1. LOGO Spring 2015 Lecture 9: Chemical Kinetics Course lecturer : Jasmin Šutković 20th May 2015

2. Contents International University of Sarajevo 1. Factors that affect the chemical kinetics 2.Reaction rates and rate laws 3.Methods of Determining Reaction Orders 4.Half-Lives and Radioactive Decay Kinetics 5.The Collision Model of Chemical Kinetics 6.Catalysis

3. 1. Factors That Affect Reaction Rates  Chemical kinetics – Study of reaction rates, or the changes in the concentrations of reactants and products with time – By studying kinetics, insights are gained into how to control reaction conditions to achieve a desired outcome  Chemical kinetics of a reaction depend on various factors 1. Reactant concentrations 2. Temperature 3. Physical states and surface areas of reactants 4. Solvent and catalyst properties

4. Concentration Effects  Two substances cannot react with each other unless their constituent particles come into contact; if there is no contact, the rate of reaction will be zero.  The more reactant particles that collide per unit time, the more often a reaction between them can occur.  The rate of reaction usually increases as the concentration of the reactants increases.

5. Example

6. Temperature Effects  Increasing the temperature of a system increases the average kinetic energy of its constituent particles.  As the average kinetic energy increases, the particles move faster, so they collide more frequently per unit time and possess greater energy when they collide, causing increases in the rate of the reaction.  Rate of all reactions increases with increasing temperature and decreases with decreasing temperature.

7. Homogenous VS Heterogeneous solutions  Homogeneous Mixtures are composed of two or more components that are equally (uniformly) distributed throughout the system, examples: Salt dissolved in water, Soapy water, Hydrochloric acid & water, Alcohol & water, Sugar dissolved in water,  Heterogeneous Mixtures are composed of two or more components that are unequally (not uniformly) distributed though out the system, examples: Sand & water (liquid & solid), Oil & water (immiscible liquids), Table salt crystals & sugar crystals (solids & solids), REMINDER !!

8. Phase and Surface Area Effects  If reactants are uniformly dispersed in a single homogeneous solution, the number of collisions per unit time depends on concentration and temperature.  If the reaction is heterogeneous, the reactants are in two different phases, and collisions between the reactants can occur only at interfaces between phases; therefore, the number of collisions between the reactants per unit time is reduced, as is the reaction rate.

9. Solvent Effects  The nature of the solvent can affect the reaction rates of solute particles.  Solvent viscosity is also important in determining reaction rates. 1. In highly viscous solvents, dissolved particles diffuse much more slowly than in less viscous solvents and collide less frequently per unit time. 2. Rates of most reactions decrease rapidly with increasing solvent viscosity.

10. Viscosity ….REMINDER!!! The viscosity of a fluid is a measure of its resistance on gradual deformation by shear stress or tensile stress.

11. Catalyst Effects  Catalyst is a substance that participates in a chemical reaction and increases the rate of the reaction without undergoing a net chemical change itself.  Catalysts are highly selective and often determine the product of a reaction by accelerating only one of several possible reactions that could occur.

12. 2. Reaction Rates The rate of a reaction is the speed at which a reaction happens. If a reaction has a low rate, that means the molecules combine at a slower speed than a reaction with a high rate. Some reactions take hundreds, maybe even thousands, of years while others can happen in less than one second.

13. Reaction Rates cont.… – Expressed as the concentration of reactant consumed or the concentration of product formed per unit time A + 2 B → 3 C – Units are moles per liter per unit time (M/s, M/min or M/h) – To measure reaction rates 1. initiate the reaction; 2. measure the concentration of the reactant or product at different times as the reaction progresses; 3. plot the concentration as a function of time on a graph; 4. calculate the change in the concentration per unit time.

14. Reaction Rates cont.…  There is another big idea for rates of reaction called collision theory.  The collision theory says that as more collisions in a system occur, there will be more combinations of molecules bouncing into each other.  If there are a higher number of collisions in a system, more combinations of molecules can occur. The reaction will go faster and the rate of that reaction will be higher.

15. Concentration  More substance in a system, greater chance that molecules will collide and speed up the rate of the reaction  less of something, there will be fewer collisions and the reaction will probably happen at a slower speed. Sometimes when you are in a chemistry lab, you will add one solution to another. When you want the rate of reaction to be slower, you will add only a few drops at a time instead of the entire beaker.!!!

16. Temperature  Raising the temperature of a system, the molecules bounce around a lot more because they have more energy.  Bouncing results in collision and collision in combinations. When you lower the temperature, the molecules are slower and collide less.  That temperature drop lowers the rate of the reaction.  Back to the chemistry lab! Sometimes you will mix solutions in ice so that the temperature of the system stays cold and the rate of reaction is slower.

17. Pressure:  Pressure affects the rate of reaction, especially when you look at gases.  When you increase the pressure, the molecules have less space in which they can move. That greater density of molecules increases the number of collisions.  When you decrease the pressure, molecules don't hit each other as often. The lower pressure decreases the rate of reaction.

18. Interaction of all factors

19. Reaction Rates cont.… – The change in the concentration of either the reactant or the product over a period of time. – For a simple reaction ( A  B), rate =  [B] = –  [A]  t  t – Square brackets indicate concentration; and  means “change in.” – Concentration of A decreases with time; and the concentration of B increases with time.

20. Rate Laws cont.…. Reaction orders – For a reaction with the general equation aA + bB  cC + dD, the experimentally determined rate law has the form rate = k[A] m [B] n. – The proportionality constant, k, is called the rate constant. 1. Value is characteristic of the reaction and reaction conditions 2. A given reaction has a particular value of the rate constant under a given set of conditions, such as temperature, pressure, and solvent

21. Orders of Reaction – summary In a zero order reaction the rate=k since anything to the power of 0 is 1. Therefore the rate of reaction does not change over time and the [A] (for example) changes linearly. In a first order reaction, the rate and concentration are proportional. This means that if the concentration is doubled, the rate will double. And finally, in a second order reaction, if the concentration is doubled, the rate will increase by a factor of 4 (2 2 ). The speed at which the [A] changes

22. A zero-order reaction is a reaction that proceeds at a rate that is independent of reactant concentration. Typically with increasing or decreasing reactants concentrations will not affecting the observed reaction. This means that the rate of the reaction is equal to the rate constant, k, of that reaction. One can write their rate in a form such that the exponent of the reactant in the rate law is 0 rate = –  [A] = k [reactant] 0 = k(1) = k  t Zero-Order Reactions

23. Reaction rate is directly proportional to the concentration of one of the reactants- depends linearly only on one reactant concentration. – Have the general form A  products – Differential rate for a first-order reaction is rate = –  [A] = k[A]  t – If the concentration of A is doubled, the rate of the reaction doubles; if the concentration of A is increased by a factor of 10, the rate increases by a factor of 10 – Units of a first-order rate constant are inverse seconds, s –1 – First-order reactions are very common First-Order Reactions

24. Two kinds of second-order reaction 1. The simplest kind of second-order reaction is one whose rate is proportional to the square of the concentration of the reactant and has the form 2A  products. – Differential rate law is rate = –  [A] = k[A] 2 2  t – Doubling the concentration of A quadruples the rate of the reaction – Units of rate constant is M –1  s –1 or L/mol  s Second-Order Reactions

25. 2. The second kind has a rate that is proportional to the product of the concentrations of two reactants and has the form A + B  products. – Reaction is first order in A and first order in B – Differential rate law for the reaction is rate = –  [A] = –  [B] = k[A] [B]  t  t – Reaction is first order both in A and in B and has an overall reaction order of 2 Second-Order Reactions

27. Half-Lives  Another approach to describe reaction rates is based on the time required for the concentration of a reactant to decrease to one-half its initial value.  The period of time is called the half-life of the reaction, written as t ½.  The half-life of a reaction is the time required for the reactant concentration to decrease from [A] 0 to [A] 0 /2.  If two reactions have the same order, the faster reaction will have a shorter half-life and the slower reaction will have a longer half-life.

28. Radioactive Decay Rates  Radioactivity, or radioactive decay, is the emission of a particle or a photon, that results from the spontaneous decomposition of the unstable nucleus of an atom.  The rate of radioactive decay is an intrinsic property of each radioactive isotope, independent of the chemical and physical form of the radioactive isotope.

29. Radioactive Decay Rates Because radioactive decay is a first-order process, the time required for half of the nuclei in any sample of a radioactive isotope to decay is a constant, called the half-life of the isotope Isotopes with a short half-life decay more rapidly, undergoing a greater number of radioactive decays per unit time than do isotopes with a long half-life

30. Radioisotope Dating Techniques Using the half-lives of isotopes, one can estimate the ages of geological and archaeological artifacts. Techniques that have been developed for this application are known as radioisotope dating techniques. Radiometric or radioactive dating is a technique used to date materials such as rocks or carbon, usually based on a comparison between the observed abundance of a naturally occurring radioactive isotope and its decay products, using known decay rates.

31. Activation Energy  A minimum energy (activation energy, E a ) is required for a collision between molecules to result in a chemical reaction.  Reacting molecules must have enough energy to overcome electrostatic repulsion and a minimum amount of energy to break chemical bonds so that new ones may be formed.

32. Catalysis – Substances that increase the rate of a chemical reaction without being consumed in the process – A catalyst does not appear in the overall stoichiometry of the reaction it catalyzes, but it must appear in at least one of the elementary steps in the mechanism for the catalyzed reaction – Catalyzed pathway has a lower E a (activation energy ) – Because of its lower E a, the rate of a catalyzed reaction is faster than the rate of the uncatalyzed reaction at the same temperature

34. Catalysis A catalyst decreases the height of the energy barrier, and its presence increases the rates of both the forward and the reverse reactions by the same amount There are three major classes of catalysts 1. Heterogeneous catalysts 2. Homogeneous catalysts 3. Enzymes

35. Enzymes Enzymes are catalysts that occur naturally in living organisms and are almost all protein molecules A reactant in an enzyme-catalyzed reaction is called a substrate. Enzymes can increase reaction rates by enormous factors and tend to be very specific, typically producing only a single product in quantitative yield.

36. Enzymes Enzymes are expensive, and often cease functioning at temperatures higher than 37ºC, and have limited stability in solution. Enzyme inhibitors cause a decrease in the rate of an enzyme-catalyzed reaction by binding to a specific portion of an enzyme and thus slowing or preventing a reaction from occurring.

37. Readings …  First,follow the slides and get more info from the Book… Homework 4 topic (submit by 20 th May):  How we determine the age of dinosaurs ? Write one page essay.