1. Chemical Kinetics Unit 10

2. Kinetics –Kinetics – the study of the speeds of chemical reactions and the mechanisms by which reactions occur. Reaction RatesAlso referred to as “Reaction Rates” Rates of chemical change usually are expressed as the amount of reactant forming products per unit time. ∆[reactants] ∆ time Kinetics

4. How do reactions really occur?How do reactions really occur? –Reactant particles MUST collide! Rates of chemical reactions are described in a model called collision theory.Rates of chemical reactions are described in a model called collision theory. –Studies show most molecular collisions do NOT result in a reaction…why?? Atoms, ions, and molecules react to form products only when they collide with the proper orientation and sufficient energy.Atoms, ions, and molecules react to form products only when they collide with the proper orientation and sufficient energy. Collision Theory

6. Effective collisions are defined by two conditions:Effective collisions are defined by two conditions: –exactly the right orientation to react (head on collision is best) –enough energy to break bonds and form new ones The minimum amount of energy required is called the reaction’s activation energy.The minimum amount of energy required is called the reaction’s activation energy. –The activation energy is a barrier that reactants must get over to react –The higher the barrier the larger the amount of energy needed for the reaction to proceed Collision Theory

9. During a reaction, a particle that is neither reactant nor product forms momentarily, called an activated complexDuring a reaction, a particle that is neither reactant nor product forms momentarily, called an activated complex –if there is sufficient energy –and if the atoms are oriented properly An activated complex is a kind of transition molecule which has similarities to reactants & productsAn activated complex is a kind of transition molecule which has similarities to reactants & products –An activated complex is the arrangement of atoms at the peak of the activation-energy barrier. Collision Theory

12. Collision theory explains why some naturally occurring reactions are immeasurably slow at room temp.Collision theory explains why some naturally occurring reactions are immeasurably slow at room temp. –Carbon and Oxygen react when charcoal burns, but this reaction has a high activation energy –At room temp, the collisions of oxygen and carbon molecules aren’t energetic enough to react –But the reaction can be helped along a number of ways Collision Theory

13. It is possible to vary the conditions of the reaction, the rate of almost any reaction can be modifiedIt is possible to vary the conditions of the reaction, the rate of almost any reaction can be modified ocollision theory can help explain why the rates can be modified Several strategies can be used to speed up reactions:Several strategies can be used to speed up reactions: oIncrease the temperature oIncrease the concentration oDecrease the particle size oEmploy a catalyst Reaction Rates

14. Increasing the temperature speeds up the reaction, while lowering the temperature slows down the reactionIncreasing the temperature speeds up the reaction, while lowering the temperature slows down the reaction For every 10 o C increase in temperature, the reaction rate usually DOUBLES!For every 10 o C increase in temperature, the reaction rate usually DOUBLES! Recall, temperature is directly proportional to average kinetic energy:Recall, temperature is directly proportional to average kinetic energy: –Average KE = ½ x mass x velocity 2 At a higher temperature, there are more effective collisions – more molecules have the velocity / energy needed to reactAt a higher temperature, there are more effective collisions – more molecules have the velocity / energy needed to react Temperature

15. Just sitting out, charcoal does not react at a measurable rateJust sitting out, charcoal does not react at a measurable rate –However, when a starter flame touches the charcoal, the temperature is increased so atoms of reactants collide with higher energy and greater frequency –Some of the collisions are high enough in energy that the product CO 2 is formed o The energy released by the exothermic reaction then supplies enough energy to get more C and O 2 over the activation- energy barrier Remove the starter flame: the reaction will continue on its own. Remove the starter flame: the reaction will continue on its own.

16. The more reacting particles you have in a given volume, the higher the rate of reaction.The more reacting particles you have in a given volume, the higher the rate of reaction. Cramming more particles into a fixed volume increases the concentration of reactants:Cramming more particles into a fixed volume increases the concentration of reactants: Increasing the concentration, increases the frequency of the collisions, and therefore increasing the reaction rate.Increasing the concentration, increases the frequency of the collisions, and therefore increasing the reaction rate. Concentration

17. The smaller the particle size, the larger the surface area for a given mass of particles. Effectively is increasing the concentration.The smaller the particle size, the larger the surface area for a given mass of particles. Effectively is increasing the concentration. An increase in surface area increases the amount of the reactant exposed for collision to take place…An increase in surface area increases the amount of the reactant exposed for collision to take place… –Which increases the collision frequency and the reaction rate. Methods:Methods: –Grinding the reactants into a powder –Dissolving in a solvent. Particle Size

18. A catalyst is often the best way to speed up an reaction. In fact, some reactions simply will not go forward measurably without one. A catalyst is a substance that increases the rate of a reaction without being changed or used up during the reaction The key is that they permit reactions to proceed at lower activation energy than is normally required Catalyst

19. Catalysts lower the required activation energy by providing an alternative path or arrangement for the molecules to react. Not always understood why certain catalyst behave the way the do…they just do! By lowering the Ea threshold, more molecules in the system will have the required energy to surmount the barrier. Catalyst

22. The type of reactant substances also plays a role in reaction rates. Weak bonds = easier to break, reactants with weak bonds will react faster. –Essentially results in a low activation energy. Strong bonds = harder to break, reactants with strong bonds will react slower (high Ea) Nature of Reactants

23. Reaction Rate Laws relate the speed of a reaction to the concentrations of the reactants (or sometimes products).Reaction Rate Laws relate the speed of a reaction to the concentrations of the reactants (or sometimes products). Determined from experimentDetermined from experiment For equation: A  BFor equation: A  B –Rate = ∆[A] / ∆ t Where ∆[A] = change in molarity of reactant AWhere ∆[A] = change in molarity of reactant A ∆ t is the time over which the reaction occurred∆ t is the time over which the reaction occurred –The rate of disappearance of A is proportional to its molar concentration. –Rate = k[A] Where k is the specific rate constantWhere k is the specific rate constant Reaction Rate Laws

24. k is unique for every reactantk is unique for every reactant –Often expressed in s -1 or L / (mol * s) –If the reaction is fast, k is large. Reaction Rate OrdersReaction Rate Orders –The “order” of a reaction is the power to which the concentration must be raised to provide for the actual rate observed. –First order reactions have rates directly proportional to one reactant. If reactant concentration doubles, rate doubles. –Rate = k[A] 1 (note: powers of 1 usually not shown) Reaction Rate Laws

25. First Order Reaction Rates

26. Second Order Rate LawsSecond Order Rate Laws –Rate = k[A] 2 –Ex.: If [A] is doubled, the rate quadruples. The General Rate LawThe General Rate Law –The order of reaction for each reactant is the value of its exponent. Overall order of the reaction is the sum of exponents. –For reaction: xA + yB  Products –Rate = k[A] m [B] n –Overall order = m + n Reaction Rate Laws

27. Second Order Reaction Rates

28. “Method of initial rates” determines order by comparing rates with varying concentrations of reactants.“Method of initial rates” determines order by comparing rates with varying concentrations of reactants. When [A] doubles, initial rate doubles (= first order).When [A] doubles, initial rate doubles (= first order). When [B] doubles, rate quadruples(= second order).When [B] doubles, rate quadruples(= second order). Rate Law: Rate = k[A] 1 [B] 2 Overall order of reaction: 1+2 = 3Overall order of reaction: 1+2 = 3 Determining Reaction Order TrialInitial [A] (M) Initial [B] (M) Initial Rate (mol/L*s) 10.100 2.00 x 10 -3 20.2000.1004.00 x 10 -3 30.200 16.00 x 10 -3

29. Run #Initial [A] ([A] 0 ) Initial [B] ([B] 0 ) Initial Rate (v 0 ) 11.00 M 1.25 x 10 -2 M/s 21.00 M2.00 M2.5 x 10 -2 M/s 32.00 M 2.5 x 10 -2 M/s What is the order with respect to A? What is the order with respect to B? What is the overall order of the reaction? What is the rate law? 0 1 1 Rate = k[B]

30. Reaction rates decline over time as the concentration of reactants decrease. The rate at any given point in time is called the “Instantaneous Reaction Rate” Equal to the slope of the tangent to the curve at any given point in time. Rate = slope of tangent = (y 2 – y 1 ) / (x 2 – x 1 ) Instantaneous Reaction Rates

31. Br 2 (aq) + HCOOH (aq) 2Br - (aq) + 2H + (aq) + CO 2 (g) average rate = -  [Br 2 ] tt = - [Br 2 ] final – [Br 2 ] initial t final - t initial slope of tangent slope of tangent slope of tangent instantaneous rate = rate for specific instance in time 13.1

32. If you know the rate law and reactant concentrations, you can directly calculate the instantaneous rate. Example:Example:  2N 2 O 5  4NO 2 + O 2  Rate = k[N 2 O 5 ] where k = 1.0 x 10 -5 s -1 and concentration of N 2 O 5 is 0.350 M.  Rate = 3.5 x 10 -6 mol/(L*s) Instantaneous Rates