A carbon-carbon bond requires 348 kJ/mol to break. What is the minimum frequency of radiation with enough ene

I'm doing a practice test for my chemistry exam and I'm stuck on this problem. Please help and explain if possible. THANK YOU!!

4. Ultraviolet radiation and radiation of shorter wavelengths can damage biological molecules
because they carry enough energy to break bonds within the molecules. A carbon-carbon
bond requires 348 kJ/mol to break. What is the minimum frequency of radiation with
enough energy to break carbon-carbon bonds?
A. 8.72 x 1011 Hz
B. 5.78 x 10–19 Hz
C. 5.78 x 10–22 Hz
D. 3.44 x 10–7 Hz
E. 8.72 x 1014 Hz

The first step is to find out how much energy it takes to break a single carbon-carbon bond. In the example, they give you the energy per MOL, but you want the energy per single bond. This means simple division by Avogadros number:

348000 Joules/mol / (6.022 x 10^23 bonds / mol) = E

Where kilojoules has been converted to Joules because SI units are better for multistep calculations, and E is the energy of a single bond.

Radiation is simply light, which is the same as packets of photons. Now that you have the energy of the bond, you need a photon with the same energy as the bond to break it. What is the energy of a photon?

E = hf

where h is Planck's Constant (6.626 x 10^-34 J*s) and f is the frequency in Hz, which we need to solve for:

E/h = f

Putting it all together, we have

348000 Joules/mol / (6.022 x 10^23 bonds / mol) / (6.626 x 10^-34 J*s) = 8.72 x 10^14 Hz.

Hope this helps! :)

Yes, this is a common trick on exams. It's important to pay attention to units. If you ever see "mol" units in a problem, take a minute to consider whether or not you might need to use Avogadro's number.

Report Abuse

As a general rule, if you are doing conversions or math in general between different atoms and/or molecules, you want to work in units of moles.