Question

It takes 498. kJ/mol to break an oxygen-oxygen double bond. Calculate the maximum wavelength of light for which an oxygen-oxygen double bond could be
broken by absorbing a single photon.
Be sure your answer has the correct number of significant digits.

Answer 1

The bond energy required to break the oxygen-oxygen double bond is 240 nm

The energy of the incoming photon = $498 * 10^3 J/mol/6.02 * 10^23 moles$ = $8.27 * 10^-19 J$

Recall that;

E = hc/λ

E = energy of the photon

c = speed of light = $3 * 10^8 m/s$

λ = wavelength of photon

h = Plank's constant = $6.63 * 10^-34 Js$

λ = hc/E

λ = $6.63 * 10^-34 * 3 * 10^8/8.27 * 10^-19$

λ = $2.40 * 10^-7 m$

λ = 240 nm

Learn more: brainly.com/question/24302171

Answer 2

• Energy=498KJ =498×10^3J

We know

$\\ \sf\bull\longmapsto E=hv$

• h=planks constant
• v=frequency

$\\ \sf\bull\longmapsto v=\dfrac{E}{h}$

$\\ \sf\bull\longmapsto v=\dfrac{498\times 10^3J}{6.626\times 10^{-34}Js}$

$\\ \sf\bull\longmapsto v=75.15\times 10^{37}s^{-1}$

$\\ \sf\bull\longmapsto v=7.5\times 10^{36}s^{-1}$

Now

$\\ \sf\bull\longmapsto \lambda=\dfrac{C}{V}$

$\\ \sf\bull\longmapsto \lambda=\dfrac{3\times 10^{8}ms^{-1}}{7.5\times 10^{36}s^{-1}}$

$\\ \sf\bull\longmapsto \lambda=2.5\times 10^{-28}m$