Based on the equation : $\; \bigtriangleup E = -2.0 \times 10^{-18} J\; (\large\frac{1}{n_{2}^{2}} - \large\frac{1}{n_{1}^{2}})\;$ the wavelength of the light that must be absorbed to excite hydrogen electron from level n=1 to level n=2 will be : $\;(h=6.625 \times 10^{-34} Js\;,C=3 \times 10^{8}\;ms^{-1})$

Question

$(a)\;1.325 \times 10^{-7}\;m \qquad(b)\;1.325 \times 10^{-10}\;m\qquad(c)\;2.650 \times 10^{-7}\;m\qquad(d)\;5.300 \times 10^{-10}\;m$