1. Chapter 6 Review and Breathe

2. The Wave Nature of Light Electromagnetic radiation is one way energy travels through space. Wavelength is inversely proportional to frequency =c/

3. The Particle Nature of Light Energy is gained or lost in whole number multiples. ΔE=nh This energy is Quantized. Electromagnetic radiation is quantized as photons.

4. Dual Nature of Light Photons have energy of E photon = hc/ Energy is also E=mc 2 (c=speed of light) Thus: m=E/c 2 Substitute E=hc/ in the above equation m=hc/c 2 And finally: m=h/c Light acts as a wave and has mass

5. Do particles have wave characteristics? If a photon has mass m=h/c while it is moving… Then a particle moving at a velocity v has a wavelength using the equation m=h/v Solve for and =h/mv This is de Broglie’s equation.

6. Atomic spectrum of Hydrogen When light is passed through a prism one gets an emission spectrum. When all wavelengths are possible one gets a continuous spectrum. When energy is quantized the spectrum is a line spectrum, or discrete.

7. How do electrons emit light?

8. The Bohr model Bohr proposed that in a hydrogen atom the electron orbits the nucleus in allowed circular orbits. Each orbit has an energy associated with it. E=-2.178x10 -18 J (Z 2 /n 2 )

9. Quantum Mechanical Model Bohr’s model doesn’t work for atoms larger than hydrogen. Electrons are not behaving in a manner that agrees with the circular orbit model.

10. Wave function Electrons in the higher energy levels are acting more like standing waves than like particles. Schrödinger looked at the wave function of the electron. This wave function describes the electron’s orbital.

11. But where is the electron? Heisenberg’s uncertainty principle states that we cannot know with certainty both how fast an electron is moving and where it is. Δx * Δ(mv)> h/4  Probability distribution

12. Quantum numbers each energy level is designated by the value n which is an integer from 1 (lowest energy or "ground state") on up the number of types of orbitals possible on an energy level is also equal to n the maximum number of actual orbitals on an energy level is equal to n 2 the maximum number of electrons in an orbital is equal to 2 the maximum number of electrons on an energy level is equal to 2n 2

13. Quantum numbers Principle quantum number, n =integral values 1, 2, 3 … Represents the energy level. Angular quantum number, l = 0, 1, 2 etc for n -1. It represents the shape of the orbital. Magnetic quantum number m= -l to l. It is related to the orientation of the orbital in relationship to other orbitals in the atom.

14. Pauli Exclusion Principle In given atom, no two electrons may have the same set of quantum numbers. This is known as the Pauli Exclusion Principle. Since each orbital may hold up to two elections. Each election is assigned a separate spin. +½ and -½

15. Summary Table

16. Angular quantum numbers Take note of the nodes! For l= 0, 1, 2, and bottom row 3.

17. Aufbau Principle As protons are added to a nucleus to build up elements, electrons are added too. These electrons are added into the lower energy orbitals first.

18. Hund’s Rule The lowest energy configuration In an orbital is one having the maximum number of unpaired electrons allowed by the Pauli Principle in a set of degenerative orbitals.

19. Electron Configuration and the Periodic Table Elements’ reactivity is based on its valence electrons. The periodic table demonstrates the valence electrons of each group.