1. Chapter 6 The Periodic Table & Periodic Law

2. Section 6.1 Development of the Modern Periodic Table

4. John Newlands In 1864, noticed when the elements were arranged in order of increasing atomic mass, their properties repeated every eight elements. –THE LAW OF OCTAVES

5. Meyer & Mendeleev In 1869, published almost identical versions with the elements in order of increasing atomic mass and in columns with similar properties.

6. Mendeleev Mendeleev is given more credit than Meyer BECAUSE: –He published his table first –He better demonstrated his table Suggested some of the previously measured masses were incorrect Left blanks for not yet discovered elements

7. Mendeleev’s Predicted Properties of Ge “eka Silicon” and Its Actual Properties Table 8.1 Property Predicted Properties eka Silicon Actual Properties Ge atomic mass appearance density molar volume specific heat capacity oxide formula oxide density sulfide formula and solubility chloride formula (boiling point) chloride density 72amu gray metal 5.5g/cm 3 13cm 3 /mol 0.31J/g*K EO 2 4.7g/cm 3 ES 2 ; insoluble in H 2 O; soluble in aqueous (NH 4 ) 2 S ECl 4 ; (<100 0 C) 1.9g/cm 3 72.61amu gray metal 5.32g/cm 3 13.65cm 3 /mol 0.32J/g*K GeO 2 4.23g/cm 3 GeS 2 ;insoluble H 2 O; soluble aqu (NH 4 ) 2 S GeCl 4 ; (84 0 C) 1.844g/cm 3

8. Was Mendeleev psychic???? periodic law: when arranged by atomic # elements exhibit a periodic recurrence of similar properties –Quantum-mechanical model of atom explains organization of table Development of Periodic Table

9. Mosley In 1913, using X-rays, he discovered a unique number of protons in the nuclei of atoms for each element. Today the elements are arranged in order of increasing atomic number PERIODIC LAW –There is a periodic repetition of chemical and physical properties of the elements when they are arranged in order of increasing atomic number

10. Arrangement of the Periodic Table Groups/Families –18 vertical columns (↑↓) –Two Labeling Systems 1.Number-and-letter system -A through 8A columns (representative elements) -1B through 8B short columns (transition elements) 2. Number system - 1-18 Periods –7 horizontal rows (↔)

12. PERIODS GROUPS/FAMILIES

13. Arrangements of the Periodic Table

14. Metals Shiny Good conductors of heat and electricity Malleable & Ductile Generally Solid at room temperature Group 1 Alkali Metals Group 2 Alkaline Earth Metals Groups 3-12 Transition Metals Lanthanide & Actinide Groups Inner Transition Metals

16. Nonmetals & Metalloids Nonmetals –Dull –Generally gases or brittle solids at room temperature –Poor conductors of heat and electricity Metalloids –Elements with physical and chemical properties of both metals and nonmetals –Rest on the “stair-step” B Si As Te At Ge Sb Po ←Metals Nonmetals →

17. Section 6.2 Classification of Elements

18. Element Placement Why are elements put into groups/families together? Because they have similar chemical properties Why do elements have similar chemical properties? Because they have the same number of valence electrons Group 1 – Alkali Metals Period 2Lithium 1s 2 2s 1 [He]2s 1 Period 3Sodium 1s 2 2s 2 2p 6 3s 1 [Ne]3s 1 Period 4Potassium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ne]4s 1 ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON

19. Recurring pattern in e - configuration is basis for periodic behavior. Main group, group # = valence e - count –Valence e - responsible for chemistry Elements in same group behave similarly

20. Dot Diagrams for Representative Elements

21. Figure 8.12 A periodic table of partial ground-state electron configurations

22. Representative Elements s-block elements –Groups 1&2, hydrogen & helium –Valence electrons occupy outermost s sublevels only p-block elements –Groups 13-18 (except helium) –Valence electrons include a full outermost s sublevel and a filled or partially filled p sublevel Period number is equal to the principle energy level where the valence electrons are located

23. Transition Elements d-block elements –Groups 3-12 –Valence electrons include a full outermost s sublevel and a filled or partially filled d sublevel The period number minus 1 equals the principle energy level where the valence electrons are located

24. Inner transition metals f-block elements –Lanthanide & Actinide Groups –Full or partially full outermost s sublevel, and full or partially full outermost f sublevel The period number minus 2 equals the principle energy level where the valence electrons are located.

26. Section 6.3 Periodic Trends

27. Atomic Radius Half the distance between two nuclei of identical atoms that are chemically bonded together Down the group –atomic radius increases Across the period –atomic radius decreases

28. Atomic Radius Decreases Atomic Radius Increases

29. Figure 8.16 Periodicity of atomic radius

30. Practice Atomic Radius Which has the larger atomic radii of the following? B or AlNa or Mg F or Cl Which has the smaller atomic radii of the following? H or HeK or Cs N or Ne Circle the one with the largest atomic radius and underline the one with the smallest. C, Si, GeV, Cr, WN, Mg, Ca

31. Ionization Energy The amount of energy required to remove an electron from the atom (how tightly an atom holds on to its electrons) Down a group –ionization energy decreases Across a period –ionization energy increases

32. Ionization Energy (IE) –Energy required for complete removal of 1 mole of e - from 1 mole of atoms Atoms w/ low IE form cations (lose e - ) Atoms w/ high IE form anions (gain e - ) Trends in Atomic Properties Na(g)  Na + (g) + e - I 1 Na + (g)  Na 2+ (g) + e - I 2 I 1 < I 2 < I 3

33. Figure 8.18 8.4 Trends in Atomic Properties Greater IE, more difficult to remove e - –Positive values, energy into atom Larger atoms easier to ionize

34. Ionization Energy Increases Ionization Energy Decreases

35. Figure 8.17 Periodicity of first ionization energy (IE 1 )

36. Practice Ionization Energy Which has the greater ionization energy? Ne or ArN or O Sc or Ti Which has the smaller ionization energy? Al, Si, PK, Rb, SrBe, Mg, Ca

37. Ionic Radius Octet Rule –Atoms tend to gain, lose, or share electrons in order to achieve a full outer energy level (typically 8 are needed) Ion –An atom that has an overall charge due to the gaining or losing of electrons

38. Figure 8.25 Main-group ions and noble gas configurations

39. Ionic Radius Comparisons Metals have LOW ionization and electron affinity –They lose electrons to form positively charged ions –Positive charged ions are smaller than the original atom Nonmetals have HIGH ionization energy and electron affinity –They gain electrons to form negatively charged ions –Negatively charged ions are larger than the original atom

40. Figure 8.29 Cation smaller than parent –e - removed, other e - feel greater Z eff Anion larger than parent –e - added, e - /e - repulsions occupy more space Trends in Properties of Monatomic ions

41. Ionic Radius Increases FOR IONIC RADIUS… MUST FOLLOW METAL/NON-METAL RULES

42. Ionic Radius Practice Which is the smaller of the two? Lithium ion or Lithium atom Chlorine ion or Chlorine atom Underline the following that will form a positively charged ion and circle the ones that will form a negatively charged ion. MgFAlCu BrNSK How will the radius of each of the above change when an ion is formed? MgFAlCu BrNSK

43. Electronegativity The ability of an atom to attract electrons in a chemical bond. Down the group –Electronegativity values decrease Across the period –Electronegativity values increase *Noble gases are the exception to this rule.

44. Electronegativity Increases Electronegativity Decreases

45. Electronegativity Practice Which has the greater electronegativity value? B or NSi or SnCr or W Which has the smaller electronegativity value? Rb, Sr, YGa, In, SnAs, Se, S

46. CUMULATIVE REVIEW Which has the smallest atomic radius between Ga, In, & Tl? –Which has the highes ionization energy? Which is the smallest: an atom of sodium, an ion of sodium, or an atom of potassium? Which has the greatest electron affinity between zinc, arsenic, or bromine? Which has the lowest ionization energy?

47. Figure 8.21 8.4 Trends in Atomic Properties

48. Electronegativity Increases Atomic Radius Decreases Ionic Radius Increases Electron Affinity Increases Ionization Energy Increases Electronegativity Decreases Atomic Radius Increases Ionic Radius Increases Electron Affinity Decreases Ionization Energy Decreases