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(a)
Interpretation:
Amount of
Concept Introduction:
Order of a reaction: The order of each reactant is represented by the exponential term of the respective reactant present in the rate law and the overall order of the reaction is the sum of all the exponents of all reactants present in the
The integrated rate law for the first order is as follows:
Half–life period: It is the time required for the reactant (substrate) concentration to reduce to the one-half of its initial concentration.
The half–life period for first order reaction is as follows:
(b)
Interpretation:
Time required to decrease
Concept introduction:
Order of a reaction: The order of each reactant is represented by the exponential term of the respective reactant present in the rate law and the overall order of the reaction is the sum of all the exponents of all reactants present in the chemical reaction. The order of the reaction is directly proportional to the concentration of the reactants.
The integrated rate law for the first order is as follows:
Half–life period: It is the time required for the reactant (substrate) concentration to reduce to the one-half of its initial concentration.
The half–life period for first order reaction is as follows:
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Chapter 14 Solutions
Chemistry & Chemical Reactivity
- The 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_forwardThe frequency factor A is 6.31 108 L mol1 s1 and the activation energy is 10. kJ/mol for the gas-phase reaction NO(g)+O3(g)NO2(g)+O2(g) which is important in the chemistry of stratospheric ozone depletion. (a) Calculate the rate constant for this reaction at 370. K. (b) Assuming that this is an elementary reaction, calculate the rate of the reaction at 370. K if [NO] = 0.0010 M and [O3] = 0.00050 M.arrow_forwardSucrose, a sugar, decomposes in acid solution to give glucose and fructose. The reaction is first-order in sucrose, and the rate constant at 25 C is k = 0.21 h1. If the initial concentration of sucrose is 0.010 mol/L, what is its concentration after 5.0 h?arrow_forward
- As with any drug, aspirin (acetylsalicylic acid) must remain in the bloodstream long enough to be effective. Assume that the removal of aspirin from the bloodstream into the urine is a lirst-order reaction, with a half-life of about 3 hours. The instructions on an aspirin bottle say to take 1 or 2 tablets every 4 hours. If a person takes 2 aspirin tablets, how much aspirin remains in the bloodstream when it is time for the second dose? (A standard tablet contains 325 mg of aspirin.)arrow_forwardThe 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_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_forward
- Define stability from both a kinetic and thermodynamic perspective. Give examples to show the differences in these concepts.arrow_forwardWhen boron trifluoride reacts with ammonia, the following reaction occurs: BF3(g)+NH3(g)BF3NH3(g)The following data are obtained at a particular temperature: (a) What is the order of the reaction with respect to BF3, NH3, and overall? (b) Write the rate expression for the reaction. (c) Calculate k for the reaction. (d) When [ BF3 ]=0.533M and NH3=0.300M, what is the rate of the reaction at the temperature of the experiment?arrow_forwardA reaction has two reactants X and Y. What is the order with respect to each reactant and the overall order of the reaction described by the following rate expressions? (a) rate=k1[ X ][ Y ]2 (b) rate=k2[ X ]2 (c) rate=k3[ X ][ Y ] (d) rate=k4arrow_forward
- A reaction has two reactants Q and P. What is the order with respect to each reactant and the overall order of the reaction described by the following rate expressions? (a) rate=k1(b) rate=k2[ P ]2[ Q ] (c) rate=k3[ Q ]2 (d) rate=k4[ P ][ Q ]arrow_forward11.64 HBr is oxidized in the following reaction: 4 HBr(g) + O2(g) —• 2 H2O(g) + 2 Br,(g) A proposed mechanism is HBr + O2 -* HOOBr (slow) HOOBr + HBr — 2 HOBr (fast) HOBr + HBr — H2O + Bn (fast) Show that this mechanism can account for the correct stoichiometry. Identify all intermediates in this mechanism. What is the molecularity of each elementary’ step? Write the rate expression for each elementary' step. Identify the rate-determining step.arrow_forwardThe 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?arrow_forward
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