KINETICS Chapter 16

KINETICS Define Kinetics. Define reaction rate. Discuss the collision theory. List the factors (nature of reactants, concentration, temperature, surface area, and catalysis) that effect reaction rates and be able to explain them according to collision theory. Be able to identify the parts of a potential energy diagram including activation energy, Ea, transition state, energy of reactants, energy of products. Be able to identify a potential energy diagram for an exothermic reaction and an endothermic reaction. Describe reaction mechanisms for simple reactions and be able to identify the parts of the multiple steps as well as determine the final balanced equation. Perform simple calculations. Define rate law

KINETICS concerned with the speed at which reactions occur. make a distinction between rate and time – AS RATE INCREASES, TIME DECREASES. The general rate equation for any chemical reaction: change in concentration   average rxn rate = [reactant or product] time change in time

COLLISION THEORY This theory states that in order for molecules to react, they must collide. The collision can involve a particle and a container wall or two particles. Rarely does it involve more than two particles.

COLLISION THEORY There are three parts: The particles must collide. The particles must collide with the correct orientation. The particles must collide with sufficient energy to form an activated complex.

COLLISION THEORY Collision theory helps explain why chemical reactions proceed through a series of steps called reaction mechanisms. Most chemical equations show three or more particles on the reactant side There is an intermediate step first where one H2 molecule collides with an O2, and then the other H2 does. If there are no collisions, there would be no reactions.

FACTORS AFFECTING RXN RATE DEMO: solid vs. liquids A. Nature of the Reactants The rate of a reaction depends on the particular reactants and the complexity of the bonds that have to be broken and formed in order for the reaction to proceed. Slight rearrangement of atoms is usually rapid at room temperature. Ionic reactions occur very fast. Why is that? Reactions involving covalent bonds take place more slowly at room temperature. The phase of the reactants can have an effect on reaction rate. Homogeneous reactions occur more rapidly than heterogeneous reactions. Reaction rates should be highest among gases.

FACTORS AFFECTING RXN RATE DEMO: 30% vs. 3% B. Concentration Remember that the rate of reaction is defined as a change in the concentration over time. WHY: Increase concentration and there are more particles in a given volume. So, the chances of the particles colliding increase.

FACTORS AFFECTING RXN RATE DEMO: clump vs. powder C. Surface Area The larger the surface area of the reactant, the greater the number of particles that are exposed for the reaction. WHY: Increasing the surface area increases the frequency of collisions.

FACTORS AFFECTING RXN RATE DEMO: cold vs. room vs. hot D. Temperature Increased temperature usually increases the reaction rate. The temperature rises and the average kinetic energy increases. The faster the molecules move, the more frequently they will collide. Increasing the temperature gives more particles the required activation energy and also increases the effectiveness of the collision.

TEMPERATURE As the temperature increases, there are more molecules with higher energies. The number of molecules which can react – those with kinetic energies greater than the activation energy – are represented by the BLUE area to the right of the black line.

FACTORS AFFECTING RXN RATE DEMO: yellow  red  yellow E. Catalysts A catalyst is a substance that increases the rate of a reaction without itself being used up in the reaction. Catalysts change the reaction mechanism for a reaction. A catalyst does not appear in the overall chemical equation. Catalysts work by lowering the amount of energy required to get the reactants to react. Opposite from catalysts are inhibitors. Inhibitors slow down the rate of a reaction by binding to the reactants to keep them from reacting.

ENERGY IN CHEMICAL RXNS Energy is required to break the bonds that hold the reactants together. The energy must be present in the reacting particles before a collision. Most collisions are elastic, but some collisions have enough energy to cause changes in the electron cloud of the colliding molecules.

ENERGY IN CHEMICAL RXNS Particles of matter have both potential and kinetic energy. Potential energy is stored energy. Kinetic energy is the energy of motion.

ENERGY IN CHEMICAL RXNS ACTIVATION ENERGY, Ea : Particles must posses a certain minimum amount of energy in order to react. Energy is required to force the reactants up and over an energy barrier . The energy required to form the activated complex is known as the activation energy.

ENERGY IN CHEMICAL RXNS TRANSITION STATE: The brief interval of bond disruption and bond formation once the reactants have enough activation energy and a place where the activated complex is formed. ACTIVATED COMPLEX : a short-lived structure formed by the colliding molecules. The activated complex exists along the reaction pathway at a point where the energy is the greatest - the transition state.

POTENTIAL ENERGY DIAGRAM

ENERGY IN CHEMICAL RXNS HEAT OF REACTION (ΔH or ΔE) is the energy difference between the reactants and products.

ENERGY IN CHEMICAL RXNS ENDOTHERMIC RXN: the energy of the products is more than the energy of the reactants

ENERGY IN CHEMICAL RXNS EXOTHERMIC RXN: the energy of the products is less than the energy of the reactants.

ENERGY IN CHEMICAL RXNS A CATALYST lowers the activation energy needed for the reaction by carrying the reaction along a different path. Result of catalyst

HANDOUT

FIGURE 1

FIGURE 2

FIGURE 3

SOLVING FOR THE RXN RATE CO + NO2  CO2 + NO Rxn rate = [NO] = [NO] at end time - [NO] at start time t end time – start time [NO] = 0.010M , t = 2s [NO] = 0.000M , t = 0s [NO] = 0.010M - 0.000M = 0.010M = 0.0050M t 2s – 0s 2s s

SOLVING FOR THE RXN RATE Positive value = concentration is increasing. Negative value = concentration is decreasing. BUT, all reaction rates are positive, so you put a negative sign in front of the equation for rate of consumption (as opposed to rate of production).

PRACTICE A  2B The concentration of A, [A], goes from 2.20M (time = 0.00 min) to 1.10M (time = 3.00 min). What is the rate of A?

THE RATE LAW The rate for the following equation can also be written using a constant: A  B Rate = K[A]   The rate law shows a mathematical relationship between reaction rate and concentration of reactants. It is a numerical value that relates reaction rate and the concentration of reactants at a given temperature.

THE RATE LAW “K” is the specific rate constant. “K” is unique for every reaction and every temperature. The reaction rate is directly proportional to the molar concentration of A. If you double the concentration of A, then the rate doubles. We will discuss “K” more in the next chapter on chemical equilibrium. A large “K” value means that A reacts rapidly to form B. What do you think a small value means?

REACTION MECHANISMS Most chemical reactions consist of a series of steps. A series of reactions that lead the reactants to become products is called a REACTION MECHANISM. Each step normally involves the collision of only two particles. A detailed reaction mechanism describes the order in which bonds break and atoms rearrange throughout the course of the chemical reactions.

2NO and 2F react to form 2NOF REACTION MECHANISMS 2NO (g) + F2 (g)  2 NOF (g) 2NO and 2F react to form 2NOF OCCURS IN TWO STEPS   Step #1: NO + F2  NOF2 Step #2: NOF2 + NO  2NOF NO + F2 + NOF2 + NO  NOF2 + 2NOF

NOF2 = intermediate product   INTERMEDIATE PRODUCT: A substance that is produced in one step of a reaction and then is consumed in a later step of that reaction. It appears first as a product and later as a reactant disappears. Multistep reaction mechanisms always involve one or more intermediate products.

REACTION MECHAMISMS CATALYST: A substance that disappears in one step of a reaction and then appears in a later step of that reaction. It appears first as a reactant and later as a product. When doing reaction mechanisms, catalysts and intermediates do not appear in the net chemical reaction.

REACTION MECHANISMS **Do not confuse the transition state with the intermediate products formed in the multisteps of a chemical equation. Each intermediate product has its own transition state and activated complex.

REACTION MECHANISMS The rate of the overall reaction is limited to the rate of the slowest step and is called the rate determining step.

Another Example N2O decomposes to N2 and O2 gases. The intermediate product O(g) is believed to be a part of the reaction mechanism. Below is a proposed 2 step mechanism consistent with this intermediate product.   1. Write the balanced equation. 2. Check off those substances that have been used. 3. Fill in the missing reactants and/or products.

Another Example Balanced eqn: Step #1: N2O  ______ + N2   Step #1: N2O  ______ + N2 Step #2: ______ + O  ______ + ______ Now, rewrite these two steps into one chemical equation, showing where the intermediate product is cancelled out on both sides of the arrow.

PRACTICE I. Ozone, O3, decomposes into diatomic oxygen, O2. The reaction mechanism involves the intermediate product singular oxygen, O. The chemical equation is: O3 + O3  O2 + O2 + O2 See if you can fill in the missing molecules in the two steps below. Remember that each individual equation must be balanced: Step 1: O3  __________ + O Step 2: O + __________  O2 + O2

PRACTICE Okay, now it is your turn to try it on your own. Remember to balance the equation. H2O2  O2 + H2O The decomposition of hydrogen peroxide involves a three-step mechanism. Iodine, I-, and H3O+ are catalysts. The H3O2+ and HOI ions are intermediate products. I’ve written parts of the steps below. Fill in the missing molecules and then write the balanced equation. (Show the intermediate steps cancelled out just like in your notes.)

PRACTICE H2O2 + H2O2  O2 + H2O + H2O Iodine, I-, and H3O+ are catalysts. The H3O2+ and HOI ions are intermediate products. Step 1: H2O2 + H3O+  H3O2+ + _________ Step 2: ___________ + I-  _______ + HOI Step 3: HOI + _______  _______ + I- + ______ Balanced chemical equation: