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 n 2

12. 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

13. 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 1s 1, if you have 2, then 1s 2

14. 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

15. 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

16. S and P Orbitals

17. D Orbitals

18. F Orbitals

19. Link to Animation http://wolgemuthe.psd401.net/chemistry/03 - elements/studyguides/orbitals.html http://wolgemuthe.psd401.net/chemistry/03 - elements/studyguides/orbitals.html

20. 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

21. Table 2 on page 110