ch 8 periodic properties of the elements n.
Skip this Video
Loading SlideShow in 5 Seconds..
Ch. 8: Periodic Properties of the Elements PowerPoint Presentation
Download Presentation
Ch. 8: Periodic Properties of the Elements

Ch. 8: Periodic Properties of the Elements

131 Views Download Presentation
Download Presentation

Ch. 8: Periodic Properties of the Elements

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Ch. 8: Periodic Properties of the Elements Dr. Namphol Sinkaset Chem 200: General Chemistry I

  2. I. Chapter Outline • Introduction • The Hydrogen Atom • Many e- Atoms • Sublevel Energy Splitting • The Aufbau Principle • The Periodic Table • Periodic Trends

  3. I. Organizing Chemical Info • When information of the elements was organized, chemistry began to advance quickly. • Element “triads” and “octaves” • Mendeleev’s periodic table • Appearance of the periodic table is due to electron configurations.

  4. II. The Hydrogen Atom • The Schrödinger equation solves the H atom exactly – gives all possible “manifestations” of the e-. • When an H atom is at its lowest possible energy, the e- exists in its 1s state.

  5. III. Many e- Atoms • The Schrödinger equation can’t solve multi-e- atoms; we only get approximate solutions. • We use quantum #’s from H atom solution to describe orbitals of other atoms.

  6. III. New Considerations • An atom with more than 1 e- is more complicated. • Two more concepts are needed to understand these larger atoms: • Electron spin • Sublevel energy splitting

  7. III. H Atoms in a Magnetic Field

  8. III. e- Spin • e- generate a small magnetic field as if they were spinning. • There are two possible directions e- can spin, so there are two possible states. • spin quantum number (ms) can be either +1/2 or –1/2. • Thus, each e- in an atom can be uniquely identified w/ four quantum #’s.

  9. III. Pauli Exclusion Principle • No two e- in the same atom can have the same 4 quantum #’s!! • H: n=1, l=0, ml=0, ms=1/2 • He has two p+, so it needs two e-: • 1st e-: n=1, l=0, ml=0, ms=1/2 • 2nd e-: n=1, l=0, ml=0, ms=-1/2 • The orbital is filled and the e- have paired spins.

  10. III. H vs. He Energy Levels • One additional e- complicates the He spectrum greater than expected.

  11. III. Reason for More Complicated Spectra • e-/e- repulsions which result in… • Different sublevel energies • Energy of an orbital depends mostly on n (orbital size) and a little on l (orbital shape).

  12. IV. Sublevel Energy Splitting • Three factors contribute to differing sublevel energies: • nuclear charge (Z) • shielding • penetration

  13. IV. Nuclear Charge • p+ in nucleus constantly pull all e-. • Higher charges attract more strongly. • Higher Z lowers orbital E by increasing e-/nucleus attraction.

  14. IV. Effective Nuclear Charge, Zeff • Electrons shield each other from the full charge of the nucleus. • There are two forms of shielding: • e-’s in same orbital • e-’s in different orbitals

  15. IV. Same Orbital Shielding

  16. IV. Inner Orbital Shielding

  17. IV. Penetration • The 3rd e- in Li “occupies” the 2s instead of 2p – WHY? • It has to do w/ orbital shapes and the e- wanting to be close to the nucleus. • To answer this question, we look at radial density curves of orbitals.

  18. IV. 2s and 2p Radial Density Curves • For the same n level, lower l values have lower sublevel energies.

  19. IV. Order of Sublevels

  20. V. The Aufbau Principle • Since e- are “lazy,” they want to “occupy” the lowest energy level possible. • Thus, if we know the energy order of sublevels, then we can “build up” the e- configurations of each atom.

  21. V. Writing e- “in” Orbitals • Two ways to represent how e- are situated in atoms: • e- configuration, nl# • orbital diagram, which uses arrows indicating e-’s and their spin

  22. V. Building Atoms • e.g. Write e- configurations and orbital diagrams for the first 6 elements (H, He, Li, Be, B, C). • Remember Hund’s rule!

  23. V. Some Sublevels Fill Out of Order • Due to shielding effects and penetration, 4s fills before 3d. • Note however that 4s empties before 3d! • Why? As 3d fills, Zeff increases and pulls 3d closer. • Can remember the filling order with the “triangle” diagram.

  24. V. Some Surprises

  25. V. Magnetic Properties • Some metals exhibit magnetism • paramagnetic: atom or ion that has unpaired e-’s • diamagnetic: atom or ion in which all e-’s are paired

  26. VI. The Periodic Table • As you go left to right on the periodic table, you are using the Aufbau principle.

  27. VI. The Periodic Table • Each region of the periodic table indicates what orbitals are being “filled.”

  28. VI. The Periodic Table

  29. VI. Important Parts of the Periodic Table • Each element placed in box w/ atomic #, atomic mass, and atomic symbol. • Atomic # increases as go L to R. • Each horizontal row is period. • Each vertical column is a group or family. • Main group elements are in groups 1,2 and 13-18.

  30. VI. Important Parts of the Periodic Table • Transition elements are in groups 3-12. • Inner-transition elements at the bottom (lanthanides and actinides). • Staircase line separates metals on L from nonmetals on R. Metalloids or semimetals lie adjacent to the line. • Some groups have special names: alkali metals, alkali earth metals, halogens, noble gases.

  31. VI. Valence Electrons • valence electrons: the outermost e- in an atom • Valence e- determine an atom’s chemistry; thus, atoms in the same vertical column have similar chemical properties. • Valence e- can be determined from the Group number.

  32. VII. Trend in Atomic Size • Why?

  33. VII. Trend in Ionization Energy • ionization energy: energy in kJ needed to completely remove 1 mole e- from 1 mole of gaseous atoms/ions • Why? • What about 1st, 2nd, 3rd, ionization energies?

  34. VII. Trend in Electron Affinity • electron affinity: energy change in kJ when 1 mole of e- added to 1 mole of gaseous atoms/ions (generally negative) • Why?

  35. VII. Trend in Ion Size • Why?