Under certain conditions the rate of this reaction is zero order in hydrogen iodide with a rate constant of 0.0017 M.s¹: Suppose a 300. mL flask is charged under these conditions with 450. mmol of hydrogen iodide. How much is left 200. s later? You may assume no other reaction is important. Be sure your answer has a unit symbol, if necessary, and round it to 2 significant digits. 2 HI (9) H2(g)+1₂(g) B = X x10

Under certain conditions the rate of this reaction is zero order in hydrogen iodide with a rate constant of 0.0017 M.s¹: Suppose a 300. mL flask is charged under these conditions with 450. mmol of hydrogen iodide. How much is left 200. s later? You may assume no other reaction is important. Be sure your answer has a unit symbol, if necessary, and round it to 2 significant digits. 2 HI (9) H2(g)+1₂(g) B = X x10

Chemistry

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Chapter12: Chemical Kinetics

Section: Chapter Questions

Problem 3RQ: One experimental procedure that can be used to determine the rate law of a reaction is the method of...

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The label on a bottle of 3% (by volume) hydrogen peroxide, H2O2, purchased at a grocery store, states that the solution should be stored in a cool, dark place. H2O2decomposes slowly over time, and the rate of decomposition increases with an increase in temperature and in the presence of light. However, the rate of decomposition increases dramatically if a small amount of powdered MnO- is added to the solution. The decomposition products are H2O and O2. MnO2 is not consumed in the reaction. Write the equation for the decomposition of H2O2. What role does MnO2 play? In the chemistry lab, a student substituted a chunk of MnO2 for the powdered compound. The reaction rate was not appreciably increased. WTiat is one possible explanation for this observation? Is MnO2 part of the stoichiometry of the decomposition of H2O2?
Candle wax is a mixture of hydrocarbons. In the reaction of oxygen with candle w ax in Figure 11.2, the rate of consumption of oxygen decreased with time after the flask was covered, and eventually' the flame went out. From the perspective of the kinetic-molecular theory, describe what is happening in the flask. FIGURE 11.2 When a candle burns in a closed container, the flame will diminish and eventually go out. As the amount of oxygen present decreases, the rate of combustion will also decrease. Eventually, the rate of combustion is no longer sufficient to sustain the flame even though there is still some oxygen present in the vessel.
The 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.
Iodomethane (CH3I) is a commonly used reagent in organic chemistry. When used properly, this reagent allows chemists to introduce methyl groups in many different useful applications. The chemical does pose a risk as a carcinogen, possibly owing to iodomethanes ability to react with portions of the DNA strand (if they were to come in contact). Consider the following hypothetical initial rates data: [DNA]0 ( mol/L) [CH3I]0 ( mol/L) Initial Rate (mol/Ls) 0.100 0.100 3.20 104 0.100 0.200 6.40 104 0.200 0.200 1.28 103 Which of the following could be a possible mechanism to explain the initial rate data? MechanismIDNA+CH3IDNACH3++IMechanismIICH3ICH3++ISlowDNA+CH3+DNACH3+Fast
At 500 K in the presence of a copper surface, ethanol decomposes according to the equation C2H5OH(g)CH3CHO(g)+H2(g) The pressure of C2H5OH was measured as a function of time and the following data were obtained: Time(s) PC2H5OH(torr) 0 250. 100. 237 200. 224 300. 211 400. 198 500. 185 Since the pressure of a gas is directly proportional to the concentration of gas, we can express the rate law for a gaseous reaction in terms of partial pressures. Using the above data, deduce the rate law, the integrated rate law, and the value of the rate constant, all in terms of pressure units in atm and time in seconds. Predict the pressure of C2H5OH after 900. s from the start of the reaction. (Hint: To determine the order of the reaction with respect to C2H5OH, compare how the pressure of C2H5OH decreases with each time listing.)
One experimental procedure that can be used to determine the rate law of a reaction is the method of initial rates. What data are gathered in the method of initial rates, and how are these data manipulated to determine k and the orders of the species in the rate law? Are the units for k. the rate constant, the same for all rate laws? Explain. If a reaction is first order in A, what happens to the rate if [A] is tripled? If the initial rate for a reaction increases by a factor of 16 when [A] is quadrupled, what is the order of n? If a reaction is third order in A and [A] is doubled, what happens to the initial rate? If a reaction is zero order, what effect does [A] have on the initial rate of a reaction?
Give at least two physical properties that might be used to determine the rate of a reaction.
In Exercise 11.39, if the initial concentration of N2Oj is 0.100 .\1. how long will it take for the concentration to drop to 0.0100 times its original value? The decomposition of N2O5 in solution in carbon tetrachloride is a first-order reaction: 2N2O5—»4NO2 + O2 The rate constant at a given temperature is found to be 5.25 X 10-4 s-’. If the initial concentration of N2O5 is 0.200 M, what is its concentration after exactly 10 minutes have passed?
Ozone, O3, in the Earths upper atmosphere decomposes according to the equation 2 O3(g) 3 O2(g) The mechanism of the reaction is thought to proceed through an initial fast, reversible step followed by a slow, second step. Step 1: Fast, reversible O3(g) O2(g) + O(g) Step 2: Slow O3(g) + O(g) 2 O2(g) (a) Which of the steps is rate-determining? (b) Write the rate equation for the rate-determining step
Nitrogen monoxide reacts with oxygen to give nitrogen dioxide. 2NO(g)+O2(g)2NO2(g) The rate law is [NO]/t = k[NO]2[O2], where the rate constant is 1.16 103 L2/(mol2 s) at 339oC. A vessel contains NO and O2 at 339oC. The initial partial pressures of NO and O2 arc 155 mmHg and 345 mmHg, respectively. What is the rate of decrease of partial pressure of NO (in mmHg per second)? (Hint: From the ideal gas law, obtain an expression for the molar concentration of a particular gas in terms of its partial pressure.)

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