1. Periodic Properties of the Elements Chapter 7

2. Effective Nuclear Charge Orbitals of the same energy are said to be degenerate. Effective nuclear charge is the charge experienced by an electron on a many- electron atom. The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons.

3. Effective Nuclear Charge Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nucleus. The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. As the average number of screening electrons (S) increases, the effective nuclear charge (Z eff ) decreases. As the distance from the nucleus increases, S increases and Z eff decreases.

4. Energies of Orbitals The result of the Effective nuclear charge on the electronic configuration is a shift of the orbital ordering for large (n = 4 or more) electronic systems.

5. Energies of Orbitals

6. Sizes of Atoms Electron Shells in Atoms Consider a simple diatomic molecule. The distance between the two nuclei is called the bond distance. If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.

7. Atomic Radii As a consequence of the ordering in the periodic table, properties of elements vary periodically. Atomic size varies consistently through the periodic table. As we move down a group, the atoms become larger. As we move across a period, atoms become smaller. There are two factors at work: principal quantum number, n, and the effective nuclear charge, Z eff.

8. Atomic Radii decreasing across a period is due to: shielding or screening effect inner electrons [He] or [Ne], etc. block the nuclear charge for 2 or 10 or __ electrons consequently the outer electrons feel a stronger effective nuclear charge Li [He] shields effective charge is +1 Be [He] shields effective charge is +2

9. Atomic Radii

10. All radii are in angstroms.

11. Ionic Radii Cations (+ ions) are always smaller than their neutral atoms

12. Ionic Radii Anions (- ions) are always larger than their neutral atoms

13. Ionization Energy minimum amount of energy required to remove the most loosely held electron from an isolated gaseous atom measure of an element’s ability to form positive ions Mg (g) + 738kJ/mol  Mg + + e - Atom (g) + energy  ion + (g) + e - first ionization energy

14. Ionization Energy second ionization energy second ionization energy energy required to remove a 2nd electron ion + + energy  ion 2+ + e - Mg + + 1451 kJ/mol  Mg 2+ + e - can have 3rd, 4th, etc. ionization energies

15. Ionization Energy generally increases as you go across a period important exceptions at Be & Mg, N & P generally decreases as you go down a group

17. Ionization Energy First 4 ionization Energies (kJ/mol) - Period 3

18. Ionization Energy these energies are exactly why these ions form Na becomes Na + Mg becomes Mg 2+ Al becomes Al 3+ Si does not form simple ions

19. Ionization Energy

20. Electronegativity The attraction of an atom to electrons This is not a measurable property BUT it is very useful in helping to predict bonding (attraction of electrons)

21. Electronegativity