1. Bohr Model of the Atom & Light
Monday, October 31st, 2016

2. Dalton’s Atomic Model

3. Plum Pudding Model (Thomson)

5. The Development of a New Atomic Model
Rutherford’s model was incomplete:
How were electrons distributed?
What prevents negative electrons from crashing into positive nucleus?
1900s-New Atomic Model (revolutionary)
Relationship between light & electrons

6. Niels Bohr (Born in Denmark 1885-1962)
Student of Rutherford

7. Niels Bohr’s Model (1913)
Electrons orbit the nucleus in circular paths of fixed energy (energy levels).

8. How many Electrons Can Fit in Each Energy Level?
Energy level is also called “principal quantum number”
Energy Level = n
n max. # electrons
1 2
2 8
3 18
4 32
5 50
6 72
Max # e = 2n2
Big Question: Why?

9. Bohr Model of an Atom Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away from the nucleus
An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)

10. Sketch the Bohr Model for the following elements – which is unreactive?
Nitrogen
Argon

11. Sketch the Bohr Model for the following elements – which is unreactive?
Nitrogen
Argon

12. } Further away from the nucleus means more energy. Increasing energy
Fifth
Further away from the nucleus means more energy.
There is no “in between” energy
Energy Levels
Fourth
Increasing energy
Third
Second
First

14. The closer an electron is to a proton, the more stable the atom!
The most stable location for an electron is as close to the nucleus as it can get… that is its ground state configuration (the lowest possible energy level the electron can be at)
But electrons can jump from energy level to energy level!
An atom is said to be in an excited state when its electrons are at an energy level higher than the ground state.

15. Electrons absorb or emit energy when they jump from one energy level to another.
A quantum of energy is the amount of energy required to move an electron from one energy level to another.

16. How do we determine (calculate) energy?
A scientist (Max Planck) determined hot objects
emit energy in packets called “quanta”
“quantum” - minimum energy that can be gained or lost by an atom
E = h
Where E = energy, h = Planck’s constant (6.626 x Js), and  = frequency (in s-1)

17. Photons are bundles of light energy that is emitted by electrons as they go from higher energy levels to lower levels.

18. High Freq Electromagnetic Spectrum Low Freq
High E Low E
Light emitted produces a unique emission spectrum.

19. Wavelength & Frequency
Speed of light (c) = 3.00 x 108 m/s
Wavelength ( ) = distance between waves
Frequency () = # waves that pass point in given time
c =  
as  increases,  decreases
as  increases,  decreases
(slinky analogy)

21. Emission Spectra White light - continuous spectrum Hydrogen atoms
- line-emission spectrum
Big Question
- Why did hydrogen atoms only
give off specific frequency
(colors) of light?

22. The “Fingerprints” of Atoms

23. The atom is quantized, i.e. only certain energies are allowed.
Atomic Spectra
The atom is quantized, i.e. only certain energies are allowed.

24. Neils Bohr Solved the Mystery
Electron circled nucleus in “orbit” of fixed energy
Absorption-electrons can “hop” from ground state to excited state
Emission-when electrons “fall” from excited state to ground state
Energy difference
corresponds to hydrogen’s
spectral lines