Unit 6: Solutions and Kinetics Rate Laws

Collision Theory In order for a reaction to occur… Reactants must collide 1 2 Collision must be at the correct orientation 3 Collision must have minimum amount of energy for bonds to break Only a small number of collisions meet the requirements and result in a reaction

1. Reactants Must Collide In order for two molecules to react, they must come in contact with one another N O F There’s no way they’ll ever react if they don’t run into one another!

2. Collision with Incorrect Orientation For a collision to result in a chemical reaction, it must occur with the correct orientation N O F This is not the correct orientation. The reaction will not happen.

2. Collision with Correct Orientation For a collision to result in a chemical reaction, it must occur with the correct orientation F N O N O F But not every collision produces a reaction—they must collide in the correct orientation to react F This is the correct orientation. The reaction will happen.

3. Collision with Not Enough Energy For a collision to result in a chemical reaction, it must occur with the minimum energy for reaction N O F The collision does not have enough energy to produce a reaction

3. Collision with Enough Energy For a collision to result in a chemical reaction, it must occur with the minimum energy for reaction F N O N O F F This collision had more energy (faster moving molecules). A reaction will occur..

Factors Affecting Reaction Rate Temperature Concentration Surface Area Catalysts

For most reactions, as temperature increases, reaction rate increases How does temperature affect the reaction rate? With higher energy molecules, collisions will have higher energy and more often result in reaction If molecules are at a higher temperature, they have a higher average kinetic energy Reactants must collide with at least energy equal to the activation energy For most reactions, as temperature increases, reaction rate increases

Concentrations of Reactants How does the concentration of reactants affect the reaction rate? Only a small fraction of the collisions meet the requirements and result in a reaction If more collisions occur, more will meet the requirements and result in a reaction More reactants mean more collisions will occur For most reactions, as reactant concentration increases, reaction rate increases

Surface Area of Reactants How does the surface area of the reactants affect the reaction rate? More reactants can collide at the same time and a fraction of those will result in reaction Larger surface area means more particles can come in contact with each other at the same time Reactants must collide in order to react As surface area increases, reaction rate increases

Catalysts Catalyst: – a substance that increases the rate of reaction without being used up creates a lower energy reaction pathway A + B + C  D + C “C” is the catalyst…it is present in the beginning and in the end Enzymes are catalysts in the body

Concentration Effect on Reaction Rates Increasing the concentration of a reactant usually increases the rate of reaction. HOWEVER, increased concentration might actually have little effect on the rate of reaction. How can we tell?

Rate Order and Rate Laws aA + bB  cC + dD General Form of Rate Law: rate = k[A]x[B]y rate order with respect to A rate order with respect to B rate constant concentration of reactant A concentration of reactant B

Rate Order and Rate Laws Rate orders are found experimentally Change the concentration of one reactant at a time to see how the rates are affected Rate units: M/s (change in Molarity per second)

Rate Law Example #1 Reaction: A + B + C  D Trial [A] [B] [C] Rate (M/sec) 1 0.1 0.05 1.52 x 10-5 2 0.3 3 0.2 3.04 x 10-5 4 6.08 x 10-5 1. What happens to the rate when [A] doubles? 2. What happens to the rate when [B] triples? 3. What happens to the rate when [C] doubles?

Rate Law Example #1 Summary: Trial [A] [B] [C] Rate (M/sec) 1 0.1 0.05 1.52 x 10-5 2 0.3 3 0.2 3.04 x 10-5 4 6.08 x 10-5 Summary: Rate = k[A]1 [B]0 [C]1 , but [B]0 = 1 so it drops out Thus Rate = k[A]1 [C]1 the overall rate order is 2 (sum of A, B, C orders) Solve for k: Using data from trial 1:

Rate Law Practice #1 Reaction: A + B  C Trial [A] [B] Rate (M/sec) 1 0.600 0.360 0.460 2 1.080 1.380 3 0.200 0.051 1. What happens to the rate when [A] triples? 2. What happens to the rate when [B] triples? 3. What is the rate law for this reaction? 4. Calculate the rate constant for this reaction.

Rate Law Practice #2 Reaction: A + B + C  D Trial [A] [B] [C] Rate (M/sec) 1 0.1 0.01 2 0.2 3 0.02 4 0.08 1. What happens to the rate when [A] doubles? 2. What happens to the rate when [B] doubles? 3. What happens to the rate when [C] doubles? 4. What is the rate law for this reaction? 5. Calculate the rate constant for this reaction.

Rate Law Practice #2 Summary: Trial [A] [B] [C] Rate (M/sec) 1 0.1 0.01 2 0.2 3 0.02 4 0.08 Summary: Rate = k[A]2 [B]1 [C]0 , but [C]0 = 1 so it drops out Thus Rate = k[A]2 [B]1 The overall rate order is 3. Solve for k: Using data from trial 3:

Rate Law Practice #3 1. What is the rate law for this reaction? Reaction: NO (g) + NO2 (g) + O2 (g)  N2O5 (g) Trial [NO] [NO2] [O2] Rate (M/sec) 1 0.1 2.1 x 10-2 2 0.2 4.2 x 10-2 3 0.3 6.3 x 10-2 4 1. What is the rate law for this reaction? 2. Calculate the rate constant for this reaction.

Rate Law – Online Simulation Reaction: H2O2(aq) + 3 I-(aq) + 2 H+(aq) → I3-(aq) + 2 H2O(l) Trial H2O (mL) H2O2 (mL) HCl (mL) KI (mL) Temp. (°C) Time (s) 1 60 10 20 25 2 50 3 40 4 5 6 45 1. What happens to the rate when you double the amount of H2O2?

Rate Law – Online Simulation Reaction: H2O2(aq) + 3 I-(aq) + 2 H+(aq) → I3-(aq) + 2 H2O(l) Trial H2O (mL) H2O2 (mL) HCl (mL) KI (mL) Temp. (°C) Time (s) 1 60 10 20 25 2 50 3 40 4 5 6 45 2. What happens to the rate when you double the amount of HCl?

Rate Law – Online Simulation Reaction: H2O2(aq) + 3 I-(aq) + 2 H+(aq) → I3-(aq) + 2 H2O(l) Trial H2O (mL) H2O2 (mL) HCl (mL) KI (mL) Temp. (°C) Time (s) 1 60 10 20 25 2 50 3 40 4 5 6 45 3. What happens to the rate when you double the amount of KI?