1. Chapter 13: Osmotic Pressure Chapter 14: Kinetics But first- NOYCE

2. Osmosis is the movement of water through a semi-permeable membrane

3. Osmotic pressure is used to determine the molecular weight of polymers  =cRT  = osmotic pressure in atm c= concentration in mol/L R= 0.0821L  atm/K  mol T= temperature in Kelvin 12.5g of a cellulose derivative was dissolved in 1L of water. The osmotic pressure of the solution was 0.00210 atm at 30°C. What is the molar mass of this molecule?

4. Low levels of protein in blood cause edema due to osmotic pressure differences

5. CONTROL OF REACTIONS Thermodynamics tells us if the products are more/less stable than the reactants Kinetics tells us if the reaction goes fast enough for us to notice

6. Back to our “organizational chart” of what controls reactivity…

7. Preview: The rate of a reaction is described by a rate equation (rate law)

8. Will an increase in temperature increase a reaction rate? 1. Yes. 2. No.

9. Will an increase in concentration increase a reaction rate? 1. Yes. 2. No.

10. Will increasing the pressure increase a reaction rate? 1. Yes. 2. No.

11. Collision theory is a qualitative explanation of how reactions occur and why rates differ Reactions occur when ◦ Molecules collide… ◦ In the correct orientation… ◦ With enough energy Consider: NO + O 3  NO 2 + O 2 Molecules collide Bonds are formed and break product molecules separate

12. What can control how fast a reaction occurs?

14. So, what controls the rate of a reaction? Number of collisions How often they collide in a shape that allows new bonds to form The energy of the colliding reactant molecules We’ll consider dependence on: 1.Concentration a. Rate laws b. Concentration vs. time relationships 2.Temperature and activation energy 3.Mechanism

15. Concentration Dependence It makes sense that as concentration increases, the number of collisions per second will increase Therefore, in general, as concentration increases, rate increases But, it depends on which collisions control the rate So, you can’t predict concentration dependence: it must be measured experimentally

16. Types of measured rates: Rate over time: Instantaneous rate: Initial rate:

17. Example of rate measurement:

18. Rate Laws (also called Rate Equations) For the reaction: 2 N 2 O 5  4 NO + O 2 Rate = k[N 2 O 5 ] For the reaction: NO 2  NO + ½ O 2 Rate = k[NO 2 ] 2 For the reaction: CO + NO 2  CO 2 + NO Rate = k[CO][NO 2 ] first order reaction second order reaction first order in CO and in NO 2 ; second order overall

19. Determining a Rate Law Determining the rate law must be done by experiment; the reaction equation does not tell you the rate law Two methods: Initial Rates and the Graphical Method Method of Initial Rates Measure the rate of the reaction right at the start. Vary the starting concentrations Compare initial rates to initial concentrations

20. Determining a Rate Law: Initial Rate Method Isolation of variables: Vary only one concentration at a time and keep temperature constant If concentration doubles and: ◦ Rate does not change, then zero order ◦ Rate doubles, then first order ◦ Rate quadruples, then second order General Rule:

21. Initial Rate Method: Example 1 What is the rate law?

22. Initial Rate Method: Example 2

23. Concentration-Time Relationships

24. Example 1

25. The decomposition of nitrous oxide at 565 o C, 2 N 2 O  2 N 2 + O 2 is second order in N 2 O. If the reaction is initiated with [N 2 O] equal to 0.108 M, and drops to 0.940 M after 1250 s have elapsed, what is the rate constant? Example 2