Achieve HW #4

[KDeguzman_Dis3K]

"A liquid is exposed to infrared radiation with a wavelength of 9.26×10−4 cm. Assume that all the radiation is absorbed and converted to heat. How many photons are required for the liquid to absorb 18.08 J of heat?"

I know for this question you need to use the equation E = hc/ lambda and make sure to convert the wavelength from cm to meters and joules to kilojoules.
E= (6.626*10^-34)(3*10^8)/0.00000926 = 2.15 * 10^-20 J
2.15 * 10^-20 J -> 2.15 * 10^-23 kJ
18.08 * 1 photon / 2.15*10^-23 = 8.4 * 10^-23

My answer keeps getting marked wrong, what step of the process is incorrect?

[Lauren Woodward 1I]

Your mistake was converting Joules to Kilojoules before the final step. The energy of a single photon that you calculated first was given in Joules/photon. To find the number of photons using the given value of 18.08, the Joules units have to be able to cancel each other out so only photons will be left over.

[Grace Chang 1E]

Hello!

First, you don't need to change J to kJ; J is used for Planck's constant, so you want to keep it this way (unless it explicitly mentions kJ in any way). Next, if you divide 18.08 by 2.15*10^-20 (which I obtained without converting to kJ), you should get 8.42x10^20. This might've just been a calculator error!

I hope that helps :)