1. Journal Entry 1.What is rate? 2.Do all reactions occur at the same rate? 3.Give examples of reactions that have different rates? 4.Give examples of reactions that occur at different rates under different conditions 5.Give examples of processes that cannot be controlled 6.Give examples of processes that can be controlled

2. Practice Redox Problems Ag (s) + HNO 3(aq)  AgNO 3(aq) + NO (g) + H 2 O (l) C 3 H 8 O (aq) + CrO 3(g) +H 2 SO 4(aq)  Cr 2 (SO 4 ) 3(aq) + C 3 H 6 O (aq) + H 2 O (l) I - (aq) + HSO 4 - (aq)  I 2(s) + SO 2(g) (acidic) CrO 4 2- (aq) + S 2- (aq)  S (s) + CrO 2- (g) Sb (s) + HNO 3(aq)  Sb 2 O 5(s) + NO (g) + H 2 O (l) KOH (aq) + Cl 2(g)  KCl (aq) + KClO (aq) + H 2 O (l) Zn (s) + NO 3 - (aq)  Zn 2+ (aq) + NO (g) (acidic) MnO 4 - (aq) + SO 3 2- (aq)  MnO 2(s) + SO 4 2- (aq) (basic)

3. Unit 3: Kinetics Lesson1: Reaction Rate

4. Different kinds of rates Rate of reaction Rate of reading Rate of population growth

5. Chemical Kinetics The study of whether or not a reaction will occur How fast a reactant disappears or how fast a product appears Fireworksvs. Digestion

6. Average reaction rate = -∆Reactant= ∆Product Time the symbol ∆ means ``the change in`` Therefore, = [Reactant] final - [Reactant] initial Time final -Time initial = [Product] final - [product] initial Time final -Time initial

7. 2NH 3(g)  3H 2(g) + N 2(g) Hydrogen is formed 3 times faster than nitrogen ∆[H 2 ] = 3 x ∆[N 2 ] OR1 x ∆[H 2 ] = ∆[N 2 ] ∆t ∆t3 ∆t ∆t

8. One mole of nitrogen forms, two moles of ammonia decompose -∆[NH 3 ] = 2 x ∆[N 2 ]OR-1 x ∆[NH 3 ] = ∆[N 2 ] ∆t ∆t 2 ∆t ∆t 2NH 3(g)  3H 2(g) + N 2(g)

9. Overall reaction rate relationship: -1 x ∆[NH 3 ] = ∆[N 2 ] = 1 x ∆[H 2 ] 2 ∆t ∆t 3 ∆t rate[H2] = 6.0 x 10 -2 mol/L∙s Rate [N2] = 2.0 x 10 -2 mol/L∙s Rate[NH3] = -4.0 x 10 -2 mol/L∙s = 4.0 x 10 -2 mol/L∙s 2NH 3(g)  3H 2(g) + N 2(g)

10. Average rate of any reaction aA + bB  cC + dD Can be determined by using the inverse of each coefficient in the chemical equation 1 x ∆[C] = 1 x ∆[D] = -1 x ∆[A] = -1 x ∆[B] c ∆t d ∆t a ∆t b ∆t

11. [C 4 H 9 Cl] (M) vs Time (sec)

12. Instantaneous Rate Find the slope of the line. Inst. Rate = ∆[C 4 H 9 Cl] = 0.025mol/L – 0.065 mol/L ∆t600s – 200s = 1 x 10 -4 mol/L ***BEDMAS***

13. Determining Reaction Rate monitoring of mass, pH, and conductivity Mg (s) + 2HCl (aq)  MgCl 2(aq) + H 2(g) -Mass will decrease (H 2 escaping) -pH will increase (getting more basic as HCl is used up) -Use change in conductivity to determine reaction rate of reactant ions forming product ions

14. Determining Reaction Rate Monitoring of pressure – When reactions involve gases, the pressure of the system changes 2N 2 O 5  4NO 2 + O 2 -pressure increases as two moles of N 2 O 5 decompose into 5 moles of gaseous products

15. Monitoring colour and volume – Absorption of light is directly proportional to concentration (mol/L) – In gases, they can be collected in an inverted tube & measured by displacement Determining Reaction Rate

16. Monitoring of temperature – Using a temperature probe – Increase for exothermic (heat given off or exiting!) – Decrease for endothermic (heat absorbed) Determining Reaction Rate

17. Why reaction rates? Ex. Knowing reaction rate can help doctors control insulin production in the body Ex. Industrial chemists might want to speed up ammonia production for fertilizers

18. Reaction RateReaction Time Describes the change over time that a reaction proceeds at. Merely the amount of time that a reaction takes to occur.

19. Example: @ 508°C 2HI (g)  H 2(g) + I 2(g) Time (s)[HI] (mol/L 00.1000 500.0716 1000.0558 1500.0457 2000.0387 2500.0336 3000.0296 3500.0265 GRAPH [HI] VS. TIME

21. Calculate the average rate from 0-350 seconds Avg rate = C f – C i T f -T i

22. Calculate the instantaneous rate at 0 seconds (draw a tangent line) Inst. rate = C f – C i T f -T i

23. Calculate the instantaneous rate at 100 seconds (draw a tangent line) Inst. rate = C f – C i T f -T i