Electronic Structure and Periodic Trends

1. Electronic Structure and Periodic Trends Chapter 5

3. Experiment 1 • Add an elemental gas to a cathode ray tube and get ----- colors • Hydrogen (H2) purple blue • Neon (Ne) red orange • Helium (He) yellow pink • Argon (Ar) lavender • Xenon (Xe) blue

4. Experiment 2 • Shine white light through a prism -- rainbow • A prism separates light of different wavelength, each color represents a different wavelength.

5. Experiment 3 • Shine the colored light from our gas discharge tubes through a prism  get distinct bands of color (light). • http://jersey.uoregon.edu/vlab/elements/Elements.html

6. Bohr model of the atom • Electrons orbit the nucleus like little planets (planetary model) each with its own energy. Electrons can move from one energy level to another by absorbing or releasing energy. • Energy is released as radiant energy or light.

10. Quantum of energy • the smallest quantity of energy that can be emitted (or absorbed) in the form of electromagnetic radiation.

11. Schrodinger’s quantum mechanical model of the atom •  E = H •  is the wave function or orbital • 2 (probability function) represents the probability of finding an electron at any given position in an atom.

12. s orbitals • ·spherical in shape Electron density map Representation of volume of orbital

14. p orbitals • ·dumbbell shaped • ·three different spatial orientations

15. d orbitals

16. f orbitals Complex shapes 7 different orientations

19. Electronic configuration of the atoms •  Rules for filling orbitals •  1.Lowest energy orbitals are filled first. • 2.Only 2 electrons (of different spin) allowed in each orbital. • 3. When sublevels are filling, fill each orbital with 1 electron of same spin and then pair openly when necessary

26. Phosphorous

30. Valence Electrons • The electrons that occupy the outermost s and p orbitals of an atom

33. Lewis Electron Dot Structures • For elements • Composed of elemental symbol + dots representing the outer shell or valence electrons • For oxygen --

35. Electronic Configuration of the Ions • Cations - Electrons are removed from the highest energy occupied orbital • Anions - Electrons are added to the lowest energy unoccupied orbital • For transition metals -- The highest ns electrons are removed first (even though they are not the last added)

36. Isoelectronic • Isoelectronic species have the same electron configuration. • Atoms tend to gain or lose electrons to become isoelectronic with noble gases • Ne 1s2 2s2 2p6 • Na+1 1s2 2s2 2p6 • F-1 1s2 2s2 2p6

37. Lewis Electron Dot Structures • For ions • Add or subtract dots for electrons gained or lost to form ion. • For O2- – 1s22s22p6

38. Periodic Properties • Metallic Character

40. Periodic Properties • Atomic Size – determined by how far the outermost electrons are from the nucleus

42. Ionic Radii • Cations -- radius decreases due to an increase in Zeffective • Anions -- radius increases due to crowding of more electrons into a shell

45. Periodic Properties • Ionization Energy - The amount of energy required to remove the outermost electron form an isolated neutral atom in the gaseous state.

47. Electron Affinity • The energy change that occurs when an electron is added to an atom (or ion) in the gaseous state. Frequently costs nothing but actually yields energy therefore EA’s are usually negative.

48. Descriptive Chemistry • Alkali Metals • Alkaline Earths • Aluminum • Halogens • Nobel Gases

49. Mendeleev’s Original Periodic Table