The rate of reaction is markedly affected by temperature. Chemical Kinetics & k versus T Two theories were developed to explain the temperature effects. 1. 2. 2 UV radiation strikes a CFC molecule causing a chlorine atom to break away the chlorine atom collides with an O 3 molecule Take the destruction of ozone in the stratosphere. CFC: chlorofluorocarbons Watch animation (http://www.ucar.edu/learn/images/o3split.gif) What is happening at the molecular level? forms chlorine monoxide and frees an O 2 molecule a free oxygen atom collides with the chlorine monoxide molecule the two oxygen atoms join and release the chlorine atom the free chlorine atom is free to destroy more O 3 molecules. Go back to the basics! 3 4 1. Molecules must interact. Think collisions! Collisions must occur. 2. The more molecules in a confined space, the more likely collisions will occur. Think probability! Rate is proportional to concentrations of the reactants. 3. Collisions must occur at the right place. Think orientation! If the atom approaches the center O atom of the O 3 molecule, that O atom will not transfer. When O 3 and collide and react, the rate of the reaction will be proportional to the product of the [O 3 ] and []. 5 6 1
is based on three postulates: 1. Chemical reactions in the gas phase are due to the collision of the reactant particles. 4. Bonds are broken and new bonds are formed. Think energy! At the time of collision, bonds are stretched and broken as new bonds are made. Breaking these bonds and rearranging the atoms during the collision requires the input of energy. 2. A collision only results in a reaction if a certain threshold energy is exceeded. 3. A collision only results in a reaction if the colliding particles are correctly This takes us to the &. We will expand on these postulates in the following slides. 7 8 1. Chemical reactions result from collisions of the reactant particles. H 2 (g) + I 2 (g) 2 HI (g) This makes sense! At room temperature and pressure - 10 10 collisions per second If every collision results in the formation of HI, the reaction would be over in much less than a second! In reality, at room temperature, this reaction proceeds very slowly. About one in 10 13 collisions produces a reaction. How do we explain this? 9 10 REVIEW: To describe a collection of gases that is randomly moving about with a range of speed, we use a DISTRIBUTION function. They are known as the Maxwell-Boltzmann distribution of speeds. At lower temperatures, molecular velocity are lower. At higher temperatures, more molecules move with the higher velocity. The shaded area represents the number of molecules traveling fast enough to collide with sufficient energy to react. The area for the Higher temperature curve is bigger. The distributions of molecular speeds of CO 2 molecules at three different temperatures. As the molar mass decreases, the fraction of the molecules moving at higher speed increases. 11 12 2
Activation Energy Reaction MUST overcome an energy threshold called the ACTIVATION ENERGY or activation barrier. Only faster moving molecules will collide with sufficient kinetic energy to overcome the activation barrier. is the fraction of the molecules present in a gas which have energies equal to or in excess of activation energy at a particular temperature. Assume a reaction has an activation energy of 50 kj mol -1 At 20 C (293 K) the value of the fraction is: Raise the temperature by 10 C (ie 303 K) the value of the fraction is: is the fraction of the molecules present in a gas which have energies equal to or in excess of activation energy at a particular temperature. At around room temperature, by increasing the temperature by 10 C, the fraction of the molecules able to react has almost doubled. As a result, this causes the rate of reaction to almost double. 13 14 Orientation or Steric Factor CO (g) + NO 2 (g) NO (g) + CO 2 (g) Above 600K, the reaction proceeds by the collision between CO and NO 2. View animation K + CH 3 I KI + CH 3 Only effective collision will lead to formation of products. Experimentally, This is the correct orientation of K and CH 3 I that will lead to the formation of the products. Rate = k [CO] [NO 2 ] Flash animation: http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/collis11.swf 15 16 Frequency Factor, A, in Arrhenius Equation Orientation or Steric Factor O (g) + N 2 O (g) N 2 (g) + O 2 (g) A is a constant that indicates how many collisions have the correct orientation that could lead to the formation of products. The units of the A are identical to those of the rate constant. (eg For first order reaction, A has the units of s 1, is of the order of 10 13 s -1 ) Oxygen collides with the nitrogen end of N 2 O will result in NO REACTION. Oxygen collides with the oxygen end of N 2 O will result in producing the products. Activated complex forms in the reaction mixture. 17 18 3
To progress from reactants to products, the molecules must collide with enough energy to pass over the activation barrier. PE KE Reaction Coordinate diagram As bond rearrangement occurs, an unstable intermediate species, called an activated complex, exists in the reaction mixture. Transition State PE KE The energy difference between the reactants and the high point of the diagram is the activation energy barrier. E a, activation energy, is the minimum amount of energy required for the reaction to occur. Reaction energy, is the energy difference between the products and the reactants. Macroscopic measurements of E a and k are the result of many individual collisions. As E a increases, k decreases. 19 Activation energy is of the order of kj/mole. The Reaction Coordinate diagram shows the energy profile of an exothermic reaction since P.E. of the products < P.E. of the reactants 20 This diagram shows the energy profile of an endothermic reaction since P.E. of the products > P.E. of the reactants The energy difference between the reactants and the high point of the diagram is the activation energy barrier. E a, activation energy, is the minimum amount of energy required for the reaction to occur. Reaction energy, is the energy difference between the products and the reactants. An endothermic reaction always has a greater E a and a slower rate. CO (g) + NO 2 (g) NO (g) + CO 2 (g) Above 600K, the reaction proceeds by the collision between CO and NO 2. Reactants Activated complex Products 21 22 Summary For an equilibrium (reversible) reaction In this example, Ea reverse > Ea forward The Assumes a collision between reactants needs to happen before a reaction can take place. the majority of collisions do not lead to a reaction, but only those in which the colliding species have: The During a reaction, an increase in potential energy corresponds to an energy barrier over which the reactant molecules must pass if the reaction is to proceed. The transition state occurs at the maximum of this energy barrier. Therefore, k forward > k reverse K eq > 1 Product is favoured for this exothermic reaction. A kinetic energy greater than a certain minimum called the activation energy, Ea The correct spacial orientation with respect to each other. The energy difference between the reactants and the potential energy maximum is referred to as the activation energy. The energy difference between the potential energy of the products and reactants is referred to as the reaction energy. 23 24 4
Summary Arrhenius Equation k is essentially a probability that a collision between reactants will lead to products. k can vary with temperature and pressure. 25 5