SECTION 1. THE REACTION PROCESS CHAPTER 17. REACTION KINETICS SECTION 1. THE REACTION PROCESS “Kinetics” refers to motion. Kinetic energy = energy of motion Reaction kinetics (chemical kinetics) is the study of how fast chemical reactions go (reaction rates).

The Reaction Process Some reactions happen essentially as fast as the reactants come together. They are limited by the rate of diffusion. For example, an acid-base reaction, mixing solutions of HCl and NaOH:

H3O+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → 2H2O(l) + Na+(aq) + Cl-(aq) The reaction occurs as soon as H3O+ and OH- diffuse to be near each other. We will be focusing on reactions in which one or more other steps are required in addition to the reactants simply coming together.

Reaction Mechanisms The reaction mechanism is the step-by-step sequence of reactions by which the overall chemical change occurs. It includes chemical species that do not appear in the overall reaction. These are called intermediates.

Overall reaction: H2(g) + I2(g) → 2HI Reaction mechanism: Example: Overall reaction: H2(g) + I2(g) → 2HI Reaction mechanism: step 1: I2 → 2I step 2: I + H2 → H2I step 3: H2I + I → 2HI ___________________________________________________ overall I2 + H2 → 2HI Iodine atoms (I) and H2I are intermediates.

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What happens in many reactions is described by collision theory. In order for reactions to occur between substances, their particles must collide. What happens in many reactions is described by collision theory. Not every collision brings about a reaction. There are two requirements: favorable orientation and enough energy.

Particle Collisions

1. Orientation. Ex. – NO + Cl2 → NOCl + Cl (structure: O=N-Cl) The chlorine must hit the nitrogen end of the molecule for there to be a reaction: reaction no reaction

2. Energy: there must be enough energy to break existing bonds before new bonds can form. Ex. – 2H2 + O2 → 2H2O ∆H is negative, but first H-H and O-O bonds must break. ∆H for this is positive.

Activation Energy A transitional structure that results from an effective collision, and that persists while old bonds are breaking and new ones are forming, is called an activated complex. The minimum energy needed to produce the activated complex is the activation energy.

Energy diagram: activated complex reactants products activation energy = Ea reactants energy ∆E products __________________________________________ course of reaction

An intermediate would appear between two peaks of energy: reactants ∆E products __________________________________________ course of reaction

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Therefore reaction rates increase with increasing temperature. Effect of temperature As temperature increases, particles have more kinetic energy, so it is more likely that they will collide with energy equal to or greater than the activation energy. Therefore reaction rates increase with increasing temperature.

SECTION 2. REACTION RATE The change in concentration of reactants per unit time as a reaction proceeds is called the reaction rate. Homogenous reaction: reactants and products are in a single phase (usually liquid or gas) Heterogeneous reaction: involves reactants in two different phases

Rate-Influencing Factors Nature of Reactants: The rate of reaction depends on the particular reactants and bonds involved. Surface Area: For a heterogeneous reaction, the phases must come in contact. For a solid, only particles at the surface can react. So increasing the surface area (such as by cutting into pieces or making a powder) increases the rate.

3. Temperature (already discussed) 4. Concentration. The frequency of effective collisions will increase if the concentration of one or more reactants is increased. 5. Presence of a catalyst. A catalyst increases the rate of a chemical reaction without being consumed.

A catalyst provides an alternative mechanism with a lower activation energy, so reactants are more likely to have enough energy to react. Enzymes are biological catalysts. Metals are often used as catalysts.

How does ∆E compare for the overall reaction? as-bio-and-chem.blogspot.com

How does ∆E compare for the overall reaction? no effect as-bio-and-chem.blogspot.com

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Rate Law: an equation that relates reaction rates and concentrations of reactants. It applies to a specific reaction at a specific temperature. It must be determined experimentally. Usually this is done by changing the concentration of one reactant while keeping the concentration of all others constant, and measuring how the reaction rate changes.

The general form for the rate law is given by the following equation: R = k[A]n[B]m where R = reaction rate; k = a constant; [A] and [B] are molar concentrations of reactants A and B; and n and m are usually small integers (possibly zero).

The power to which a reactant is raised is called the order in that reactant. Ex. – if n = 2, the reaction is 2nd order in reactant A. An order of 0 means that the rate does not depend on the concentration, as long as some is present.

The sum of all the exponents is the order of the reaction or overall order. 2H2 + 2NO → N2 + 2H2O If you keep [NO] constant and vary [H2], it is found that: R  [H2] If you keep [H2] constant and vary [NO], it is found that: R  [NO]2.

So the rate law is: R = k[H2][NO]2 n = 1; it is 1st order in H2 m = 2; it is 2nd order in NO overall, n + m = 3; it is a 3rd order reaction

Other examples (a and b are reactions in smog formation (p. 573)): a. 3NO → N2O + NO2 R = k[NO]2 Order in NO = Overall order = b. NO2 + CO → NO + CO2 R = k[NO2]2 Order in NO2 = Order in CO = Overall order =

Other examples (a and b are reactions in smog formation (p. 573)): a. 3NO → N2O + NO2 R = k[NO]2 Order in NO = 2 Overall order = 2 b. NO2 + CO → NO + CO2 R = k[NO2]2 Order in NO2 = 2 Order in CO = 0 Overall order = 2

c. NH4+ + NO2- → N2 + 2H2O R = k[NH4+][NO2-] Order in NH4+ = Order in NO2- = Overall order = d. BrO3- + 5Br- + 6H+ → 3Br2 + 3H2O R = k[BrO3-][Br-][H+]2 Order in BrO3- = Order in Br- = Order in H+ =

c. NH4+ + NO2- → N2 + 2H2O R = k[NH3+][NO2-] Order in NH3+ = 1 Order in NO2- = 1 Overall order = 2 d. BrO3- + 5Br- + 6H+ → 3Br2 + 3H2O R = k[BrO3-][Br-][H+]2 Order in BrO3- = 1 Order in Br- = 1 Order in H+ = 2 Overall order = 4

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The orders in the rate law may or may not match the coefficients in the balanced equation. Determining k: once the order of reactants have been determined, the rate law can be rearranged: and k can be calculated from the experimental value of R at any concentration of A and B.

Example of rate law calculation (Sample Problem C, p. 575) For the reaction A + B → C rates were determined at three concentrations of A and B: Experiment [A] (M) [B] (M) Rate (M/s) 1 1.2 2.4 8.0 x 10-8 2 4.0 x 10-8 3 3.6 7.2 x 10-7

To determine the order in A: For which two experiments does [A] change and [B] remains constant? 1 and 3 How much does [A] increase? 3 times How much does the rate increase?

Since the concentration went up 3 times and the rate went up 9 = 32 times, the order in A is 2. To determine the order in B: For which two experiments does [B] change and [A] remains constant? 2 and 1 How much does [B] increase? 2 times

How much does the rate increase? Since the concentration went up 2 times and the rate went up 2 = 21 times, the order in B is 1. So the rate law is: R = k [A]2 [B]

To get k, rearrange the rate law and substitute the values from any experiment. Using the values from experiment 1:

The form of the rate law depends on the reaction mechanism. Rate Laws and Reaction Pathway The form of the rate law depends on the reaction mechanism. For a reaction that occurs in a single step, the reaction rate of that step is proportional to the products of the concentrations of the reacting particles.

Ex. = A and B collide to make two particles of C: A + B → 2C R = k [A] [B] In reverse: two particles of C collide to make A + B: 2C → A + B R = k’ [C] [C] = k’ [C]2

For a reaction occurring in more than one step, often one step is much slower than others. This will determine the overall rate; it is called the rate-determining step. The rate law will depend on what happens in that step.

Ex.: NO2 + CO → NO + CO2 The mechanism is thought to be: 1. NO2 + NO2 → NO3 + NO (slow) 2. NO3 + CO → NO2 + CO2 (fast) Since the rate is determined by step 1, the rate law is: R = k [NO2]2

The reaction is zero order in CO, because CO only participates in step 2. Adding more CO will not make the reaction go any faster.