Finding photon wavelengths that break chemical bonds at 250–800 kJ/mol

[ducmod]

Homework Statement

Hello!
Here is the quote of the exercise:
Excessive exposure to sunlight increases the risk of
skin cancer because some of the photons have
enough energy to break chemical bonds in biological
molecules. These bonds require approximately
250–800 kJ / mol of energy to break. The energy of a
single photon is given by E = hc / lambda where E is the
energy of the photon in J, h is Planck’s constant 6.626 * 10^(-34) J s
and c is the speed of light 3 * 10^8 m/s
Determine which kinds of light
contain enough energy to break chemical bonds in
biological molecules by calculating the total energy
in 1 mol of photons for light of each wavelength.
(a) infrared light (1500 nm)
(b) visible light (500 nm)
(c) ultraviolet light (150 nm)

Of course, everyone knows that it's ultraviolet, but I can't come up with a correct math. Please, explain
my mistakes:
(6.626 * 10^(-34) J s * 3 * 10^8 m/s) / (150 * 10^(-9)) = 1.3252 * 10^(-17)
which I don't see how to fit in 250–800 kJ / mol of energy

Thank you!

 

[Bystander]

in 1 mol of photons

 

[ducmod]

Do you mean that 1.3252 * 10^(-17) is for one photon, and I have to multiply this number by avogadro number?
I didn't quite understand your remark.

 

[Borek]

Do you mean that 1.3252 * 10^(-17) is for one photon, and I have to multiply this number by avogadro number?

That's exactly what you are asked to calculate, aren't you?