3) It takes $7.21 imes 10^{-19} \mathrm{~J}$ of energy to remove an electron from an iron atom.
a) What is the maximum wavelength of light that can do this?
b) What is the ionization energy of iron in $\mathrm{kJ} / \mathrm{mol}$ ?

Question

3) It takes $7.21 	imes 10^{-19} \mathrm{~J}$ of energy to remove an electron from an iron atom.
a) What is the maximum wavelength of light that can do this?
b) What is the ionization energy of iron in $\mathrm{kJ} / \mathrm{mol}$ ?

Accepted Answer

***A*** ∻Key Concept∻ ⚹Energy and wavelength relationship⚹ ∻Explanation∻ ⚹The maximum wavelength of light that can remove an electron from an atom is inversely proportional to the energy required to remove the electron, as described by the equation $E = \frac{hc}{\lambda}$, where $E$ is the energy, $h$ is Planck's constant, $c$ is the speed of light, and $\lambda$ is the wavelength.⚹ ***B*** ∻Key Concept∻ ⚹Ionization energy calculation⚹ ∻Explanation∻ ⚹The ionization energy in kJ/mol can be calculated by converting the given energy from joules to kilojoules and then multiplying by Avogadro's number ($6.022 imes 10^{23} ext{ mol}^{-1}$) to convert from energy per atom to energy per mole.⚹◊What is the relationship between energy and wavelength in terms of ionization energy in Chemistry?⍭ Generate me a similar question◊

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