2-7 Spectroscopy ElectronOrbits.exe

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:

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 = 6.62607004 × 10-34 m2 kg / s

Electromagnetic Spectrum with Visible Light. HIGH ENERGY LOW ENERGY

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).

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:

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

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

www.youtube.com/watch?v=CBrsWPCp_rs Min: 9-13

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