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Chapter 15

Chapter 15. Kinetics. Kinetics. Deals with the rate of chemical reactions Reaction mechanism – steps that a reaction takes Haber Process: uses iron oxide as a catalyst N 2 + 3H 2  2NH 3 Spontaneous – will occur on its own – not necessarily fast. Reaction Rate.

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Chapter 15

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  1. Chapter 15 Kinetics

  2. Kinetics • Deals with the rate of chemical reactions • Reaction mechanism – steps that a reaction takes • Haber Process: uses iron oxide as a catalyst • N2 + 3H2 2NH3 • Spontaneous – will occur on its own – not necessarily fast

  3. Reaction Rate • Change in the concentration of a reactant or product over time • Rate = Δ[A]/Δt • A- Molarity (final – initial) • Units – moles/L/s or M/s • Reactants can have neg. rates – change to a +

  4. Average rates – M changes over a given time period • Instantaneous rate – found at any given second. • Rate is affected by 2 factors: • 1. Coefficients in the balanced equation. • 2. Time

  5. Rate Laws • Concerned only with the concentration of the reactants • Rate = k[A]n • k = rate constant; n = order – must be determined by experimental data • Rate is directly related to concentration

  6. 2 types of rate laws: • 1. differential - (rate law) – shows how reaction depends on concentration. • 2. Integrated – shows how concentration depends on time

  7. Determining the order of the differential rate law: • Order is not the same as the coefficient in the balanced equation • Following method only works for differential rates (concentration and rate are given) • First you must find the order

  8. To determine the order: • If there are 2 or more reactants: • Find where one of the reactants concentrations remains the same will the other reactant’s concentration changes. • Compare rates by the division method demonstrated. • It is easier if the larger concentration goes on top. • But whichever experiment goes on top for the concentration, the rate for the same experiment must be on top in the other side.

  9. Order: • 1st order – concentration doubles – rate doubles • 2nd order – concentration doubles – rate quadruples • 3rd order – concentration doubles – rate increases by a factor of 8 • 0 order – concentration doubles – rate stays the same

  10. Do the same thing for each reactant in the table. Then put them together in one rate law. • Overall reaction order – sum of all the orders • General pattern for units: • L(n-1)/mol(n-1)*s • N is the overall reaction order • To calculate the rate constant – plug #’s of any row back in the equation and solve for k

  11. Integrated rate laws: • Use this method when concentration and time are given • Must graph integrated rate laws to determine the order. • Slope is related to the rate constant.

  12. To determine the concentration at a certain time, plug back into the appropriate equation for the rate law to solve. The initial concentration is the M at time = 0. • To determine the rate, plug back into the rate law to solve. • Pseudo- first order rate law – if more than 1 reactant • Can determine the rate law by graphing one reactant at a time. • Then put them together in one rate law.

  13. Factors affecting reaction rates: • 1. Concentration – (from rate laws) • 2. Temperature – speed up when temp. is increased • 3. Catalyst – speeds up a reaction without being used up itself

  14. How does temperature affect the reaction rate? • Collision model – molecules must collide to react • Higher the temp – higher the KE – more collisions

  15. Two requirements for reactants to collide: • 1. The molecules must be oriented correctly to allow bond formation • 2. The collisions must have enough energy to produce a reaction

  16. Activation Energy - Ea • Energy required to break a chemical bond and produce a new one in a chemical reaction • To be broken the KE must be high enough to overcome its bond and convert it to PE in the new product • If the Ea is low – reaction happens fast • If the Ea is high – reaction happens slow

  17. Arrhenius Equation • Ln(k) = -Ea/R(1/T) +ln(A) • Ln(k) – y axis; (1/T) – x axis • When graphed a straight line should form. • R= 8.31J/mol*K • Units for Ea = J/mol • Slope = -Ea/R

  18. Ln(k2/k1) = Ea/R(1/T1 – 1/T2) • **Use this form if temperature changes** • Catalyst – lower the Ea • Does not affect the overall energy difference of the reaction • 2 types of catalyst: • A. Homogeneous – same phase as the reacting molecule • B. Heterogenous – different phase (usually 2 gases being absorbed on the surface of a solid)

  19. Reaction Mechanisms: • Series of steps for a chemical reaction • Must meet 2 requirements: • 1. Sum of the elementary steps must give overall balanced equation for the reaction • 2. The mechanism must agree with the experimentally determined rate law

  20. Elementary steps – reaction whose rate law can be written by looking at the # of species colliding • 3 types of molecularity (same as the order) • A. Unimolecular – 1 molecule – 1st order • B. Bimolecular – 2 species – 2nd order • C. Termolecular – 3 species – 3rd order – rare since all have to hit at the same time

  21. Intermediate – a species that is neither in the overall reactant or product but is used up during the reaction • To determine the rate law look at the rate determining step. • Rate determining step – in multi-step reaction it is the slowest step

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