Chemical Kinetics 1 PDF

MED-102 General Chemistry Chemical Kinetics 1  LOBs covered List the factors that affect the rate of a chemical reaction Explain how various factors affect the rate of a chemical reaction Derive the Rate Law given experimental data  What is a Reaction Rate? The...

MED-102 General Chemistry Chemical Kinetics 1  LOBs covered List the factors that affect the rate of a chemical reaction Explain how various factors affect the rate of a chemical reaction Derive the Rate Law given experimental data  What is a Reaction Rate? The speed of a chemical reaction Can be very fast (e.g. explosion of TNT) Can be very slow (e.g. iron rusting) Can have intermediate speed (e.g. milk going sour) Concentration change Rate = Time change  Factors Affecting the Reaction Rate Physical state of the reactants Reactant concentrations Reaction temperature Presence of a catalyst WATCH: https://www.youtube.com/watch?v=-4HXaUBbv04  Factor #1 Physical state of the reactants – Increase surface area of solid reactants Solid tablet versus powder form – Increase number of successful collisions Collisions must have enough energy Collisions must have correct orientation  Factor #2 Reactant concentrations – Higher molarities of aqueous solutions – Higher partial pressures of gas reactants  Factor #3 Reaction temperature – Higher kinetic energy Higher collision frequency Higher collision intensity  Factor #3 – Revision Slide The plot above shows Boltzmann distributions. They describe the fractions of molecules having a particular kinetic energy. The vertical dashed black line gives the minimum energy required for the reaction occur. It is usually known as the activation energy. The solid blue distribution curve is at lower temperature. The dark blue region to the right of the activation energy shows the fraction of molecules that have sufficient energy to react successfully. The dashed blue distribution curve is at a higher temperature. The blue areas to the right of the activation energy and under the dashed blue curve show the fraction of molecules having enough energy to react successfully. We can clearly see that at higher temperature, more molecule can react successfully.  Factor #4 Presence of a catalyst Substance that alters the reaction rate Fully recovered in initial form at reaction end WATCH: https://www.youtube.com/watch?v=m_9bpZep1QM Measuring the reaction rate H2(g) + I2(g) → 2 HI(g) Measuring the reaction rate – Revision Slide The red curve shows the amount of either H2 or I2 as the reaction proceeds. We see that these reactants are depleted as the reaction proceeds. The slope of the red curve is negative, pointing down. The blue curve shows the amount of HI product as the reaction proceeds. We see that the amount of HI increases gradually. The slope of the blue curve is positive, pointing up. The slope at a given time is found by drawing a right-angle triangle at the target time, and finding Δy/Δx. The slopes of both curves have their maximum magnitudes at Time = 0. After that, the slope magnitudes decrease steadily. Eventually, when the reaction is completed, the slopes will become flat. We can relate the overall rate of the reaction to the slopes. We can measure either the rate of depletion of reactants, or the rate of formation of products. Average Rate of Reaction Concentration change Rate = Time change  Average Rate of Reaction – Revision Slide Concentration change Rate = Time change Let’s do the first Average Rate. The time interval is 10 minutes, as we see from the first column. In this time interval, the moles of A have gone from 1.00 to 0.74, which is a change of 1.00 – 0.74 = 0.26 mol. Therefore, the Average Rate = 0.26 mol / 10 min = 0.026 mol/min. If we continue to do this for every new interval, we get the set of Average Rates that we see in the very last column. We see that the Average Rate goes from 0.026 to 0.020, etc., and there is a steady decrease in the Average Rate. This is not surprising. The Average Rate decreases, and this is consistent with the slope decreasing, or flattening out as the reaction proceeds. Instantaneous Rate of Reaction Tangent to concentration versus time curve We can do this at any desired specific time Initial Rate = Instantaneous Rate at t = 0 5-Minute Break Why does the reaction slow down? As the amounts of reactants are reduced, there are fewer and fewer collisions leading to products The slope of the Concentration versus Time plot decreases with time  Rate Law and Reaction Order If we write the general reaction a A + b B → Products Then we can write an expression showing how the rate depends on the concentrations of the reactants m n Rate = k[A] [B] RATE LAW  Rate Law Rate = k[A]m[B]n k is the rate constant m and n are the reaction orders with respect to each reactant m + n is the overall reaction order It is important to point out that the reaction orders m and n are not related to the stoichiometric coefficients a and b of the balanced chemical equation. We find m and n by doing specific experiments called initial rate experiments.  Examples of rate laws Rate = k[NO]3[O2] – 3rd order in NO, 1st order in O2 – 4th order overall Rate = k[C6H8]2[H2]2 – 2nd order in both reactants – 4th order overall Rate = k[ClO2][OH-]1/2 – 1st order in ClO2, ½ order in OH- – 3/2 overall order Experimental Determination of a Rate Law 2 ClO2(aq) + 2 OH-(aq) → ClO3-(aq) + ClO2-(aq) The initial rate was measured under different initial concentrations Trial [ClO2], M [OH-], M Initial Rate M/s 1 0.060 0.030 0.0248 2 0.020 0.030 0.00276 3 0.020 0.090 0.00828 – Determine the rate law for this reaction. – Determine the value of the rate constant, including its units. Experimental Determination of a Rate Law Trial [ClO2], M [OH-], M Initial Rate M/s 1 0.060 0.030 0.0248 2 0.020 0.030 0.00276 3 0.020 0.090 0.00828 Rate = k[ClO2 ]m[OH− ]n To determine the order m we need to vary [ClO2] while keeping [OH-] fixed To determine the order n we need to vary [OH-] while keeping [ClO2] fixed In Trials 1 and 2 [OH-] is fixed while [ClO2] is tripled. What happens to the Initial Rate? 0.0248/0.00276 = 9. Therefore, 3m = 9. Thus, m = 2 (second order) In Trials 2 and 3 [ClO2] is fixed while [OH-] is tripled. What happens to the Initial Rate? 0.00828/0.00276 = 3. Therefore 3n = 3. Thus, n = 1 (first order) Experimental Determination of a Rate Law Trial [ClO2], M [OH-], M Initial Rate M/s 1 0.060 0.030 0.0248 2 0.020 0.030 0.00276 3 0.020 0.090 0.00828 Rate = k[ClO2 ]2[OH− ]1 In order to find the value of k, we select any one of the three trials. Solve Rate Law for k and use Trial 1 data Rate 0.0248 M/s −2 −1 k= = = 229.6 M s [ClO2 ]2 [OH− ] (0.060 M)2 (0.030 M) Once we have the orders m and n and the value of k, we can calculate the Initial Rate for any desired concentrations  Determining a Rate Law – Revision Slide There is a very nice video on YouTube that explains in much detail how we do these calculations WATCH: https://www.youtube.com/watch?v=lTNidJFIvXU  Summary for Revision The speed of a chemical reaction, or the reaction rate, is given as a concentration change divided by a time change. Four factors affect the reaction rate: the state of the reactants, concentration of reactants, the reaction temperature, and the presence of a catalyst. A successful collision is one that has a minimum of required energy called the activation energy. We can measure the reaction rate either by tracking the rate of depletion of reactants, or the rate of formation of products. A concentration versus time plot shows curves that have specific slopes that change as a function of time. The curve related to the depletion of reactant has a negative slope, and the curve related to the formation of product has a positive slope. We can relate either curve to the overall reaction rate. We can find the average rate of reaction over a finite time interval, or the instantaneous rate of the reaction at a specific time. The rate of a reaction decreases as the reaction proceeds because as reactant amounts decrease there are fewer successful collisions taking place. The Rate Law shows how the rate of the reaction depends on the reactant concentrations. It contains reactant concentrations, reactant orders, and a rate constant. The Rate Law can be determined experimentally only, through initial rate experiments.