1. Chapter 5- The Periodic Law 5.1-History of the Periodic Table 5.2-Electron Configuration & the Periodic Table 5.3-Electron Configuration & Periodic Properties

2. 5.1-History of the Periodic Table Pages 123-127

3. Mendeleev  Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass. Elements with similar properties were grouped together. There were some discrepancies.

4. Mendeleev  Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered elements.

5. Moseley  Henry Moseley (1913, British) Organized elements by increasing atomic number. Resolved discrepancies in Mendeleev’s arrangement. Periodic Law-the physical and chemical properties of the elements are periodic functions of the atomic numbers.

6. Organization of the Elements  Periodic table-an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

7. Additions to Mendeleev’s Periodic Table  Noble gases Group 18 Argon discovered in 1894 Took so long to discover because very unreactive  Lanthanides 14 elements with atomic numbers from 58-71 Placed below the periodic table to conserve space  Actinides 14 elements with atomic numbers 90-103 Also placed below periodic table

8. 5.2-Electron Configuration & the Periodic Table Pages 128-139

9. Periods & Blocks of the Periodic Table  Length of period (row) determined by how many electrons can occupy the sublevels being filled. 1 st period-1s sublevel being filled with 2 electrons  2 elements, H & He 3 rd period-3s & 3 p sublevels being filled with 2+6 electrons  8 elements  Periodic table is divided into “blocks” based on the filling of sublevels with electrons.

10. Blocks of the Periodic Table

11. Determining Period from Configuration  An element’s period can be determined by looking at its electron configuration  The highest occupied energy level corresponds to the element’s period As: [Ar]3d 10 4s 2 4p 3  4 in 4p 3 indicates that the highest energy level that electrons occupy is the 4 th. Therefore, As is located in the 4 th period of the periodic table.

12.  Metals  Nonmetals  Metalloids Metallic Character

13.  Main Group Elements  Transition Metals  Inner Transition Metals Areas of the Periodic Table

14. s-Block Elements: Groups 1 & 2  Chemically reactive metals  Include the alkali metals and the alkaline earth metals

15. Alkali metals  Group 1 metals  ns 1  Silvery appearance and very soft  Not found pure naturally because so reactive  Because of extreme reactivity with moisture, usually stored under kerosene  Video: Disposal of Surplus SodiumDisposal of Surplus Sodium  Video: Alkali Metals in WaterAlkali Metals in Water

16. Alkaline-Earth metals  Group 2 metals  ns 2  Harder, denser, & stronger than alkali metals  Also too reactive to be found free in nature (but less reactive than Gp. 1)  Video: Magnesium/silver nitrate mixture reacting with waterMagnesium/silver nitrate mixture reacting with water

17. d-Block Elements: Groups 3-12  Metals with typical metallic properties  Called “transition elements”  Typically less reactive than Gps. 1&2, & some are extremely unreactive  d sublevels first appears at the 3 rd energy level  Fills after 4s  Variations from expected in d-block, so elements in the same group do not necessarily have the same outer e- configuration

18. p-Block Elements: Groups 13-18  p and s-block elements together called “main-group elements”  Total number of electrons in highest energy level=group # - 10 Group 17 elements have 17-10=7 outer “valence” electrons  Properties of p-block elements vary greatly since metals, nonmetals, and metalloids are contained here

19. p-block Elements  Halogens Group 17 nonmetals Most reactive nonmetals  React with most metals to form salts  Metalloids Fall on both sides of a “stair-step” line separating metals and nonmetals Semi-conductors

20. f-Block Elements: Lanthanides & Actinides  Lanthanides Top row of f-block 14 elements Shiny metals similar in reactivity to the alkaline-earth metals  Actinides Bottom row of f-block 14 elements All radioactive 1 st 4 elements found naturally on Earth; remainder only lab-made elements

21. 5.3-Electron Configuration & Periodic Properties Pages 140-154

22. Remember the Periodic Law  When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

23.  ½ the distance between the nuclei of identical atoms that are bonded together  Increases to the LEFT and DOWN Atomic Radius

24. Li Ar Ne K Na

25.  Why larger going down? Higher energy levels have larger orbitals Shielding - core e - block the attraction between the nucleus and the valence e -  Why smaller to the right? Increased nuclear charge without additional shielding pulls e - in tighter Atomic Radius

26.  First Ionization Energy-energy required to remove one electron from a neutral atom  Increases UP and to the RIGHT Ionization Energy

27.  First Ionization Energy Ionization Energy K Na Li Ar Ne He

28.  Why opposite of atomic radius? In small atoms, e - are close to the nucleus where the attraction is stronger  Why small jumps within each group? Stable e - configurations don’t want to lose e - Ionization Energy

29.  Successive Ionization Energies Mg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ Large jump in I.E. occurs when a CORE e - is removed. Ionization Energy

30. Al1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ  Successive Ionization Energies Large jump in I.E. occurs when a CORE e - is removed. Ionization Energy

31.  Energy change that occurs when an electron is acquired by a neutral atom  Tends to become less negative (less energy released) DOWN and to the LEFT Electron Affinity

32.  Ionic Radius Cations (+)  lose e -  smaller © 2002 Prentice-Hall, Inc. Anions (–)  gain e -  larger Ionic Radius

33. Electronegativity  A measure of the ability of an atom in a chemical compound to attract electrons  Most electronegative element is fluorine Given arbitrary value of 4; all others relative

34.  Which atom has the larger radius? BeorBa CaorBr Ba Ca Examples

35.  Which atom has the higher 1st I.E.? NorBi BaorNe N Ne Examples

36.  Which has the greater electonegativity? KorLi AlorCl Li Cl Examples

37.  Which particle has the larger radius? SorS 2- AlorAl 3+ S 2- Al Examples