1. IIIIII The Periodic Table & Periodic Law I. Development of the Modern Periodic Table

2. A. Mendeleev zDmitri Mendeleev (1869, Russian) yOrganized elements by increasing atomic mass yElements with similar properties were grouped together yThere were some discrepancies

3. A. Mendeleev zDmitri Mendeleev (1869, Russian) yPredicted properties of undiscovered elements

4. B. Moseley zHenry Moseley (1913, British) yOrganized elements by increasing atomic number yResolved discrepancies in Mendeleev’s arrangement yThis is the way the periodic table is arranged today!

5. C. Modern Periodic Table zGroup (Family) zPeriod

6. 1. Groups/Families zVertical columns of periodic table zNumbered 1 to 18 from left to right zEach group contains elements with similar chemical properties

7. 2. Periods zHorizontal rows of periodic table zPeriods are numbered top to bottom from 1 to 7 zElements in same period have similarities in energy levels, but not properties

8. zMain Group Elements zTransition Metals zInner Transition Metals 3. Blocks

9. Lanthanides - part of period 6 Actinides - part of period 7 Overall Configuration

10. IIIIII II. Classification of the Elements (pages 182-186) Ch. 6 - The Periodic Table

11. A. Metallic Character zMetals zNonmetals zMetalloids

12. 1. Metals zGood conductors of heat and electricity zFound in Groups 1 & 2, middle of table in 3-12 and some on right side of table zHave luster, are ductile and malleable

13. a. Alkali Metals zGroup 1 z1 Valence electron zVery reactive zElectron configuration yns 1 zForm 1 + ions zCations yExamples: Li, Na, K

14. b. Alkaline Earth Metals zGroup 2 zReactive (not as reactive as alkali metals) zElectron Configuration yns 2 zForm 2 + ions zCations yExamples: Be, Mg, Ca, etc

15. c. Transition Metals zGroups 3 - 12 zReactive (not as reactive as Groups 1 or 2), can be free elements zElectron Configuration yns 2 (n-1)d x where x is column in d-block zForm variable valence state ions zCations yExamples: Co, Fe, Pt, etc

16. 2. Nonmetals zNot good conductors zFound on right side of periodic table – AND hydrogen zUsually brittle solids or gases

17. a. Halogens zGroup 17 (7A) zVery reactive zElectron configuration yns 2 np 5 zForm 1 - ions – 1 electron short of noble gas configuration zAnions yExamples: F, Cl, Br, etc

18. b. Noble Gases zGroup 18 zUnreactive, inert, “noble”, stable zElectron configuration yns 2 np 6 full energy level zHave a 0 charge, no ions zExamples: He, Ne, Ar, Kr, etc zElements form compounds to have electron configurations like noble gases

19. 3. Metalloids zSometimes called semiconductors zForm the “stairstep” between metals and nonmetals zHave properties of both metals and nonmetals zExamples: B, Si, Sb, Te, As, Ge, Po, At

20. B. Chemical Reactivity zAlkali Metals zAlkaline Earth Metals zTransition Metals zHalogens zNoble Gases

21. C. Valence Electrons zValence Electrons ye - in the outermost energy level zGroup #A = # of valence e - (except He) 1A 2A 3A 4A 5A 6A 7A 8A

22. C. Valence Electrons zValence electrons = yelectrons in outermost energy level zYou can use the Periodic Table to determine the number of valence electrons zEach group has the same number of valence electrons 1A 2A 3A 4A 5A 6A 7A 8A

23. A. Periodic Law zWhen elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

24. B. Chemical Reactivity zFamilies ySimilar valence e - within a group result in similar chemical properties

25. zAtomic Radius ysize of atom © 1998 LOGAL zIonization Energy yEnergy required to remove an e - from a neutral atom © 1998 LOGAL zElectronegativity C. Other Properties

26. zAtomic Radius = ½ the distance between two identical bonded atoms 1. Atomic Radius

27. zAtomic Radius yIncreases to the LEFT and DOWN 1. Atomic Radius

28. zWhy larger going down? yHigher energy levels have larger orbitals yShielding - core e - block the attraction between the nucleus and the valence e - zWhy smaller to the right? yIncreased nuclear charge without additional shielding pulls e - in tighter 1. Atomic Radius

29. zFirst Ionization Energy = Energy required to remove one e - from a neutral atom. 2. Ionization Energy K Na Li Ar Ne He

30. zFirst Ionization Energy yIncreases UP and to the RIGHT 2. Ionization Energy

31. zWhy opposite of atomic radius? yIn small atoms, e - are close to the nucleus where the attraction is stronger zWhy small jumps within each group? yStable e - configurations don’t want to lose e - 2. Ionization Energy

32. zSuccessive Ionization Energies yMg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ yLarge jump in I.E. occurs when a CORE e - is removed. 2. Ionization Energy

33. yAl1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ zSuccessive Ionization Energies yLarge jump in I.E. occurs when a CORE e - is removed. 2. Ionization Energy

34. 3. Electronegativity zThe measure of the ability of an atom in a chemical compound to attract electrons zGiven a value between 0 and 4, 4 being the highest

35. zWhy increase as you move right? yMore valence electrons, need less to fill outer shell zWhy increase as you move up? ySmaller electron cloud, more pull by + nucleus 3. Electronegativity

36. zWhich atom has the larger radius? yBe orBa yCa orBr Examples

37. zWhich atom has the higher 1st I.E.? yNorBi yBa orNe Examples

38. zWhich element has the higher electronegativity? yCl or F yBe or Ca Examples