CHEM 162: Gilbert Chapter 15 page 1 Chapter 15: Chemical Kinetics Problems: 15.1-15.74, 15.81-15.115, 15.117-15.118, 15.120, 15.122-15.129, 15.133, 15.137 Consider the formation of rust (or oxidation of iron). a. It can occur over a period of several years (e.g., the rusty winch below). b. It can occur quite rapidly (e.g., the combustion of iron in pure oxygen). 15.1 CARS, TRUCKS, AND AIR QUALITY The increased use of automobiles in the last 50 years has created air quality problems, often in the form of photochemical smog (smoke fog). The brownish gas we often equate with smog (e.g. see layer over Los Angeles in the photo at the right) is actually NO2 gas, which has a reddish-brown color at high concentrations: NO2 forms from the following series of reactions: step 1: N2(g) + O2(g) → 2 NO(g) ΔH°=180.6 kJ step 2: 2 NO(g) + O2(g) → 2 NO2(g) ΔH°= -114.2 kJ Although N2 and O2 are abundant in air, the formation of NO requires very high temperatures, so it only forms inside internal combustion engines and fossil fuel power plants or when lightning strikes. When NO escapes into the atmosphere, it reacts with more O2 to form NO2, usually with the help of energy provided by sunlight—hence, the name “photochemical” smog.CHEM 162: Gilbert Chapter 15 page 2 CHEMICAL KINETICS AND COLLISION THEORY chemical kinetics: study of the factors that influence reaction rates—i.e., how quickly reactions occur collision theory (or collision model): molecules must collide to react Now, consider the reverse reaction or decomposition of NO2: 2 NO2(g) → 2 NO(g) + O2(g). System before any reactions occur System after a few NO2 molecules react Ex. 1: a. As the reaction proceeds, the [NO2] __________. ↑ ↓ stays the same b. As the reaction proceeds, the [NO] __________. ↑ ↓ stays the same c. As the reaction proceeds, the [O2] __________. ↑ ↓ stays the same d. As the [NO2] decreases, is the reaction more likely or less likely to occur? Explain. e. As the [NO] and [O2] increase, is the reaction more likely or less likely to occur? Explain.CHEM 162: Gilbert Chapter 15 page 3 15.2 REACTION RATES We can define the rate of a reaction quantitatively as the change in concentration of a reactant or product per unit time: ratetime in change ionconcentrat product) (orreactant in change= t [product] or [reactant] ΔΔΔ= Note: A reaction rate must be positive since it expresses how quickly a reaction occurs. For example, consider the decomposition of NO2, a brown gas responsible for smog: 2 NO2(g) → 2 NO(g) + O2(g) The reaction rate can be expressed in terms of the formation of a product: rate Δt[NO]Δ = or rate Δt][OΔ 2 = The reaction rate can also be expressed in terms of the consumption of a reactant: However, because the concentration of NO2 decreases, Δ[NO2] is negative, so a negative sign is used since reaction rates are positive. rate Δt][NOΔ 2−= Consider the following data on reactant and product concentrations for the reaction at 300°C: 2 NO2(g) → 2 NO(g) + O2(g) Ex. 1: Consider the data at the right to calculate the rate of the reaction with respect to NO2, NO, and O2 in parts a and b below: a. Calculate the reaction rate with respect to NO—i.e., the rate of production of NO— between t=0 s and t=50 s. Show the final answer with the correct units and sig figs. rate ===ifift - t[NO]-[NO] Δt[NO]ΔCHEM 162: Gilbert Chapter 15 page 4 b. Calculate the reaction rate with respect to O2 (i.e., the rate of production of O2) and with respect to NO2 (i.e., rate of consumption of NO2) between t=0 s and t=50 s. c. Fill in the boxes below: rate of consumption of NO2 rate of production of NO rate of production of O2 d. Write the chemical equation for the reaction then explain any differences in the rates above. e. Use the data on the previous page to calculate the reaction rate with respect to NO2, NO, and O2 between t=50 s and t=250 s. rate of consumption of NO2 rate of production of NO rate of production of O2 f. How do these reaction rates compare to those calculated in part a? Thus, the “reaction rate” must be defined with respect to a specific reactant or product or a general reaction rate can be written accounting for the stoichiometric coefficients, as follows: reaction rate Δt][OΔ Δt[NO]Δ 212 ==Δt][NOΔ 212 −=CHEM 162: Gilbert Chapter 15 page 5 Ex. 2 Consider the following reaction: A + 3 B → 2 C Write the reaction rate in terms of all the reactants and products in the reaction. Average Reaction Rates versus Instantaneous Rates Thus, as the reaction proceeds, the concentration of any reactants decreases while the concentration of any products increases. Consider the following plot of the concentrations of reactants and products over time for the reaction, 2 NO2(g) → 2 NO(g) + O2(g), and the reaction rate data with respect to the [O2]: Time period (in s) Δt]2[OΔ 0 s → 50 s 50 s → 100 s 100 s → 150 s 150 s → 200 s 200 s → 250 s 2.2×10-5 M/s 1.4×10-5 M/s 1.0×10-5 M/s 6.0×10-6 M/s 6.0×10-6 M/s As we calculated previously, the reaction rate decreases over time. Note that the reaction rates in the table above are average reaction rates—i.e., they are the average rates over 50-second intervals. By calculating the average reaction rate over shorter and shorter intervals of time, we can determine the rate for a specific instant in time. → an instantaneous rateCHEM 162: Gilbert Chapter 15 page 6 The instantaneous rate can be determined by calculating the slope of a line tangent to the concentration curve at that point in time. Ex. 1: a. The plot above shows the tangent drawn at t=100 s for the NO2 curve and at t=250 s for both NO and O2. Calculate the instantaneous rate at each of these points. b. How do the rates compare? Is this expected? Ex. 2: Consider the following reaction, 2 N2O5(g) → 4 NO2(g) + O2(g). a. Write the rate expression in terms of the consumption of N2O5 and the formation of NO2 and O2. b. If oxygen is forming at a rate of 0.015M/s at a given instant, calculate the rate of consumption of N2O5 and the rate of formation of NO2 at that same instant. c. How do the rates