Question

1.. Suppose that you shine light of energy 1050 kJ/mol on an H atom in the ground state. What happens to the light and to the electron?

Group of answer choices

a. The light is not absorbed and the final energy of the electron is -1312 kJ/mol.

b.. The light is absorbed and the final energy of the electron is −328 kJ/mol.

c.. The light is not absorbed and the final energy of the electron is −1050 kJ/mol.

d.. The light is absorbed and the final energy of the electron is −1050 kJ/mol.

2.. How many different energies of light are emitted from an H atom with the electron in the n=4 energy shell?

Group of answer choices

a. 3

b. 4

c. 5

d. 6

3..Would a photon of visible light (photons with wavelengths = 400 - 700 nm) have sufficient energy to excite an electron in a hydrogen atom from n = 1 to n = 5?

NOTE: You'll need to calculate how much energy is required for the transition from n = 1 to n = 5. Then convert that energy to wavelength!

Group of answer choices

a. Yes! the electron would need to absorb a photon with a wavelength equal to 93 nm to excite it from n=1 to n=5. Visible light does have sufficient energy.

b. No! the electron would need to absorb a photon with a wavelength equal to 410 nm to excite it from n=1 to n=5. Visible light does not have sufficient energy.

c. Yes! the electron would need to absorb a photon with a wavelength equal to 410 nm to excite it from n=1 to n=5. Visible light has sufficient energy.

d. No! the electron would need to absorb a photon with a wavelength equal to 93 nm to excite it from n=1 to n=5. Visible light does not have sufficient energy.

Answer 1

Question 3.

$\Delta$E = hc/$\lambda$ = hc* R_H * (1/n₁² - 1/n₂²)

i.e. $\Delta$E = 6.625*10^-34 J.s * 3*10⁸ m.s^-1 * 1.097*10⁷ m^-1 * (1/1² - 1/5²)

i.e. $\Delta$E = 2.093*10^-19 J

And 1/$\lambda$ = 1.097*10⁷ m^-1 * (1/1² - 1/5²) = 1.053*10⁷ m^-1, i.e. $\lambda$ = 94.96 nm ~ 93 nm

Therefore, the correct answer is option D.