The degradation of CF3CH2F (an HFC) by OH radicals in the troposphere is first order in each reactant and has a rate constant of k=1.6×108M−1s−1 at 4∘C.
If the tropospheric concentrations of OH and CF3CH2F are 8.1×105 and 6.3×108 molecules cm−3, respectively, what is the rate of reaction at this temperature in M/s?
Express your answer using two significant figures.
If you can show work please that would be fantastic!
rate = k(CF3CH2F)(OH)
Okay... I already know that. I plugged in all the numbers and I get 8.2*10^22, which is not the right answer. Can someone show the work for the problem please?
molecules is not M. I would convert molecules to mol which will be mols/cc and convert that to mols/L which is M.
To calculate the rate of the reaction at a given temperature, we can use the first-order rate equation:
Rate = k * [OH] * [CF3CH2F]
Given:
Rate constant (k) = 1.6×10^8 M^−1 s^−1
[OH] = 8.1×10^5 molecules cm^−3
[CF3CH2F] = 6.3×10^8 molecules cm^−3
First, we need to convert the concentrations from molecules cm^−3 to M (molar concentration). To do this, we'll use Avogadro's number (6.022×10^23 molecules/mol) and convert cm^−3 to L (liters).
[OH] in M = ([OH] in molecules cm^−3) / (Avogadro's number * 1000 cm^3/L)
[OH] in M = (8.1×10^5) / (6.022×10^23 * 1000)
[CF3CH2F] in M = ([CF3CH2F] in molecules cm^−3) / (Avogadro's number * 1000 cm^3/L)
[CF3CH2F] in M = (6.3×10^8) / (6.022×10^23 * 1000)
Now, we can plug in these values into the rate equation:
Rate = (1.6×10^8 M^−1 s^−1) * ([OH] in M) * ([CF3CH2F] in M)
Rate = (1.6×10^8) * ((8.1×10^5) / (6.022×10^23 * 1000)) * ((6.3×10^8) / (6.022×10^23 * 1000))
After solving this expression, we get the rate of reaction in M/s.