being used up or permanently changed chemically.  A catalyst works by changing the energy pathway for a chemical reaction. It provides an alternative route (mechanism) that lowers the Activation Energy meaning more particles now have the required energy needed to undergo a successful collision.

 What is activation energy?  The least amount of energy needed to permit a particular chemical reaction.

 There are 2 types of catalysts:  Homogeneous catalyst: Homogeneous catalysts are in the same phase as the reactants.  Heterogeneous catalyst: Heterogeneous catalysts are present in different phases from the reactants (for example, a solid catalyst in a liquid reaction mixture), whereas homogeneous catalysts are in the same phase (for example, a dissolved catalyst in a liquid reaction mixture).

 Example of Homogeneous catalyst  2H 2 O 2 (aq)+ KI(aq)  2H 2 O(l)+O 2 (g)  Example of Heterogeneous catalyst  Decomposition of H 2 O 2 in presence of MnO 2  Hydrogen peroxide is a solution while manganese dioxide is a solid and can be easily separated.

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