1. Subtleties of the shell structure of the atom Physically speaking why does the 4s level in neutral atoms lie below the 3d? The s orbital has a small fraction of its probability density close to the nucleus. 3d orbitals do not have such inner regions, as they only have planar nodes Hence an s electron from a higher shell will sometimes occur at lower energy than a d electron in a lower shell

2. Effective nuclear charge From Li to Ne, nuclear charge increases from 3 to 10 The charge that a 2s or 2p electron feels is different due to the shielding from the electrons in the 1s orbital 2 s orbital penetrate into the 1s orbital and therefore are shielded less on average than p orbitals Note: Shielding effect increases as the number of e’s increases. This is the result of additional shielding from the 2 s and 2 p e’s -1.77 -1.42 -2.78 -3.15 -3.51 -3.87 Note: As Z* increases orbitals shrink towards nucleus as e’s are held more tightly dues to stronger electronic interactions. E

3. Effect on atomic size Consider the size change from F to I Decrease strongly Increase significantly Increase gently Consider the change in size of the atoms from Li to F Consider the size change from Li to Rb Consider the size change from F to I

4. Size of Atoms and Ions Atomic radius decreases along the period, and increases down the group The radius of an anion is larger than its neutral atom. Removing the electron decreases shielding without changing the charge of the nucleus. Valence electrons of a cations are in a lower energy shell than in the neutral atom, decreasing the ionic radius. Adding the extra electron increases shielding without changing the charge of the nucleus. The radius of a cation is smaller than its neutral atom Ie. Z* is smaller. ie., Z* is larger.

5. Sizes of monatomic ions Anions are larger than cations This is always true across a period of the table Ions in each group of the table get larger in size down the group Isolectronic ions decrease in size across the period, as Z* increases dramatically. Ex) N 3- to F - Na + to Al 3+

6. Ionization energy The energy that must be absorbed in order to remove a valence electron from a neutral atom in the gas phase

7. Z* and its effect on size and IE 3+ e-e- e-e- e-e- 9+ e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- LiF r = 152 pm EA 1 = 520 kJ/mol r = 71 pm EA 1 = 1681 kJ/mol Z*= 1.28 Z*= 5.13 > < <

8. Periodic distribution of IE 1 values List of the IE 1 in kJ/mol for the elements IE increases across the period IE increases up the group Z* increases Shielding effect decreases ie. Z* increases

9. Enthalpy of Electronic Attraction Energy released when an element attracts an extra electron into the lowest-energy unoccupied orbital to form an anion For large Z* e’s are held closely to the nucleus therefore e-n interactions will be stronger for an additional electron coming in. Compare Li (Z* = 1.28) with F (Z* = 5.13 )  H EA increases in magnitude across period  H EA decreases in magnitude down the group Negative since energy is released  H EA always negative

10. Electronegativity (  ) i) How strongly does an element hold onto its own electrons ? ii) How strongly is an element able to attract electrons from other elements? A combination of ionization energy and enthalpy of electronic attraction. Which element(s) should have the highest electronegativity? Which element(s) should have the lowest electronegativity?

11. Pauling Electronegativity General trend in element electronegativity

12. The Periodic Table Dmitri Mendeleev, In 1869, noticed that elements exhibited similar behaviour, in groups, in the ratios in which they form molecules with H and O. Elements arranged by increasing mass so that similar elements form columns Incomplete – noble gases are missing Blank spaces left for elements yet to be discovered. In 1913, H.G.J. Moseley- the periodic table is more descriptive if the elements were in order of increasing atomic number rather than increasing mass. Order based on atomic mass causes misalignments.

13. The Modern Periodic Table name & atomic weight 12.011 Carbon

14. Law of Periodicity “The properties of the elements are periodic functions of atomic number.” Metals – Conducting, Ductile Metalloids - Semiconductors Ductile ? Nonmetals – insulators not ductile Group Period Repetition of properties Similar chemical properties