91Ó°ÊÓ

A photon’s wavelength connects energy through Planck’s law.

$E\_photon=hv=\frac{hc}{\mathrm{λ}}$

Planck’s constant h is$6.626.{10}^{-34}J{S}^{-1}$ and Speed of light c is role="math" localid="1663779143148" $3.{10}^{8}m{s}^{-1}$ .

Avogadro's constant $N_A$multiplied by the energy of one photon equals the energy of one mole of photons.

E$=E_{photon}.N_A=9.04.{{10}^{-19}}^{}J.6.022.{10}^{23}mo{l}^{-1}=544115Jmo{l}^{-1}=544KJmo{l}^{-1}$

Standard reduction potentials ( $∆E^{°1}$) obtained from the table 14-4 are,

$$\begin{gathered} E_{H20=-0.815V} \\ V{E}_{NAD{P}^{+}=-0.320V} \\ ∆E°=E_{NADP+}+E_{H20} \end{gathered}$$

Free energy change is:

$∆G^{°}=-nF∆E^{°}$

where n is number of electrons and F is Faraday constant, $96.485J{V}^{-1}mo{l}^{-1}$.

$∆G^{°}=-4.96.485J{V}^{-1}mo{l}^{-1}.-1.135V=438041Jmol=438KJmo{l}^{-1}$

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The number of moles required is calculated as:

$\frac{438KJmo{l}^{-1}}{544KJmo{l}^{-1}}=0.8$

Hence, 0.8 moles of photons are needed for production of 1 mole of oxygen