The bond energy of the O2 molecule is 495.
a) calculate the minimum frequency and max wavelength of light required to break this bond.
b) how many photons of this wavelength would it take in order to break open the bonds of all the oxygen molecules in a 384 mL container filled with oxygen at a pressure of 183.5 KPa and a temp of 36 degrees C? Assume 1 photon is used to break 1 oxygen bond.
i don't even know how to begin to solve this problem?!?!?!
Chemistry - DrBob222, Wednesday, November 7, 2007 at 5:41pm
That is 495 what? J/mol? kJ/mol? kJ/molecule. J/atom? just what?
Use E = hc/lambda where E is Joules, h is Planck's constant = 6.626 x 10^-34, c = speed of light in m/s (3 x 10^8 m/s)
and lambda will be in meters. To determine frequency, use c = lamda x freq.
Use PV = nRT. I would convert kPa to atmospheres, T remember is in Kelvin, R is 0.08205 L*atm/mol*K, V must be in liters. Calculate n and from that determine how many molecules of O2 you have. The way I read the problem 1 photon is required to break 1 bond; therefore, you simply need to know how many bonds there are to be broken.
Post your work if you get stuck.