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(C) The energy required to break one mole of $H_2$ bonds is $436 \ kJ \ mol^{-1}$.To break a single bond, the energy required $(E)$ is $\frac{436 	im

Vedclass · Accepted Answer

$274$

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(C) The energy required to break one mole of $H_2$ bonds is $436 \ kJ \ mol^{-1}$.To break a single bond, the energy required $(E)$ is $\frac{436 	imes 10^3 \ J \ mol^{-1}}{6.023 	imes 10^{23} \ mol^{-1}} \approx 7.24 	imes 10^{-19} \ J$.The relationship between energy and wavelength is $E = \frac{h \cdot c}{\lambda}$, where $h = 6.626 	imes 10^{-34} \ J \cdot s$ and $c = 3 	imes 10^8 \ m \cdot s^{-1}$.Rearranging for wavelength: $\lambda = \frac{h \cdot c}{E}$.$\lambda = \frac{6.626 	imes 10^{-34} \ J \cdot s 	imes 3 	imes 10^8 \ m \cdot s^{-1}}{7.24 	imes 10^{-19} \ J} \approx 2.745 	imes 10^{-7} \ m$.Converting to nanometers: $\lambda \approx 274.5 \ nm \approx 274 \ nm$.

The energy required to break one mole of hydrogen-hydrogen bonds in $H_2$ is $436 \ kJ$. What is the longest wavelength of light required to break a single hydrogen-hydrogen bond (in $nm$)?

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