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CH_14_6_per_page - Contents 14-1 The Rate of a Chemical...

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Unformatted text preview: Contents 14-1 The Rate of a Chemical Reaction 14-2 Measuring Reaction Rates 14-3 Effect of Concentration on Reaction Rates: The Rate Law 14-8 Theoretical Models for Chemical Kinetics Chapter 14: Chemical Kinetics Dr. Berk & Dr. Koel CHM 25 Spring 2007 1 of 2 of 30 Chemical Kinetics The Rate of a Chemical Reaction We can use thermodynamics to tell if a reaction is product- or reactant-favored. But this gives us no info on HOW FAST reaction goes from reactants to products. KINETICS — the study of REACTION RATES and their relation to the way the reaction proceeds. 3 of 30 Reaction rate = change in concentration of a reactant or product with time. 2 Fe3+(aq) + Sn2+ → 2 Fe2+(aq) + Sn4+(aq) t = 38.5 s Δt = 38.5 s [Fe2+] = 0.0010 M Δ[Fe2+] = (0.0010 –0) M Δ[Fe2+] 0.0010 M = 38.5 s = 2.610-5 M s-1 Rate of formation of Fe2+= Δt 4 of 30 Rates of Chemical Reaction 2 Fe3+(aq) + Sn2+ → 2 Fe2+(aq) + Sn4+(aq) General Rate of Reaction aA+bB→cC+dD Rate of reaction = rate of disappearance of reactants Rate of formation of Sn4+= Δ[Sn4+] Δt = 0.0005 M 38.5 s = 1.310-5 M s-1 =- 1 Δ[B] 1 Δ[A] =b Δt a Δt 1 Δ[Fe3+] Δ[Sn4+] 1 Δ[Fe2+] == Δt Δt 2 Δt 2 = rate of appearance of products = 1 Δ[D] 1 Δ[C] = d Δt c Δt 5 of 30 6 of 30 CHM 25 SP07 1 Determining a Reaction Rate Blue dye is oxidized with bleach. Its concentration decreases with time. The rate — the change in dye conc with time — can be determined from the plot. Determining a Reaction Rate Graphing concentration versus time does not give a straight line because the rate of the reaction changes during the course of the reaction. 2 N2O5 → 4 NO2 + O2 Three “ types”of rates • initial rate • average rate • instantaneous rate 7 of 30 8 of 30 Determining a Reaction Rate Factors Affecting Rates Concentrations and physical state of reactants and products Temperature Catalysts 9 of 30 10 of 30 Effect of Concentration on Reaction Rates: The Rate Law a A + b B….→ g G + h H …. Rate of reaction = k[A]m[B]n … . Rate constant = k Overall order of reaction = m + n + … . 11 of 30 Consider the following reaction: 2 NO (g) + Cl2 (g) → 2 NOCl (g) The experimentally determined rate equation for this reaction is Rate = k [NO]2[Cl2] What is the order in terms of NO? ____ What is the order in terms of Cl2? ____ What is the overall order? ____ 12 of 30 CHM 25 SP07 2 Effect of Concentration on Reaction Rates: The Rate Law Interpreting Rate Laws Rate = k [A]m[B]n[C]p If m = 1, rxn. is 1st order in A Rate = k [A]1 If [A] doubles, then rate goes up by factor of _____ If m = 2, rxn. is 2nd order in A. Rate = k [A]2 Doubling [A] increases rate by ________ If m = 0, rxn. is zero order. Rate = k [A]0 If [A] doubles, rate ________ 13 of 30 14 of 30 Deriving Rate Laws Deriving Rate Laws Rate of rxn = _______________ Here the rate goes up by ______ when initial conc. doubles. Therefore, we say this reaction is _________________ order. Now determine the value of k. Use expt. #3 data— Derive rate law and k for CH3CHO(g) --> CH4(g) + CO(g) from experimental data for rate of disappearance of CH3CHO [CH3CHO] (mol/L) 0.10 0.20 0.30 0.40 Disappear of CH3CHO (mol/L• ) sec 0.020 0.081 0.182 0.318 15 of 30 Expt. 1 2 3 4 Using k you can calc. rate at other values of [CH3CHO] at same T. 16 of 30 EXAMPLE Establishing the Order of a reaction by the Method of Initial Rates. Use the data provided establish the order of the reaction with respect to NH4+ and NO2- and also the overall order of the reaction. NH4 Experiment 1 2 3 +(aq) EXAMPLE NH4+(aq) + NO2-(aq) → N2(g) + H2O(l) Experiment 1 Initial rate (mol/L*s) 1.35 x 10-7 2.70 x 10-7 5.40 x 10-7 17 of 30 Initial [NH4+] 0.100 M 0.100 M 0.200 M Initial [NO2-] 0.0050 M 0.010 M 0.010 M + NO2 -(aq) → N2(g) + H2O(l) Initial [NO2-] 0.0050 M 0.010 M 0.010 M Initial rate (mol/L*s) 1.35 x 10-7 2.70 x 10-7 5.40 x 10-7 Initial [NH4+] 0.100 M 0.100 M 0.200 M 2 3 First consider it in terms of NH4+, keeping the NO2- constant. Now reverse it, keep the NH4+ constant, and look at the NO2-. 18 of 30 CHM 25 SP07 3 Theoretical Models for Chemical Kinetics Collision Theory Activation Energy (Ea) Kinetic-Molecular theory can be used to calculate the collision frequency. gases 1030 collisions per second. If each collision produced a reaction, the rate would be about 106 M s-1. Actual rates are on the order of 104 M s-1. In For a reaction to occur there must be a redistribution of energy sufficient to break certain bonds in the reacting molecule(s). Activation Energy: The Still a very rapid rate. Only a fraction of collisions yield a reaction. 19 of 30 minimum energy above the average kinetic energy that molecules must bring to their collisions for a chemical reaction to occur. 20 of 30 MECHANISMS & Activation Energy Cis Transition state Trans Activation energy barrier Conversion of cis to trans-2-butene requires twisting around the C=C bond. Rate = k [trans-2-butene] 21 of 30 22 of 30 Energy involved in conversion of trans to cis butene energy Activated Complex -266 kJ 4 kJ/mol trans See Figure 15.14 23 of 30 +262 kJ Reaction passes thru a TRANSITION STATE where there is an activated complex that has sufficient energy to become a product. ACTIVATION ENERGY, Ea = energy req’ to form activated complex. d cis Here Ea = 262 kJ/mol 24 of 30 CHM 25 SP07 4 Also note that trans-butene is MORE STABLE than cis-butene by about 4 kJ/mol. Therefore, cis ---> trans is EXOTHERMIC This is the connection between thermodynamics and kinetics. Activation Energy and Temperature Reactions are faster at higher T because a larger fraction of reactant molecules have enough energy to convert to product molecules. In general, In general, differences in differences in activation activation energy cause energy cause 25 of 30 reactions to vary reactions to vary from fast to slow. from fast to slow. 26 of 30 Collision Theory Collision Theory If activation barrier is high, only a few molecules have sufficient kinetic energy and the reaction is slower. As temperature increases, reaction rate increases. Orientation of molecules may be important. 27 of 30 28 of 30 Transition State Theory The activated complex is a hypothetical species lying between reactants and products at a point on the reaction profile called the transition state. 29 of 30 30 of 30 CHM 25 SP07 5 ...
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