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Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
Chemical kinetics    ok1294986988
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Chemical kinetics ok1294986988

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  1. <ul><li>Chemical kinetics </li></ul>
  2. kinetics <ul><li>6.1 Rates of Reaction (2h) </li></ul><ul><li>Rate of reaction can be defined as “the decrease in the concentration of reactants per unit time or the increase in the concentration of product per unit time”. </li></ul><ul><li>Consider the reaction </li></ul><ul><li>2H 2 (g) + 2NO(g) -> 2H 2 O(g) + N 2 (g) </li></ul><ul><li>The rate of disappearance of H 2 is the same as the rate </li></ul><ul><li>of disappearance of NO. It is also the same as the rate </li></ul><ul><li>of appearance of H 2 O and HALF the rate of </li></ul><ul><li>appearance of N 2 from the stoichiometry of the </li></ul><ul><li>reaction in the balanced equation. </li></ul>
  3. INTRODUCTION TO REACTION RATES: <ul><li>Work in your table groups to write down an example of </li></ul><ul><ul><li>A very, very fast chemical reaction </li></ul></ul><ul><ul><li>A chemical reaction that takes about 5 – 10 minutes to occur. </li></ul></ul><ul><ul><li>A chemical reaction that takes several years to happen </li></ul></ul><ul><ul><li>A chemical reaction that takes thousands or millions of years to occur. Cooking/digesting food are both examples of medium rate chemical </li></ul></ul><ul><ul><li>An example of a chemical reaction, perhaps related to food, that you would want to </li></ul></ul><ul><ul><ul><li>slow down </li></ul></ul></ul><ul><ul><ul><li>speed up </li></ul></ul></ul><ul><li>Can you write down a link between how fast a chemical reaction is and the time it takes for that reaction to occur? </li></ul>
  4. Rate is……………….. <ul><li>For the disappearance of a reactant: </li></ul><ul><li>Change in some property/ time unit </li></ul><ul><li>The value is given a – sign </li></ul><ul><li>For the appearance of a product: </li></ul><ul><li>The value is given a + sign </li></ul>
  5. <ul><li>At Grade 10 level you learnt that chemical reactions occur when the particles in the reacting substances collide with each other with sufficient energy to produce new product chemicals. </li></ul><ul><li>Why do gases, liquids and solutions tend to react faster than solids at the same temperature? </li></ul><ul><li>Controlling reaction rates is an economic necessity. We need to produce chemicals as cheaply as possible, which usually means as quickly as possible. </li></ul><ul><li>Working in your table groups, try to list at least 5 different factors which affect how fast a chemical reaction occurs. </li></ul>
  6. Collision Theory <ul><li>As the reactants heat up, the particles move faster and so collide more often and with more energy. When they react, they must have enough energy to overcome the activation energy needed by the reaction. A rough rule of thumb which applies to many reactions (think about cooking!) is that a temperature rise of 10K approximately doubles the rate of the reaction. Mathematically, this may take some explaining. </li></ul><ul><li>Reaction Rates and Industrial Processes: </li></ul><ul><li>Work in your table group and choose one of the following industrial processes. Produce an OHT / PowerPoint to help you tell the rest of the class the key facts about your chosen process. </li></ul><ul><li>1. Haber Process </li></ul><ul><li>2. Contact Process </li></ul><ul><li>3. Manufacture of Margarine </li></ul><ul><li>4. Cat cracking </li></ul><ul><li>5. Biological washing powder </li></ul><ul><li>6. Enzyme technology </li></ul>
  7. Collision Theory states that for a reaction to take place particles must collide with sufficient energy to overcome the Activation energy for that particular reaction
  8. METHODS USED IN KINETICS TO FOLLOW A REACTION <ul><li>Method 1: Loss in mass as the reaction progresses. e.g in a reaction which produces CO 2 </li></ul><ul><li>Method 2: Volume of gas evolved as a reaction progresses. e.g in a gas syringe or by displacement of water in a graduated cylinder. </li></ul><ul><li>Method 3: Using a pH probe if there is a change in acidity as the reaction progresses. </li></ul>
  9. Method 2 Using a gas syringe
  10. More methods for following the course of a reaction <ul><li>Method 4: Use a conductivity stick if there is a change in the number of ions in solution. </li></ul><ul><li>Method 5: Use a colorimeter if there is a change in colour of one of the species. </li></ul><ul><li>There are several more methods, but I am sure you now understand that we can use any suitable method which measures changes as the reaction progresses. </li></ul>
  11. What does a real reaction look like ? What is happening to the rate of the reaction with time? How can we find the rate at any instant When is the reaction finished?
  12. What is going on during the reaction? <ul><li>At the start: The reaction is fast because here we have the highest concentration of reactants, therefore the greatest number of successful collisions. </li></ul><ul><li>As reaction progresses: the rate declines as there are fewer particles to collide. </li></ul><ul><li>At the end : the slope is zero (flat) no more particles to react </li></ul>
  13. When is the reaction HALF completed? <ul><li>The rate at any instant is the gradient (slope) at any instant. </li></ul><ul><li>Clearly, if the rate changes all the time, the time for ½ the reaction is not ½ the time it takes for the complete reaction. </li></ul>Let’s say 48cm 3 of CO 2 was given off on completion. Half reaction when 24cm 3 , Read off graph time to give 24cm 3
  14. Factors affecting rate <ul><li>Concentration in terms of Collision Theory </li></ul><ul><li>Particle size of solids </li></ul><ul><li>clearly if the same of solid is used but the particles are smaller (bigger surface area) will provide more places for reaction </li></ul><ul><li>Temperature: </li></ul><ul><li>clearly particles will have a higher collision energy at higher temp. and more successful collisions. </li></ul>
  15. Fact 1: An increase in the concentration of a solution, or the pressure of a gas, results in an increase in the reaction rate Explanation : If the concentration of a solution, or the pressure of a gas, is increased, then there are more particles in a given volume. Therefore there will be more collisions in a set amount of time and the probability of more successful collisions becomes higher. As a result the reaction rate will increase.
  16. Fact 2: An increase in surface area of a solid results in a increase in the rate of a reaction. Explanation : If the surface area of a solid is increased, there are more particles exposed to the other reactant. Therefore there will be more successful collisions in a set amount of time and the rate increases
  17. Temperature effect and collisions <ul><li>The increased Kinetic Energy of collisions at higher temperatures only accounts for a small proportion of the increased rate. </li></ul><ul><li>On average a 10 o C rise doubles the rate of a reaction. </li></ul><ul><li>The major reason for this is the proportion of particles, at a higher temperature, which now have an energy greater than the Activation energy. </li></ul>
  18.  
  19. Fact 3: An increase in temperature results in an increase in reaction rate. Explanation : If the temperature is increased, the average kinetic energy of the particles increases. There is therefore a greater chance of more collisions having an energy greater than the activation energy and reaction rate increases. Also (less importantly) because particles are moving faster, there will be more collisions in a given time, therby increasing the rate
  20. What the syllabus says……. <ul><li>6.2.5 State and explain qualitatively the Maxwell-Boltzmann energy distribution curve for a fixed amount of gas at different temperatures and its consequences for changes in reaction rate. </li></ul><ul><li>Students should be able to explain why the area under the curve is constant and does not change with temperature. </li></ul>
  21. Catalyts (Enzymes) <ul><li>Grade 10: A catalyst is a substance that increases the rate of a reaction but is not used up by the reaction. </li></ul><ul><li>SL and HL -level: Catalysts affect the rate of a reaction by providing an alternative reaction pathway with a lower activation energy </li></ul>
  22.  
  23.  
  24. Fact 4: Addition of a catalyst can lead to an increase in reaction rate Explanation : A catalyst acts by enabling a reaction to proceed via a route of lower activation energy. There will therefore be more collisions of sufficient energy to cause a reaction to occur, and rate will increase
  25. SUMMARY
  26. Orders, rate equations and the rate constant
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  28.  
  29. Mechanisms * Train Station Analogy <ul><li>People can: </li></ul><ul><ul><li>Exit a train at 20 per second </li></ul></ul><ul><ul><li>Ascend the escalator at 10 per second </li></ul></ul><ul><li> Pass the ticked barrier at 50 per second </li></ul><ul><li>People will only pass the ticket barrier at only 10 per second because of the limiting escalator. </li></ul><ul><li>This is the rate determining step which is indicated by the experimentally determined rate equation </li></ul>
  30. Equation Example <ul><li>The reaction </li></ul><ul><li>NO 2 ( g ) + CO ( g ) -> NO ( g ) + CO 2 ( g ) </li></ul><ul><li>occurs in two steps: </li></ul><ul><li>NO 2 + NO 2 -> NO + NO 3 (slow step) </li></ul><ul><li>NO 3 + CO -> NO 2 + CO 2 (fast step) </li></ul><ul><li>The slow step is the rate determining step and the rate equation should be </li></ul><ul><li>Rate = k[NO 2 ] 2 </li></ul>
  31. Orders and Reaction mechanism
  32.  
  33. Rate = k [ A ] Rate = k [(CH 3 ) 3 CCl
  34.  
  35. Going further; Mr Arrhenius
  36.  

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