Certain atoms emit photons of light with an energy of 3.820 ✕ 10−19 j. calculate the frequency (in hz) and wavelength (in nm) of one of these photons.

frequency hz

and

wavelength nm

what is the total energy (in kj) in 1 mole of these photons?

kj

ν = 5,765x10¹⁴ Hz

λ = 520 nm

230,0 kJ/mole

Explanation:

To convert energy yo frequency you need to use:

E = hν

Where E is energy (3,820x10⁻¹⁹ J)

h is Planck's constant (6,626x10⁻³⁴ Js)

And ν is frequency, replacing ν = 5,765x10¹⁴ s⁻¹ ≡ 5,765x10¹⁴ Hz

To convert frequency to wavelength:

c = λν

Where s is speed of light (2,998x10⁸ ms⁻¹)

ν is frequency (5,765x10¹⁴ s⁻¹)

And λ is wavelength, replacing: λ = 5,200x10⁻⁷ ≡ 520 nm

If 1 photon produce 3,820x10⁻¹⁹ J, in mole of photons produce:

3,820x10⁻¹⁹ J × = 230040 J/mole ≡ 230,0 kJ/mole

I hope it helps!

answer: mercury

explanation:

d: the voltage of the secondary coil is less than that of the primary coil.

im not suree sorry if wrong