1. Quantum Physics
Atomic spectra and atomic energy states

2. How do you excite an atom?
Heating to a high temperature
Bombarding with electrons
Having photons fall on the atom

3. The opposite of excitement
A. Lee in “Action”

4. Atomic spectra When have you seen it? How does it work?
When a gas is heated to a high temperature, or if an electric current is passed through the gas, it begins to glow.

5. Atomic spectra
When a gas is heated to a high temperature, or if an electric current is passed through the gas, it begins to glow.

6. Types of spectra
What are the 3 types of atomic spectra?

7. Emission spectrum
If we look at the light emitted (using a spectroscope) we see a series of sharp lines of different colours. This is called an emission spectrum.

8. Absorption Spectrum
Similarly, if light is shone through a cold gas, there are sharp dark lines in exactly the same place the bright lines appeared in the emission spectrum.
Light source
gas
Some wavelengths missing!

9. How do we use them?
Scientists had known about these lines since the 19th century, and they had been used to identify elements (including helium in the sun), but scientists could not explain them.

10. Rutherford
At the start of the 20th century, Rutherford viewed the atom much like a solar system, with electrons orbiting the nucleus.
What’s the problem with this model?
However, under classical physics, the accelerating electrons (centripetal acceleration) should constantly have been losing energy by radiation (this obviously doesn’t happen).

11. Niels Bohr
In 1913, a Danish physicist called Niels Bohr realised that the secret of atomic structure lay in its discreteness, that energy could only be absorbed or emitted at certain values.

12. The Bohr Model
We say that the energy of the electron (and thus the atom) can exist in a number of states n=1, n=2, n=3 etc. (Similar to the “shells” or electron orbitals that chemists talk about!)
n = 1
n = 2
n = 3

13. The Bohr Model
The energy level diagram of the hydrogen atom according to the Bohr model
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
High energy n levels are very close to each other
Energy eV
Electron can’t have less energy than this
-13.6

14. The Bohr Model
An electron in a higher state than the ground state is called an excited electron. It can lose energy and end up in a lower state.
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
High energy n levels are very close to each other
Wheeee!

15. Atomic transitions
If a hydrogen atom is in an excited state, it can make a transition to a lower state. Thus an atom in state n = 2 can go to n = 1 (an electron jumps from orbit n = 2 to n = 1)
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
Wheeee!
electron

16. Atomic transitions
Every time an atom (electron in the atom) makes a transition, a single photon of light is emitted.
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
electron

17. Atomic transitions
The energy of the photon is equal to the difference in energy (ΔE) between the two states. It is equal to hf. ΔE = hf
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
electron
ΔE = hf

18. The Lyman Series
Transitions down to the n = 1 state give a series of spectral lines in the UV region called the Lyman series.
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
Lyman series of spectral lines (UV)

19. The Balmer Series
Transitions down to the n = 2 state give a series of spectral lines in the visible region called the Balmer series.
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
Balmer series of spectral lines (visible) UV

20. The Pashen Series
Transitions down to the n = 3 state give a series of spectral lines in the infra-red region called the Pashen series.
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
Pashen series (IR)
visible UV

21. Emission Spectrum of Hydrogen
The emission and absorption spectrum of hydrogen is thus predicted to contain a line spectrum at very specific wavelengths, a fact verified by experiment.
Which is the emission spectrum and which is the absorption spectrum?

22. Pattern of lines
Since the higher states are closer to one another, the wavelengths of the photons emitted tend to be close too. There is a “crowding” of wavelengths at the low wavelength part of the spectrum
n = 1 (the ground state)
n = 2
n = 3
n = 4
n = 5
-13.6
Energy eV
Spectrum produced

23. Limitations of the Bohr Model
Can only treat atoms or ions with one electron
Does not predict the intensities of the spectral lines
Inconsistent with the uncertainty principle (see later!)
Does not predict the observed splitting of the spectral lines