1. Chemical Periodicity Chapter 6

2. Chemical Periodicity The periodic table is arranged in rows according to increasing atomic number. Physical and chemical properties of the elements can be explained by electron configurations or positions in the periodic table. Recall that the periodic table is organized in blocks according to subshell being filled.

3. The Periodic Table The periodic table is organized into groups or columns with similar electron configurations. –Na and K –O and S For this reason, the elements in a group will have similar chemical and physical properties. –Above examples and also halogens and inert gases.

4. The Periodic Table Main group elements –‘A’ group elements (periodic table) –Where is the last electron placed? –Representative elements are the first three period of the main group elements H to Ar Noble gases (part of the A group) –Group VIIIA –Where is the last electron? Outermost shell? Demonstration: Noble gas electron configuration

5. The Periodic Table d-transition elements –B group elements (identify on the table) –How many series or periods? –Where does the last electron go? –The number of the d subshell being filled is one less than the current shell number. Why?

6. The Periodic Table f-transition elements –Termed as the inner-transition elements. –Number of series or periods? –Where is the last electron placed? Why is the f-transition elements termed as inner transition elements? Note: Electron buildup for f-transition elements can be complicated (Appendix B)

7. Properties of the Elements The outermost electrons or the valence electrons are largely responsible for the chemical and physical properties of the elements. The trends in the physical properties that will be discussed can be explained in a large degree to the valence electrons.

8. Properties of the Elements The forces experienced by the valence electrons can determine trends observed in many atomic properties. What type of forces can the electrons experience in an atom? –Charge from the nucleus. A strong attraction between the electron and the nucleus lowers the electron energy. The electron is more tightly bound to the atom. –Interactions among the electrons. There are repulsions between electrons in the same orbital. Energy?

9. Factors that Influence the Energy of the Electron 1.Shielding from the nucleus – the valence electrons are shielded by inner electrons. –Z eff = Z -  (effective nuclear charge) Z = nuclear charge,  = amount of shielding Li example –Electrons with lower l quantum numbers shield those with higher l quantum numbers. Why? –Electrons in the same l do not shield each other significantly. Z eff increases from left to right.

10. Trend in Z eff within a Period/Row Z eff increases from left to right within a period. What consequence will this have on the valence electrons? Energy of the electron?

11. Factors that Influence the Energy of the Electron 2.Distance from the nucleus –Attraction between opposite charges How will the decreasing d affect the energy of the electron? How will the energy of the electron be affected by shell and subshell numbers?

12. Combining These Two Affects Discuss the dependencies: Decreasing the value of n? Increasing Z eff ?

13. Factors that Influence the Energy of the Electron 3.Interactions among electrons –Placing electrons in the same orbital creates repulsions. What affect will this have on the electron? The chemical periodicity of the properties can be discussed based on the preceding factors.

14. Atomic Radii The size of an atom is roughly determined by the ______________. –Typical sizes range from 0.3 to 3 Å. The closer the outermost electron to the nucleus the ______ the energy. –Qualitative relationship

15. Atomic Radii What does the relationship predict going down a group/column? What does the relationship predict going across a period? Look at chart on page 241. STM image of atoms STM operation

16. Ionization Energy Minimum amount of energy required to remove the most __________ electron from a gaseous atom. –1 st ionization energy (IE 1 ) Mg(g) +738 kJ/mol  Mg + (g) + e - –2 nd ionization energy (IE 2 ) Mg + (g) + 1451 kJ/mol  Mg 2+ (g) + e - Why is IE 2 >IE 1 ?

17. Ionization Energy How is energy of the electron related to the ionization energy? Examine Figure 6-1 and Table 6-1. What are the trends? Why?

18. Ionization Energy Ionization energy down a column? Ionization energy across a period? –Na through Ar –Explain the trend This is a periodicity pattern. Holds well for Group A.

19. Ionization Energies – Electron Configurations Trends for IE’s in the group of the representative elements can partially be explained by electron configurations. –Why? –Group IA, the alkali metals, has the lowest IE’s. 1 st and 2 nd ionization energies for Na –Group IIA trends 1 st, 2 nd, and 3 rd ionization energies Why is there a large relative increase in the 3 rd ionization energy?

20. Ionization Energies - Reactivity Elements with low or high ionization energies form _______________. –Which groups in A? Elements with intermediate values of IE ___________ upon reacting. –Which groups in A?

21. Electronegativity A measure of the relative tendency of an atom to _________ upon combining chemically with another atom. –Elements with high electronegativities? –Elements with low electronegativities? –Elements with intermediate electronegativities?

22. Electronegativity Representative elements –Increases from left to right Why? –Decreases going down a group Why? Type of bonding that will occur in a compound

23. Electron Affinity The amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge. –Measure of an atom’s ability to form a negative ion. Cl(g) + e -  Cl - (g) + 349 kJ (releases energy) Mg(g) + e - + 231 kJ  Mg - (g) (takes energy)