1. The Quantum Model of the Atom
Section 4.2

2. Bohr’s Problems
Why did hydrogen’s electron exist around the nucleus only in certain allowed orbits?
Why couldn’t the electron exist in a limitless number of orbits with slightly different energies?
Why didn’t the Bohr model work for all atoms?

3. Louis de Broglie French scientist in the 1920s
Observed that the behavior of electrons is similar to the behavior of waves
Suggested that electrons be considered waves
Other experiments confirmed that electrons have wave-like properties

4. Electrons are Like Light
Electrons can be diffracted.
Electrons can show interference patterns

5. Electrons as Particles and Waves
Electrons were determined to also have a dual wave-particle nature
So……..where are they in the atom?
Werner Heisenberg was a German physicist
Tried to detect electrons

6. Heisenberg Uncertainty Principle
Electrons have about the same energy as a photon of light
Try to detect an electron, and the photon of light knocks it off its course
Heisenberg uncertainty principle: it is impossible to determine simultaneously both the position and the velocity of an electron

7. Schrödinger Wave Equation
Erwin Schrödinger was an Austrian physicist
Said that electrons travel in waves
Only waves of specific energies, and therefore frequencies provide solutions to the equation

8. Quantum Theory
The foundation for modern quantum theory was laid by Heisenberg’s uncertainty principle and Schrödinger’s wave equation
Quantum theory: describes mathematically the wave properties of electrons and other very small particles moving very fast

9. Quantum Theory
Electrons do not travel around the nucleus in neat orbits
They exist in certain regions called orbitals
Orbital: a three-dimensional region around the nucleus that indicates the probable location of an electron

10. Quantum Numbers
Quantum numbers: numbers that specify the properties of atomic orbitals and the properties of electrons in orbitals
The first 3 quantum numbers result from solutions to the Schrödinger equation
The fourth quantum number describes a fundamental state of the electron that occupies an orbital

11. Principle Quantum Number
It indicates the main energy level occupied by the electron
Symbolized by n
As n increases, the electron’s energy and its average distance from the nucleus increases
Total number of orbitals in a main energy level is indicated by n2

13. Angular Momentum Quantum Number
It indicates the shape of the orbital
Sublevels: orbitals of different shapes
For a specific main energy level, the number of orbital shapes possible is equal to n
The different shapes are designated s, p, d, and f, each with a specific number of orbitals

14. Basic Shapes

15. More Energy level n=1: only 1 sublevel, s
s orbitals (in the s sublevel) are spherical
One s orbital can hold 2 electrons
There is one s orbital in each energy level
Designated as 1s, 2s, 3s, 4s, etc.
If you have one electron in the 1s orbital, it is designated as 1s1, if you have 2, then 1s2

16. Other Sublevels
The p sublevel has 3 orbitals, each capable of holding 2 electrons for a total of 6 electrons
The d sublevel has 5 orbitals, each capable of holding 2 electrons for a total of 10 electrons
The f sublevel has 7 orbitals for a total of 14 electrons

17. Magnetic Quantum Number
It indicates the orientation of an orbital around the nucleus
S orbitals are spherical, so they only have one possible orientation
P orbitals are “dumbbell” shaped and have 3 possible orientations

18. S and P Orbitals

19. D Orbitals

20. F Orbitals

21. Spin Quantum Number
Indicates the two fundamental spin states of an electron in an orbital
Has only two possible values: +½ and -½
A single orbital can hold a maximum of two electrons, but they must have opposite spin states

23. Table 2 on page 110