Based on the equation: $$\Delta\mathrm{E} = -2.0 \times 10^{-18} \mathrm{~J} \left(\frac{1}{\mathrm{n}_2^2} - \frac{1}{\mathrm{n}_1^2}\right)$$ the wavelength of the light that must be absorbed to excite hydrogen electron from level $n = 1$ to level $\mathrm{n} = 2$ will be: ($\mathrm{h} = 6.625 \times 10^{-34} \mathrm{Js}$, $\mathrm{C} = 3 \times 10^8 \mathrm{~ms}^{-1}$) (1) $1.325 \times 10^{-7} \mathrm{~m}$ (2) $1.325 \times 10^{-10} \mathrm{~m}$ (3) $2.650 \times 10^{-7} \mathrm{~m}$ (4) $5.300 \times 10^{-10} \mathrm{~m}$
Based on the equation:
$$\Delta\mathrm{E} = -2.0 \times 10^{-18} \mathrm{~J} \left(\frac{1}{\mathrm{n}_2^2} - \frac{1}{\mathrm{n}_1^2}\right)$$
the wavelength of the light that must be absorbed to excite hydrogen electron from level $n = 1$ to level $\mathrm{n} = 2$ will be: ($\mathrm{h} = 6.625 \times 10^{-34} \mathrm{Js}$, $\mathrm{C} = 3 \times 10^8 \mathrm{~ms}^{-1}$)\\
(1) $1.325 \times 10^{-7} \mathrm{~m}$\\
(2) $1.325 \times 10^{-10} \mathrm{~m}$\\
(3) $2.650 \times 10^{-7} \mathrm{~m}$\\
(4) $5.300 \times 10^{-10} \mathrm{~m}$