Your SlideShare is downloading. ×
Chemical kinetics
Upcoming SlideShare
Loading in...5
×

Thanks for flagging this SlideShare!

Oops! An error has occurred.

×

Introducing the official SlideShare app

Stunning, full-screen experience for iPhone and Android

Text the download link to your phone

Standard text messaging rates apply

Chemical kinetics

168
views

Published on

GCE A LEVEL TOPIC: (A2) CHEMICAL KINETICS …

GCE A LEVEL TOPIC: (A2) CHEMICAL KINETICS

PLEASE DOWNLOAD BECAUSE THERE ARE MANY ANIMATIONS THAT HIDE SOME OF THE CONTENTS (THE ANIMATIONS DO NOT PLAY DURING THE PREVIEW)

Published in: Science, Technology, Business

0 Comments
1 Like
Statistics
Notes
  • Be the first to comment

No Downloads
Views
Total Views
168
On Slideshare
0
From Embeds
0
Number of Embeds
0
Actions
Shares
0
Downloads
10
Comments
0
Likes
1
Embeds 0
No embeds

Report content
Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel
No notes for slide

Transcript

  • 1. + Reaction Kinetics (A2) Prepared by: Faiz Abdullah
  • 2. + AS recap In As level, you should know what is:  Rate of reaction  Collision theory  Boltzmann distribution of energy  Catalyst  Activation Energy  Effect of temperature, concentration and catalyst on rate of reaction
  • 3. + For A2, we are dealing with DATA MANIPULATION
  • 4. + Order of Reaction What do you mean by order?
  • 5. + Measuring rate of reactions Usually, we are looking at HOW FAST THE CONCENTRATION OF REACTANTS IS FALLING AT ONE TIME. Suppose we have: A(aq)+B(g)  Products We can measure rate of reaction with: -The decreasing concentration of A in mol/dm3 in 1 min. -The decreasing volume of gas B in 1min
  • 6. + Orders of reaction Suppose of you have: A + B  products From the experiment: you found out that…. When [A] doubles, rate doubles. Rate of reaction is proportional to [A] Therefore, order with respect to A is 1 When [A] doubles, rate increases four times Rate of reaction is proportional to [A]2 Therefore, order of reaction w.r.t. A is 2 When [A] doubles, rate of reaction does not change Rate of reaction does not depend on [A] Therefore, order w.r.t A is 0 WARNING: YOU CANNOT DEDUCE THE ORDER OF REACTION JUST BY LOOKING AT THE EQUATION!!!! ORDERS OF REACTIONS ARE ALWAYS FOUND BY DOING EXPERIMENTS
  • 7. + Rate equation Suppose A + 2B + C  products From experiments, we found out that: [A] doubles, rate of reaction doubles Therefore, order w.r.t A is 1 [B] doubles, rate of reaction increases by 4 Therefore, order w.r.t B is 2 [C] doubles, rate of reaction does not change Therefore, order w.r.t C is 0 Rate Equation: Rate = k [A] [B]2 Where k is rate constant
  • 8. + Why is C in the chemical equation but not in the rate equation???? Some reactions occur in multiple steps: Step 1: A + 2B  2C + D Step 2: C + D  products Final : A + 2B + C  products SOME STEPS ARE SLOW AND SOME ARE FASTER SLOW FAST THIS STEP IS THE RATE-DETERMINING STEP OVERALL REACTION RATE DEPENDS ON THE SLOW STEP WHEN you measure rate of reaction, what you are actually Measuring is the rate of the determining step!!!
  • 9. + Rate constant, k Rate = k [A] [B] Rate constant is constant (does not change value) only when concentrations of reactants are changing. RATE CONSTANT CHANGES WHEN: 1. TEMPERATURE CHANGES 2. ADDING CATALYST
  • 10. + DEDUCING ORDER BY INTIAL RATES METHOD Run Initial [A]/mol Initial [B]/mol Initial rate/mols-1 1 1.00 1.00 1.25 x 10-2 2 1.00 2.00 2.5 x 10-2 3 2.00 2.00 2.5 x 10-2 Can you find the rate equation? Can you find k? Make [A] constant, [B] x 2, rate x 2 Order w.r.t B is 1 Make [B] constant, [A] x 2, rate same Order w.r.t A is 0 Rate = k [B] k=1.25 x 10-2 s-1
  • 11. + Deducing order from graphs First- order reaction Zero-order reaction
  • 12. + Rate concentration graph
  • 13. + Graphs Summary
  • 14. + Half-life of first-order reaction Half life  time taken to get half of the final concentration
  • 15. + Half-life equation Half-life (in seconds) can be used to find k, rate constant: T1/2 = ln (2) / k
  • 16. + EXPERIMENTAL TECHNIQUES FOR STUDYING RATES ① Sampling followed by titration ② Using a colorimeter ③ Measurement of gas evolved
  • 17. + CATALYST ① Homogenous: catalyst same phase as the reactants ② Heterogeneous: catalyst and reactants different phases
  • 18. + CATALYSIS In AS, you need to know FOUR SPECIFIC EXAMPLES
  • 19. + 1) HABER PROCESS N2 + 3H2  2NH3 Catalyst: Iron Note: Transition metals are good at acting as catalysts because their atoms have unfilled d-orbitals. Gases are adsorbed on to the surface of the metal, forming weak bonds. ① Formation of bonds with the metal surface weakens the bonds within the gas molecules ② The orientation of the adsorbed molecules may be favorable for the reaction THIS IS A HETEROGENEOUS SYSTEM
  • 20. + 2) Catalytic converters in vehicle exhausts
  • 21. +  Catalytic converters in vehicle exhausts aim to remove a number of pollutant gases from vehicle exhausts.  Pollutants: Nitrogen oxides, carbon monoxides.  Inside the ceramic honeycomb, it has a very thin coat of: o Platinum and palladium: oxidize CO and unburnt hydrocarbons o Platinum and rhodium: reduce NOx to N2
  • 22. + 3) Nitrogen oxides in the atmosphere  Studies on acid rain have concluded that in the atmosphere the presence of oxides of nitrogen, particularly NO2, increases rate of oxidation of SO2  SO3.  NO2 remains unchanged and is thought to form a weak intermediate with SO2.
  • 23. + 4) The role of Fe2+ in I-/S2O8 2- reaction Step 1: S2O8 2- + 2Fe2+  2 SO4 2- + 2Fe3+ Step 2: 2Fe3+ + 2I-  2Fe2+ + I2 Overall reaction: S2O8 2- + 2I-  2Fe2+ + I2 Fe2+ does not change overall Although there are two steps, Ea is lowered overall