Topic 4

1. Topic 4 • Periodicity –  the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of the periodic table of elements.

2. The “Shell Model” • The Shell Model (Shells = energy levels) • The model is based on coulomb’s law and qualitatively predicts ionization energies, which can be measured in the lab. • Understanding how the shell model is consistent with the experimental data is key

3. Coulomb’s Law • Typically used to determine the energy of interaction (attractive or repulsive) between two ions • E = energy in joules • r = distance between the ion’s centers • Q = the numerical ion charges

4. Can be applied Qualitatively to: • electron (-) / proton (+) attraction in an atom • electron (-) / electron (-) repulsion in an atom • Since the charge of an electron and proton are opposite, but equal the Energy is determined by the distance ( r) between them. • The smaller the distance the greater the attraction or repulsion and the larger the distance the weaker the attraction or repulsion.

5. Polyelectronic atoms • Can’t solve Schrodinger’s equation exactly. • It becomes complicated because of the repulsion of other electrons. • The best solution is to treat each electron as if it were effected by the net field of charge from the attraction of the nucleus and the average repulsions of all other electrons. • This is called effective nuclear charge. It is the total nuclear charge (from protons) minus and effect of the electron repulsions. • Link to an Effective Nuclear Charge video • http://www.youtube.com/watch?v=IvSmfgxCSNQ

6. Ionization Energy • Ionization energy the energy required to remove an electron from an atom in the gaseous state. • X(g) → X+(g) + e- • Highest energy (outermost) electron removed first. • First ionization energy (I1) is that required to remove the first electron. • Second ionization energy (I2) - the second electron

7. Trends in ionization energy • for Mg • I1 = 735 kJ/mole • I2 = 1445 kJ/mole • I3 = 7730 kJ/mole (There is a large jump in energy to remove the third electron) • The effective nuclear charge increases as you remove electrons due to an increasing nuclear charge per electron • It takes much more energy to remove a core electron than a valence electron because there is less shielding and greater stability.

8. Explain this trend • For Al • I1 = 580 kJ/mole • I2 = 1815 kJ/mole • I3 = 2740 kJ/mole • I4 = 11,600 kJ/mole

9. Across a Period • Generally from left to right, I1 increases because there is a greater effective nuclear charge (Zeff) with the same shielding. • As you go down a group I1 decreases because electrons are farther away. • ( Following Coulombs Law we can see that as the distance between the nucleus and the electrons increases the force of attraction decreases. )

10. It is not always straight forward • Zeff changes as you go across a period, so will I1 • Half filled and filled orbitals are harder to remove electrons from. • here’s what it looks like.

11. First Ionization energy Atomic number

12. Exceptions: • Beryllium Boron Decreases due to shielding from the 2s orbital • Nitrogen Oxygen Decreases due to extra electron repulsions in the doubly occupied oxygen 2p orbital

13. Electron Affinity • Is the energy change associated with the addition of an electron to a gaseous atom • X(g) + e- → X- (g) • EA generally increases left to right across the periodic table. • Explain in terms of Zeff • EA generally decreases going down the periodic table. • Explain in terms of Coulomb’s Law

14. Atomic Radius • Is half the distance between the nuclei in a molecule consisting of identical atoms. • AR decreases going left to right across the periodic table • Explain in terms of Zeff • AR increases going down the periodic table • Explain in terms of Coulomb’s Law

15. Ionic Radii(Sizes of Ions) • Size plays an important role in determining the structure and stability of ionic solids • Relative size of an ion and its “parent” atom • Positive ions are created by removal of an electron and the resulting cation is smaller. • Reasons: • An energy level is lost • Greater Zeff on the remaining electrons

16. Negative ions are created by an addition of an electron(s) resulting in an anion that is larger. • The additional electron(s) create an increase in electron repulsion. • Relative sizes of Isoelectronic Ions • The number of electrons are equal, so the number of protons determine what happens • An increase in protons from left to right increases the Zeff and decreases the size of the atom. • Na+, Mg2+, Al3+ (all have 10 electrons)

17. Typical Ionic Charges • To determine the most likely ionic charge for group A elements: • For metals – determine the number of electrons to remove to attain a noble gas configuration • For nonmetals – determine the number of electrons to be gained to attain a noble gas configuration.

18. Info in the Periodic Table • Know the groups: Alkali Metals, Alkaline Earth Metals, Halogens, and Noble Gases. • It is the number and type of valence electrons that determine an atoms chemistry. • You can get the electron configuration from it. • Metals lose electrons have the lowest IE • Non metals- gain electrons most negative (greatest) electron affinities.