1. Periodic Table 1

3. The periodic table is a systematic arrangement of the elements by atomic number (protons) Similar properties fall into vertical columns 3

4. History of the Periodic Table Three men recognized patterns in the elements. They attempted to organize the elements according to these patterns.. 4

5. History of the Periodic Table Johann Wolfgang Döbereiner – Noticed patterns in atomic mass recurring in sets of three elements Became known as “Döbereiner's triads” 5

6. History of the Periodic Table John Newlands noticed every eighth element had similar properties. Known as 'law of octaves' : 6

7. History of the Periodic Table Dmitri Mendeleev developed the first periodic table Found the repeating pattern by atomic mass and arranged them so that groups of elements with similar properties fell into vertical columns in his table. Found a problem Some elements fell into the wrong column Examples: Te & I ; Co & Ni 7

8. Mendeleev’s Periodic Table 8

9. History of the Periodic Table Henry Moseley Fixed Mendeleev’s problem by rearranging the modern table by atomic number Used X-ray spectrometer to find the atomic numbers 9

10. Arrangement of Periodic Table Periodicity– trends of properties as you go across the table or down a column 10

11. horizontal rows there are 7 Period number tells which energy level holds the valence electrons Periods 1 2 3 4 5 6 7 11

12. Groups/Families vertical columns groups 1-18; elements in the same group share chemical properties Main group elements Groups 1,2 13, 14, 15, 16, 17, 18 12

13. Types of Elements Noble gases 13

14. Metals Found on the left side of table 1, 2 or 3 valence electronsHave 1, 2 or 3 valence electrons Lose electrons to form positive ions (cations) Most are silver, shiny, solid, malleable, ductile & good heat/electrical conductors

15. Nonmetals Found on the right side of table Have 5, 6, or 7 valence electrons Gain electrons to form negative ions (anions) Brittle, dull, non-conductors, and exist in all three states (solids, liquids, gases)

16. Metalloids Elements found along the stair-step between metals and nonmetals, NOT Al Properties are in between metals & nonmetals Silicon (Si) is probably the most well- known metalloid.

17. Noble Gases odorless, colorless, monatomic gases low chemical reactivity.

18. Color Groups of the Periodic Table Alkali Metals Alkaline Earth Metals Transition Metals Lanthanide Series Halogens Noble Gases Actinide Series Inner Transitional Metals Also called inert gases because they do not react Metalloids 18

19. Properties and Electron Configuration Look- each group (column) ends with the same electron configuration. That determines many of the physical properties that the group share. 19

20. Group 1 Based on the video Alkali Metals with WaterAlkali Metals with Water 1.What properties of Alkali metals are observed? 2.What trend is observed as samples are tested with water? 3.Why weren’t hydrogen and francium tested? 20

21. Group 17 1.What are some of the physical properties of the halogens? Halogen 21

22. Group 18 Note: In the video “Group 0” is an old name for Group 18. 1.Why are the noble gases un-reactive? 2.If all neon signs were made of pure neon gas, what colors would we have? 3.What are uses for noble gases other than in neon lights? 4.How can a physical property be used to tell the difference between noble gases? 5.Radon was not tested. Predict what a balloon filled with Radon would do when dropped from the roof and why. Noble Gases 22

23. Summary of Groups, Props. & Electrons NOVA Video 1.What is the relationship between electron configuration and group number on the periodic table? 2.Why are halogens and alkali metals highly reactive, but not the noble gases? NOVA Video 23

24. Periodic Table Trends Patterns on the periodic table – Atomic Radius – Ionic Radius – Electronegativity – Ionization Energy – Reactivity 24

25. Periodic Trends- similarities of elements based on where they are in the table Depend on two things: Effective Nuclear Charge- The attraction the valence electrons have for the protons in the nucleus. Electron Shielding Effect- Inner shell electrons blocking valence electrons from the nucleus. 25

26. Effective Nuclear Charge Watch this video And this Effective Nuclear Charge is abbreviated Zeff Smart folks have noticed that the zeff for each group is equal to the number of valence electrons. 26

27. Atomic Radius Atomic radius is half the distance between the centers of two atoms measured in angstroms. The more energy levels, the ________ the atomic radius. (larger/smaller) The higher the effective nuclear charge, the ________ the atomic radius. (larger/smaller) larger smaller 27

28. Atomic Radius Trend Atomic radius increases as you move down a group Atomic radius decreases as you move from left to right in a period Down the group the number of energy levels increase so the number of shielding electrons increase. The nucleus cannot pull in the valence electrons. That makes a bigger atom. Across the period the number of protons increases while the number of shielding electrons stays the same. This make the nucleus pull in the valence electrons. That makes a smaller atom. 28

29. Ions Cations Form from metals Lose electrons Metal have low effective nuclear charge holding on to the valence electrons. Anions Form from nonmetals Gain electrons Nonmetals have high effective nuclear attraction on the valence electrons 29

30. Ions Metals lose electrons to form cations Li Li + F F-F- Nonmetals gain electrons to form anions Ionic radius is smaller than atomic radius energy level is lost or “shed” Ionic radius is larger than atomic radius because the electrons outnumber the protons. The nucleus has less control of the valence electrons. 30

31. Electronegativity Electronegativity s a measure of how strongly atoms attract bonding electrons to themselves An assigned number “rates” the electronegativity (from 0.7 to 4.0) – Low electronegativity = cannot attract valence electrons – High electronegativity = can attract valence electrons

32. Electronegativity Trend Electronegativity values increase as you move from left to right in any period. Within any group, electronegativity values decrease as you go down. Biggest IE = Fluorine Smallest IE = Francium 32

33. Electronegativity- EN- the tendency of an atom to pull shared electrons to itself. High EN= Big pull F9pF9p Li 3p Be 4p B 5p C 6p N 7p O 8p Factors affecting Electronegativity -Size of the atom/distance- small size/distance the nucleus has a stronger attraction for electrons Why does the trend decrease down a group? 33

34. Ionization Energy Ionization Energy – the energy needed to remove the outermost electron in an atom. How hard is it to steal an electron Increases as you go right in a period Larger nuclear charge – more protons pulling on the electrons Atom is smaller – outer electrons are closer to the nucleus; easier to pull in electrons Decrease as you go down in a group More energy levels – Radius is larger; outer electrons are farther from the nucleus; more difficult to gain electrons 34

35. Ionization Energy Pattern 35

36. Ionization Energy Decreasing ionization energy Increasing ionization energy 36

37. Metal Reactivity Trend Metal Activity depends on the attraction the metal has for the nonmetals electrons. Trend Increases as you move down a group Decreases as you move from left to right in a period *The most reactive metal is francium decreasing metal reactivity increasing metal activity

38. Nonmetal Activity Trend Non-Metal Activity refers to how easily nonmetals gain e- to form anions Trend Decreases as you move down a group Increases as you move from left to right in a period *The most reactive nonmetal is fluorine increasing nonmetal activity decreasing nonmetal activity