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Kinetics Question

  1. Feb 19, 2014 #1
    1. The problem statement, all variables and given/known data
    Consider the following reaction at 282 K.

    2 A + 2 B → C + D​

    where rate = rate=k[A]2. An experiment was performed for a certain number of seconds where [A]0 = 0.000303 M and 0 = 1.63 M. A plot of ln[A] vs time had a slope of -7.27. What will the rate of this reaction be when [A] = = 0.503 M?

    Rate (M/s)=

    The correct answer is 0.348 M/s.

    2. Relevant equations
    I'm not sure where to start here, but I tried using these equations, because I was given the slope, and the slope for ln[A] vs time is given in a first order reaction:

    ln[A]t = -kt + ln[A]o
    Rate = k[A]
    k = -slope

    3. The attempt at a solution
    Basically I just tried plugging into k = 7.27 into rate=k[A]2 were [A] and are both 0.503 M.

    Thank you for the help!
  2. jcsd
  3. Feb 19, 2014 #2
    Look at the values of [A] and at the beginning of the experiment. From the specified stoichiometry, how much do you think the concentration of is going to change relative to its initial value during the experiment? Will this really be significant?

  4. Feb 19, 2014 #3
    Oh, do I have to use an ICE table?
  5. Feb 20, 2014 #4
    I don't know what an ICE table is, but, whatever it is, you don't need to use it on this problem.

  6. Feb 20, 2014 #5
    ICE as in initial concentration, change in concentration, and end concentration.

    How do you solve this problem?
  7. Feb 20, 2014 #6
    The solution to this problem starts out by reconsidering the questions I asked in post #2. Let me ask in another way. If B had reacted with all the A in the experiment, what would its final concentration have been? What percentage change would this have made in the concentration of B?

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