1. The Periodic Law Chapter 5 Sophia Nolas, Yen Dinh, Chris Fleming, Jane Smigiel

2. SECTION ONE (pg 123) History of the Periodic Table

3. Dmitri Mendeleev Dmitri Mendeleev He created a periodic table of the elements, grouping them with other elements with similar properties.

4. Henry Moseley Henry Moseley His observations on his data led to both the modern definition of atomic number and how it is the basis for the organization of the periodic table.

5. Periodic Law Periodic Law : The physical and chemical properties of the elements are periodic functions of their atomic numbers. Periodic Table Periodic Table : An arrangement of the elements where similar properties fall in the same column, or group.

6. Lanthanides Lanthanides : Are the 14 elements with atomic numbers from 58 to 71. These elements are so similar in chemical and physical properties that it was a tedious task for many chemists to identify them. Actinides Actinides : Are the 14 elements with atomic numbers from 90 to 103. Lanthanides and Actinides sit below the main portion of the periodic table to save space.

7. SECTION TWO (pg 128) Electron Configuration and the Periodic Table

8. S sublevel = 2 electrons P sublevel = 6 electrons D sublevel = 10 electrons F sublevel = 14 electrons

9. Electron Configuration Electron Configuration : Carbon (C) = 1s 2 2s 2p 2 Hydrogen (H) = 1s Copper (Cu) = 1s 2 2s 2 2p⁶ 3s 2 3p⁶ 4s 2 3d⁹ Coefficient Coefficient = the period where element is located, (EXCEPTION = d orbitals are off by -1.) Letter Letter = what sublevel Exponent Exponent = how many elements in the period and sublevel to reach “goal” element (group number).

10. Orbital sublevels (the letters in the electron configuration)

11. Electron configuration explanation http://www.khanacademy.org/video/more-on- orbitals-and-electron- configuration?playlist=Chemistry

12. A simpler way to show electron configuration: Take previous noble gas and put it in [ ]. Then find the electron configuration from there. Example: Silver – Ag Previous noble gas = Kr [Kr] 5s 2 4d⁹

13. Alkali Metals – Group 1 Group 1 of the periodic table One electron in S sublevel Do not occur in nature as free elements Good conductors with low melting points Ductile and malleable (can be cut with a knife) Very reactive (with water to make hydrogen gas) Usually held in kerosene due to the high reacting level Used in lights, electricity technology

14. Alkaline-Earth Metals – Group 2 Group 2 of the periodic table Contain electron pair in S sublevel Harder and denser than alkali metals (AM) Too reactive to be found in nature, but not as reactive as AM

15. Halogens – Group 17 Most reactive nonmetals React vigorously with metals

16. Questions! Name the group, period and block of [Xe]6s 2 Write the electron configuration of Group 1 element in the 3 rd period Which element is more reactive?

17. Answers! Group = 2 nd Period = 6 th Block = s [Ne] 3s The first one

18. SAMPLE PROBLEMS Give the group, period, and block for these elements (without periodic table), and then determine the name of the element (with the periodic table)  [Ne] 3s²3p 1  [He] 2s 1  [Ar] 4s 2 3d 9 Write the outer electron configuration of these elements (without the periodic table) and then identify the element (with the periodic table)  Group 2, Period 6  Group 7, Period 4  Group 14, Period 3 FOR MORE PRACTICE, GO TO PAGES 133, 136, 138, AND 139 IN THE TEXTBOOK

19. SECTION THREE (pg 140) Electron Configuration and Periodic Properties

20. The way in which the elements arranged on the periodic table corresponds to not only their atomic number but also electron configuration – This creates periodic trends in the properties of these elements

21. Atomic Radii One-half the distance between the nuclei of identical atoms that are bonded together Trend Trend: Increases from top to bottom and from left to right (Towards francium) – Due to the decrease of positive and negative charges from less protons and electrons and the attractive forces between them – The electrons are pulled closer to the nucleus from right to left & from bottom to bottom on the periodic table (which is why the A.R. decreases in these direction)

22. ATOMIC RADII CHART

23. Ionic Radius One half the diameter of an ion in an ionic compound – Cation: positive ions that have lost an electron, thus are smaller than neutral atoms – Anion: negative ions that have gained an electron, thus are larger than neutral atoms Trend Trend: Increases from top to bottom and from right to left

24. Electronegativity How readily an atom will attract an electron(s) in a chemical compound in order to achieve noble gas configuration – Combination of Ionization energy &n Electron affinity Trend Trend: Increases from bottom to top and from left to right

25. ELECTRONEGATIVITY Chart (Key: more red=more electronegative)

26. Ionization Energy The energy required in order to remove and electron from an atom to create an ion – A + energy (A+) & e- Increases from bottom to top and from left to right – As you move down the groups, the more outer shell electrons you have, creating a shielding effect on the outermost electrons, making it easier to remove Electrons from the atom

27. Electron Affinity The energy change that occurs when and electron is added to a neutral atom – Can be endothermic or exothermic Increases from bottom to top and from left to right – Second electron affinity is always endothermic: it requires energy to force electron to bond to the atom

28. SAMPLE PROBLEMS List the following elements from biggest to smallest atomic radius – He, Fr, Zn, Ge, C List the following from highest to lowest electronegativity – V, Os, Cd, Sn, Mg List the following from highest to lowest ionization energy – H, Al, P, Sr, Re List the following from highest to lowest electron affinity – O, As, Li, Au, Rb FOR MORE PRACTICE GO TO PAGES 142, 146, AND 152

29. FOR ADDITIONAL PRACTICE… Chapter Summery + Reviewpg 155 Periodic Table pg 130