1. 2-7 Spectroscopy
ElectronOrbits.exe

2. The internal structure of the atom was deduced through a series of experiments during the early 1900s. Several of these experiments involved shining light into the atom to observe the effects. To understand the results, a brief summary of light is presented below:

3. Wavelength (λ) = distance between 2 similar points on wave.
Frequency (ν) = # waves/sec (Hz)
Relationships:
λ • ν = c
where c = speed of light =
3.0 x 108 m/s
E = h • ν where h = constant
 h = × m2 kg / s

4. Electromagnetic Spectrum with Visible Light.
HIGH
ENERGY
LOW
ENERGY

5. Bohr used the following emission spectrum to develop his idea of quantized energy levels and quantum leaps:
Note how the emission spectrum contains 4 separate lines. This “discrete line spectrum” is in contrast to the “continuous spectrum of the visible spectrum (the rainbow: ROYGBIV).

6. Interpretation of the spectral experiment: energy added to the atom is absorbed by the electron, which makes a quantum leap from a lower energy state to a higher energy state. The electron doesn’t remain excited very long, and returns to a lower energy state, giving off a color of light whose color exactly matches the energy level difference:

7. E e- Light E e- n = 4 ____________________________
n = 3 ____________________________
n = 2 ____________________________
n = 1 ____________________________ E e-
Light E e-

8. The energy levels are represented by an n, with n=1 being the “ground state”, n=2 is the first “excited state”, etc.
For an electron configuration the ground state represents all the electrons in their lowest energy level configurations. If any of the electrons are in a higher energy state, then the atom’s electrons are in an excited state

9.
Min: 9-13

10. Spectroscope lab Ground state of Be: 1s22s2
Excited state of Be: 1s22s12p1
Demos
Ca, Cu(II), Li, K, Na
peach green Red Lavender yellow
Spectroscope lab