In a first-order reaction, suppose that a quantity X of a reactant is added at regular intervals of time, Δt. At first the amount of reactant in the system builds up; eventually, however, it levels off at a saturation value given by the expression a saturation value
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Chemistry: Principles and Reactions
- The Raschig reaction produces the industrially important reducing agent hydrazine, N2H4, from ammonia, NH3, and hypochlorite ion, OCl−, in basic aqueous solution. A proposed mechanism is Step 1: Step 2: Step 3: What is the overall stoichiometric equation? Which step is rate-limiting? What reaction intermediates are involved? What rate law is predicted by this mechanism?arrow_forwardAt 573 K, gaseous NO2(g) decomposes, forming NO(g) and O2(g). If a vessel containing NO2(g) has an initial concentration of 1.9 102 mol/L, how long will it take for 75% of the NO2(g) to decompose? The decomposition of NO2(g) is second-order in the reactant and the rate constant for this reaction, at 573 K, is 1.1 L/mol s.arrow_forward(Section 11-5) A rule of thumb is that for a typical reaction, if concentrations are unchanged, a 10-K rise in temperature increases the reaction rate by two to four times. Use an average increase of three times to answer the questions below. (a) What is the approximate activation energy of a typical chemical reaction at 298 K? (b) If a catalyst increases a chemical reactions rate by providing a mechanism that has a lower activation energy, then what change do you expect a 10-K increase in temperature to make in the rate of a reaction whose uncatalyzed activation energy of 75 kJ/mol has been lowered to one half this value (at 298 K) by addition of a catalyst?arrow_forward
- Gaseous azomethane (CH3N2CH3) decomposes to ethane and nitrogen when heated: CH3N2CH3(g) CH3CH3(g) + N2(g) The decomposition of azomethane is a first-order reaction with k = 3.6 104 s1 at 600 K. (a) A sample of gaseous CH3N2CH3 is placed in a flask and heated at 600 K for 150 seconds. What fraction of the initial sample remains after this time? (b) How long must a sample be heated so that 99% of the sample has decomposed?arrow_forwardMany biochemical reactions are catalyzed by acids. A typical mechanism consistent with the experimental results (in which HA is the acid and X is the reactant) is Step 1: Step 2: Step 3: Derive the rate law from this mechanism. Determine the order of reaction with respect to HA. Determine how doubling the concentration of HA would affect the rate of the reaction.arrow_forwardThe hydrolysis of the sugar sucrose to the sugars glucose and fructose, C12H22O11+H2OC6H12O6+C6H12O6 follows a first-order rate equation for the disappearance of sucrose: Rate =k[C12H22O11] (The products of the reaction, glucose and fructose, have the same molecular formulas but differ in the arrangement of the atoms in their molecules.) (a) In neutral solution, k=2.11011s1 at 27 C and 8.51011s1 at 37 C. Determine the activation energy, the frequency factor, and the rate constant for this equation at 47 C (assuming the kinetics remain consistent with the Arrhenius equation at this temperature). (b) When a solution of sucrose with an initial concentration of 0.150 M reaches equilibrium, the concentration of sucrose is 1.65107M . How long will it take the solution to reach equilibrium at 27 C in the absence of a catalyst? Because the concentration of sucrose at equilibrium is so low, assume that the reaction is irreversible. (c) Why does assuming that the reaction is irreversible simplify the calculation in pan (b)?arrow_forward
- The decomposition of iodoethane in the gas phase proceeds according to the following equation: C2H5I(g)C2H4(g)+HI(g) At 660. K, k = 7.2 104 sl; at 720. K, k = 1.7 102 sl. What is the value of the rate constant for this first-order decomposition at 325C? If the initial pressure of iodoethane is 894 torr at 245C, what is the pressure of iodoethane after three half-lives?arrow_forwardYou are studying the kinetics of the reaction H2(g) + F2(g) 2HF(g) and you wish to determine a mechanism for the reaction. You run the reaction twice by keeping one reactant at a much higher pressure than the other reactant (this lower-pressure reactant begins at 1.000 atm). Unfortunately, you neglect to record which reactant was at the higher pressure, and you forget which it was later. Your data for the first experiment are: Pressure of HF (atm) Time(min) 0 0 0.300 30.0 0.600 65.8 0.900 110.4 1.200 169.1 1.500 255.9 When you ran the second experiment (in which the higher pressure reactant was run at a much higher pressure), you determine the values of the apparent rate constants to be the same. It also turns out that you find data taken from another person in the lab. This individual found that the reaction proceeds 40.0 times faster at 55C than at 35C. You also know, from the energy-level diagram, that there are three steps to the mechanism, and the first step has the highest activation energy. You look up the bond energies of the species involved and they are (in kJ/mol): H8H (432), F8F (154), and H8F (565). a. Sketch an energy-level diagram (qualitative) that is consistent with the one described previously. Hint: See Exercise 106. b. Develop a reasonable mechanism for the reaction. c. Which reactant was limiting in the experiments?arrow_forwardThe initial rate for a reaction is equal to the slope of the tangent line at t 0 in a plot of [A] versus time. From calculus, initial rate = d[A]dt . Therefore. the differential rate law for a reaction is Rate = d[A]dt=k[A]n. Assuming you have some calculus in your background, derive the zero-, first-, and second-order integrated rate laws using the differential rate law.arrow_forward
- The decomposition of ozone is a second-order reaction with a rate constant of 30.6 atm1 s1 at 95 C. 2O3(g)3O2(g) If ozone is originally present at a partial pressure of 21 torr, calculate the length of time needed for the ozone pressure to decrease to 1.0 torr.arrow_forward11.32 The following experimental data were obtained for the reaction 2A + 3 B—C + 2D [A](mol L 1) [B](mol L ’) Rate = A(C]/Af (mol L-1 s-1) 0.127 0.15 0.033 0.127 0.30 0.132 0.255 0.15 0.066 Determine the reaction order for each reactant and the value of the rate constant.arrow_forwardThe initial concentration of the reactant in a tirst-order reaction A —» products is 0.64 rnol/L and the half-life is 30.0 s. Calculate the concentration of the reactant exactly 60 s after initiation of the reaction. How long would it take for the concentration of the reactant to drop to one-eighth its initial value? How long would it take for the concentration of the reactant to drop to 0.040 mol/L?arrow_forward
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