posted by Amy on .
A student proposes to use visible spectroscopy to measure the kinetics of the permanganate/oxalic acid reaction. The initial permanganate solution was very dark. For a 1 cm sample cell, the absorbance A = 2,500[MnO4]. As a practical rule of thumb, Beer's law is only accurate below an absorbance value of A = 1.00. This limits the initial permanganate concentration that can be used.
A: What is the maximum permanganate concentration that can be used and still have A = 1.00?
B: If the detection limit for the spectrometer is A = 0.01, what is the minimum permanganate concentration that can be detected?
C: The rate law for the reaction was determined to be:
[MnO4-]0/t = 0.052[mnO4]0^-1.29 * [H2C2O4]0^0.76
If [H2C2O4]0 = 0.30 mol/L and [MnO4] is the maximum it can be for visible spectroscopy, how long would it take for the the permanganate to be used up? Remember the rate law given above was for the average rate for the complete conversion of MnO4- to Mn2+.
A. If A = 2,500(MnO4^-) and you want A to be 1.00, Substitute 1.00 into the above and solve for (MnO4^-).
B. Substitute 0.01 for A and solve for (MnO4^-)
I don't get c.