Question

Calculate the wavelength of light required to break the bond between two chlorine atoms in a chlorine molecule. The Cl-Cl bond energy is 243$kJmo{l}^{-1}.(h=6.6\times {10}^{-34}Js;c=3\times {10}^{8}m{s}^{-1}$ , Avogadro's number $=6.023\times {10}^{23}mol{e}^{-1})$.

• A. $4.91\times {10}^{-7}m$
• B. $4.11\times {10}^{-6}m$
• C. $8.81\times {10}^{-31}m$
• D. $6.26\times {10}^{-21}m$

Answer 1

The correct option is A $4.91\times {10}^{-7}m$

The energy required to break one $Cl-Cl$ bond

$=\frac{Bondenergypermol}{Avogadr{o}^{\prime }snumber}$

$=\frac{119.255\times {10}^3}{6.023\times {10}^{23}}J$

Let the wavelength of the photon required to break one Cl-Cl bond be $\lambda$.

$\lambda =\frac{hc}{E}=\frac{6.6\times {10}^{-34}\times 3\times {10}^8\times 6.023\times {10}^{23}}{243\times {10}^3}$

$=\frac{119.255\times {10}^{-34}\times {10}^{31}\times {10}^{-3}}{243}$

$=4.91\times {10}^{-7}m$

So, the correct option is $A$