1. Bohr Model of the Atom
Why are the emission spectra of elements not a continuous spectrum?
In 1913, a Danish physicist named Niels Bohr tried to discover the answer, using hydrogen.
Bohr proposed that hydrogen had only certain allowable energy states.

2. Hydrogen’s Energy States
The lowest energy state of an atom= ground state.
When the atom gains energy= excited state.
Bohr said hydrogen, even with only 1 e-, is capable of many different excited states and these are related to the orbit of the e- in the H atom.

3. Hydrogen’s Energy States
Bohr said that the single e- of hydrogen orbited the nucleus in only certain allowed circular orbits.
The smaller the orbit, the lower the energy state (energy level); the larger the orbit, the higher the energy state (energy level)
Bohr assigned a quantum number, n to each orbit. The one closest to the nucleus he numbered “1,” the next “2,” etc.

4. An explanation of Hydrogen’s energy states…
In the ground state, (n= 1) the atom does NOT emit energy. Normal state.
But when excited, the e- jumps up to a higher energy level (n = 2, or 3…)
When this excited e- drops back to the ground state, it emits a photon equal to the difference in energy between the 2 energy levels.

5. Explanation, continued…
Because only certain orbits are possible, only certain frequencies of EM radiation can be emitted.
Think of the energy levels as rungs on a ladder – you can only go up or down on the rungs, NOT in-between.

6. ORIGIN OF THE LINES IN THE HYDROGEN EMISSION SPECTRUM

7. The problem with Bohr…
Bohr explained hydrogen’s spectral lines almost perfectly.
However, it did not explain any other element!
It also did not explain other chemical behavior.
And though his work was groundbreaking, later experiments proved his model fundamentally wrong. Electrons don’t move around the nucleus in circular orbits.
Alas, poor Bohr!

8. Enter Louis de Broglie
In the 1920s, French graduate student Louis deBroglie asked, “if waves (light) can behave like particles, can particles (electrons) behave like waves?”

9. The de Broglie Equation
deBroglie eventually derived an equation to explain the wave properties of particles.
 = _h_ mv
= wavelength of particle
h= Planck’s constant
m=mass of particle
v= velocity

10. The de Broglie Equation…
This equation proves that all moving particles have a particular wave property
Why can’t these waves be seen, like light?
They simply are too small. Even sensitive instruments can’t detect these waves.
When large objects move, when their large mass is divided into the small number of Planck’s constant, the resulting wavelength is very, very tiny.
Electrons, on the other hand, have almost NO mass, so their wavelength is easily measured.

11. Heisenberg Uncertainty Principle
The de Broglie equation ultimately was proven correct.
The next scientist to contribute to our understanding of the atom was German physicist Werner Heisenberg.
Heisenberg said that it’s impossible to make a measurement on an object without disturbing it – at least a little bit.

12. Heisenberg Uncertainty Principle
The smaller the object, the more it is disturbed by measurement.
For an electron, even shining light on it will disturb its frequency and position.
Therefore, Heisenberg concluded, “it is impossible to know precisely both the velocity and position of a particle at the same time.”
Scientists of the time found it hard to accept, but later tests proved it true.
It shows that the uncertainty of any electron’s position is VERY large.

15. Schrodinger’s Wave Equation
In 1926, Austrian scientist Erwin Schrodinger further refined de Broglie’s equation.
He developed a very complex equation that treated the hydrogen electron as a wave. It worked perfectly for not only hydrogen, but ALL elements.
It is the basis for the current model for electrons in atoms, the quantum mechanical model.

16. Basics of Quantum Mechanics
The Schrodinger equation defines the basis of energy levels and possible locations of electrons in atoms.
An electron’s location is simply a probability of where it can be located at a given time.
To make probability predictions easier, quantum mechanics assigns numbers to energy levels and shapes to sublevels.

18. Basics of Quantum Mechanics
Principal Quantum number (n) = relative size & energy of the atomic orbitals (where an electron can orbit the nucleus)
As n increases, the orbital gets larger, the e- spend more time away from the nucleus, & the energy level increases.
Therefore, the Principal Quantum number represents Principal Energy Levels.
The lowest energy level is assigned the number 1. Up to 7 levels are possible for hydrogen.

19. Energy Levels

20. Sublevels
Principal Energy levels contain sublevels (also called orbitals).
Energy level 1 has 1 sublevel, level 2 has 2 sublevels, 3 has 3 and so on.
The sublevels are named for their shapes, as follows: s=spherical; p=dumbbell; d = varying shapes (mostly lobes); and f =varying shapes (like flower petals, almost)
Each orbital can hold up to TWO electrons, maxium.