Chemical Kinetics

Collision Theory of Reactions Collision theory is simple - for a reaction to occur, particles must collide successfully! A successful collision has two components: particles must collide at the right geometry particles must collide with enough energy (the Activation energy) Factors which influence the geometry or energy of a collision will affect the rate of the reaction!

Collision Theory: Geometry A successful collision between molecules requires the correct geometry to break bonds and form products

4 Factors that Affect Rates Concentration Greater concentration of reactants means there will be more collisions and therefore faster rates! Temperature Higher temperatures means particles collide with greater kinetic energy, increasing the rate of reaction! Surface Area Increased surface area means greater chances for collisions, and therefore faster rates! Catalysts Catalysts speed up reactions without being used up. They lower the activation energy, meaning more collisions to be successful!

Transition State Theory During a chemical reaction, reactants do not suddenly convert to products The formation of products is a continuous process of chemical bonds breaking and forming At some point in a successful collision, a transitional species is formed containing “partial” bonds. This species, called the transition state or activated complex, is highly energetic and normally very short-lived in the reaction. The activation energy can be defined as the energy required to form the transition state of a reaction. Typically, it decomposes immediately into the products of the reaction or back into the reactants.

Potential Energy Diagrams

Exothermic PE Diagram

Endothermic PE Diagram Why does this diagram represent an endothermic reaction? What is present at A, B and C? Describe how to determine Ea from this diagram. Describe how to determine  H from this diagram.

Temperature & Activation Energy Collision theory predicts that as temperature increases, reaction rate will increase due to the higher average kinetic energy of the particles involved in the reaction – increasing the frequency of collisions as well as the fraction of collisions with the required activation energy. Since rate increases with temperature, we expect that the rate constant will increase with temperature also.

The Reaction Mechanism A reaction mechanism is the sequence of collisions (steps) that take place during a successful reaction. A reaction mechanism that involves more than one collision will produce reaction intermediates that are then consumed in later steps of the mechanism. These intermediates are critical in providing evidence for or against a proposed mechanism.

Reaction Mechanisms, cont’d Reactions involving more than 2 colliding particles are not likely to occur in a single step. The mechanism is likely to involve several steps. The slowest step of a mechanism is called the rate-determining step. The reaction is as fast as its slowest step!

The Citric Acid Cycle Cellular respiration involves a complicated reaction mechanism – it would be next to impossible for 7 molecules to collide successfully! Most mechanisms are made up of unimolecular, bimolecular or termomlecular elementary steps.

Concentration and Rates We have seen that “concentration” of reactants is a factor in determining reaction rates.  Greater concentration of reactants means greater frequency of collisions. More frequent collisions means more successful collisions and a faster rate! Is there a mathematical relationship between reaction rates and concentration?

The Rate Law Consider the reaction: 2 NO 2 + F 2  2 NO 2 F The rate law for the reaction describes HOW the rate depends on the concentrations of the reactants: rate = k [NO 2 ] x [F 2 ] y where “k” is the rate constant, “x” and “y” are the orders of the reaction, and “x + y” is the overall order of the reaction.

Determining the Rate Law Method of Initial Rates

Let’s find the order with respect to “NO 2 ” Use two trials where only [NO 2 ] is changing - any change in rate is thus caused by NO 2 Apply the rate law to these two trials: 2.00 = 2.00 x 1 = x So the order of reaction for NO 2 is “1”. We say the reaction is first order with respect to NO 2. Analyzing the Data:

Continue by finding the order with respect to F 2. Find two trials where [F 2 ] is the only concentration that is changing = 2.00 y 1 = y So the order of reaction with respect to F 2 is also “1”. The OVERALL order of the reaction is “2” - we say the reaction is second order overall.

Rate = k[NO 2 ][F 2 ] We can now calculate the value of the rate constant, k. For example, if we rearrange the rate law for trial 1: k = 1.00 x M -1 s -1 Notice that the units for the rate constant will depend on the units for the RATE and on the overall order of the reaction.

Rate = k[NO 2 ][F 2 ] What will happen to the rate if [NO 2 ] is doubled? What will happen to the rate if [F 2 ] is tripled? What will happen to the rate if [NO 2 ] is doubled AND [F 2 ] is tripled? What was the reaction rate in Trial 4 of the experiment?

Reaction Mechanisms & Rate Laws The rate law actually gives insight into the MECHANISM for the reaction. For the reaction between NO 2 and F 2, we found the reaction to be first order in NO 2 and in F 2 This suggests that the rate determining step of the mechanism involves a collision between 1 NO 2 molecule and 1 F 2 molecule. 2 NO 2 + F 2  2 NO 2 F Can you propose a mechanism?

One possibility: NO 2 + F 2  NO 2 F + F(SLOW) NO 2 + F  NO 2 F(FAST) _______________________ 2 NO 2 + F 2  2 NO 2 F (OVERALL) Note that “F” is an intermediate in this mechanism.

A Two-Step Mechanism Propose a simple 2-step mechanism for the reaction: Br NO  2 BrNO Step 1:Br 2 + NO  Br 2 NO(slow) Step 2:Br 2 NO + NO  2 BrNO(fast) ___________________ Overall: Br NO  2 BrNO In this mechanism, Br 2 NO is an intermediate. The first step would be the RDS.

Multi-step PE Diagrams The number of “hills” equals the number of steps Each step has a heat of reaction, activated complex and activation energy Which of these steps is the rate determining step?

Catalysts A catalyst is a substance which speeds up the rate of a reaction, without being used up in the reaction. In biological reactions, catalysts are called enzymes. They provide a new mechanism for the reaction collisions – one that has a lower activation energy than the original mechanism. A homogeneous catalyst is one that is in the same phase as the reactants. For example an aqueous species catalyzing a reaction in solution, or a gaseous species catalyzing a gas-phase reaction. An example of a heterogeneous catalyst is a solid platinum surface catalyzing a gas-phase reaction.

Potential Energy Diagram & the Effect of a Catalyst In this example, the catalyst has clearly lowered the activation energy of the reaction We also see a change from a single-step to a two- step mechanism In the new mechanism, which step is the RDS?

PE Diagram with Catalyst A catalyst provides a new mechanism with a lower Ea In this example, the new mechanism has 3 steps  H is not affected by the catalyst