Consider the following energy levels of a hypothetical atom: E4 -1.41 x 10^-19 J E3 -4.81 x 10-19 J E2 -1.15 x 10-18 J E1 -1.65 x 10 -18 J (a) What is the wavelength of the photon needed to excite an electron from E1 to E4?

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Consider the following energy levels of a hypothetical atom: E4 -1.41 x 10^-19 J E3 -4.81 x 10-19 J E2 -1.15 x 10-18 J E1 -1.65 x 10 -18 J (a) What is the wavelength of the photon needed to excite an electron from E1 to E4?

Chemistry
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First find the difference in energy between the levels, next use \(E=\dfrac{hc}{\lambda}\)
(b) What is the energy (in joules) a photon must have in order to excite an electron from E2 to E3? (c) When an electron drops from the E3 level to the E1 level, the atom is said to undergo emission. Calculate the wavelength of the photon emitted in this process.
It's pretty much all the same concept, find the energy difference between the energy levels and use Planck's equation.

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Other answers:

\[\frac{ 6.626 x 10^{-34} J/s x 3.0X10^{8}m/s }{ (-1.65 x10^{-18})-(-1.41 x 10^{-19}J) }\]
the energy difference should be positive (just add absolute value bars), but yeah that's good for the first one
0.000000132
1.32 x 10^-7
idk, sorry i can't check your answers, just put sometime into learning how to use your calc, the set up is correct.
I think that's the right answer. I'm confused how to set up b since it's asking for energy instead of wavelength.
it's just asking for the difference in energy
no need to use Planck's equation here
But you would use it for c right?

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