Subtleties of the shell structure of the atom

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

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

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. Adding the extra electron increases shielding without changing the charge of the nucleus. i.eZ* is smaller. The radius of a cation is smaller than its neutral atom Removing the electron decreases shielding without changing the charge of the nucleus. i.e.,Z* is larger. Valence electrons of a cations are in a lower energy shell than in the neutral atom, decreasing the ionic radius.

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 Isolectronicions decrease in size across the period, as Z* increases dramatically. Ex) N3- to F- Na+ to Al3+

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. e- e- e- e- e- 9+ 3+ e- e- e- e- e- e- e- Z* and its effect on size and IE Li F r = 71 pm > r = 152 pm Z*= 1.28 < Z*= 5.13 EA1 = 520 kJ/mol EA1 = 1681 kJ/mol <

8. Periodic distribution of IE1 values List of the IE1 in kJ/mol for the elements IE increases across the period Z* increases IE increases up the group Shielding effect decreases i.e 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 ) Negative since energy is released DHEAincreases in magnitude across period DHEAdecreases in magnitude down the group

10. Electronegativity (c) 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. The Modern Periodic Table name & atomic weight Carbon 12.011

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