The Quantum Model of the Atom

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