2. Chemistry 101 : Chap. 7 Periodic Properties of the Elements (1) Development of the Periodic Table (2) Effective Nuclear Charge (3) Sizes of Atoms and Ions (4) Ionization Energy (5) Electron Affinity Reading Assignment : 7.6-7.8

3. Periodic Table  Mendeleev ordered known elements according to their weight D. Mendelejeff, Zeitscrift für Chemie 12, 405-406 (1869) Dimitri Mendeleev (1834-1907)  In the modern version, elements are ordered according to atomic number  Elements within a same vertical group have similar properties

4. Periodic Table Alkali metal Alkaline earth metal Halogen rare gas Transition metals

5. Periodic Properties  Properties (e.g. reactivity) of atoms depends on…  Electron configuration  Number of electrons  How tightly electrons are bound to nucleus

6. Periodic Table and Electron Configurations Periodic Table and Electron Configurations

7. Effective Nuclear Charge  Effective nuclear charge (Z eff ) is used to measure how tightly (outer-shell) electrons are bound to nucleus Z eff = # of proton - # of core electrons + + + + + + + - - - - - - - core valence electrons  Valence electrons see nucleus and core electrons as a single unit  Valence electrons experience approximately the total charge of protons and core electrons

8. Effective Nuclear Charge  Example : What is the effective charge of 2p electron of N ?  Example : What is the effective charge of 4s electron of Ca?

9. Effective Nuclear Charge Effective charge, Z eff,, increases with the atomic numbers within a period. 11 Na 12 Mg 13 Al 14 Si …. +11-10= +1+12-10= +2+13-10= +3+14-10= +4 Z eff increases along a period  # of protons increases, but the # of core electrons stay the same along period.  Valence electrons are more strongly bound as atomic number increases within a period

10. Effective Nuclear Charge  Example : Which electron is more tightly bound: 2p electron of Ne or 3s electron of Na? Which electron would be easier to remove?

11. Atomic Radius Atomic radii decrease with the atomic numbers within a period. Increasing Z eff along a period attract valence electrons more strongly, making the atom more compact Lithium Fluorine

12. Atomic Radius Atomic radii increase with the atomic numbers within a group. 11 Na 19 K 37 Rb +11 – 10 = +1 +19 – 18 = +1 +37 – 36 = +1  Z eff essentially remains constant However, more main shells are added and principal quantum number of valence electrons increase

13. Atomic Radius atomic radius increases

14. Ion Radius cations are always smaller than their parent atoms anions are always larger than their parent atoms

15. Ion Radius  Example : Which of the following ions has the largest radius? (1) S, S 2-, O 2- (2) O 2-, Na +, Al 3+

16. Ionization Energy  Ionization Energies (IE) : Energy required to remove electrons from an atom in the gas phase  First Ionization Energy : Energy required to remove one electron from a neutral atom in the gas phase Na (g) Na + (g) + e - IE 1 = 495 kJ/mol  Second Ionization Energy : Energy required to remove second electron from an atom in the gas phase Na + (g) Na 2+ (g) + e - IE 2 = 4562 kJ/mol

17. Ionization Energy First IE increases in general with atomic numbers within a period First IE decreases with atomic numbers within a group Increasing IE

18. Ionization Energy  First IE increases with atomic numbers in a period because..  First IE decrease with atomic numbers in a group because..  Effective nuclear charge increases  decreasing the distance from the nucleus  stronger interaction between valence electron and nucleus  Atomic radius increases with little change in effective nuclear charge  weaker interaction between valence electron and nucleus

19. Ionization Energy  Example : Why there is a huge gap between the 5 th and 6 th ionization energy of nitrogen? Ionization number Ionization E (kJ/mol)

20. Ionization Energy  Example : Why there are irregularities in the first IE within a period? 2s 2 2p 1 2p 4

21. Electron Affinity  Electron Affinity (  E a ) : Energy change that occurs when an electron is added to a neutral atom in the gas phase Cl - (g) Cl (g) + e -  E a = - 349 kJ/mol 100 pm [Ne]3s 2 3p 5 167 pm [Ne]3s 2 3p 6 Energy is released ! More negative  E a means greater attraction between a given atom and an added electron

22. Electron Affinity Added electron goes to new shell Added electron goes to new subshell Added electron leads to noble gas configuration

23. Supplementary Material  Metals : Tend to loose electrons Small ionization energy (IE) Reactivity of metal is related to the IE  smaller IE = more reactive Which one is the most reactive: Na, Mg or K? Reactions (not covered in the class): (a) Metal Oxide with water Na 2 O (s) + H 2 O(l)  2NaOH (aq) (b) Metal Oxide with acid NiO(s) + 2HCl(aq)  NiCl 2 (aq) + H 2 O(l)

24. Supplementary Material  Non-Metals : Tend to gain electrons Usually have negative electron affinity (EA) Seven nonmetals that exist as diatomic molecules H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 Reactions (not covered in the class): (a) Non-Metal Oxide with water P 4 O 10 (s) + 6H 2 O(l)  4H 3 PO 4 (aq) (b) Non-Metal Oxide with base CO 2 (g) + 2NaOH(aq)  Na 2 CO 3 (aq) + H 2 O(l)