1. History of Periodic Table Chapter 5

2. History 1860s – 60 elements discovered –Cannizzaro - agreed on method to measure atomic mass –Search for relationships between properties of elements

3. Dimitri Mendeleev Organized elements by increasing atomic mass Noticed chemical and physical properties followed trend, or pattern  Periodic

4. Mendeleev’s Table

5. Henry Moseley Worked with Rutherford looking at line- spectras Noticed better pattern when elements were organized by increasing atomic # Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic #s

6. Regions of Periodic Table Group Project

7. Main Group Elements s and p block elements

8. Group 1A are the alkali metals Group 2A are the alkaline earth metals

9. Group 7A is called the Halogens Group 8A are the noble gases

10. The group B are called the transition metals

11. Top: Lanthanide Series Bottom: Actinide Series

12. Periodic Properties

13. Atomic Radii (Atomic Size) Def: half the distance between the nuclei of identical atoms that are bonded together } Radius

14. Atomic Radii - Group trends As we go down a group Another energy level… So the atoms get bigger. H Li Na K Rb

15. Atomic Radii - Periodic Trends Go across a period the radius gets smaller. Same energy level. More nuclear charge. Outermost electrons pulled in closer NaMgAlSiPSClAr

16. Ionization Energy (IE) An e - can be removed from any atom if there is enough energy A + energy  A + + e - Ion: atom or group of bonded atoms that has a (+) or(-) charge Process that results in ion formed is ionization

17. Valence Electrons Def: The e - available to be lost, gained or shared to form chemical compounds e - found in the outermost s and p sublevels

18. Ionization Energy (IE) Def: the energy needed to remove one e - from an atom (IE 1 – first IE) Atoms with HIGH ionization energy  hold on tight to their electrons

19. IE – Group Trends As you go down a group IE decreases Electron further away from nucleus Less attraction to nucleus, easier to take e -

20. IE – Periodic Trends IE generally increases from left to right Increasing nuclear charge More nuclear charge holds on tight to e - Exact opposite of atomic radius

21. IE 2 and IE 3 Energy required to remove additional e - Energies keep getting higher and higher e - that are left are being held closer to nucleus  harder to remove Pg. 155

22. Ionic Radii (Ionic Size) Cation: positive ion –Always smaller than atom –Lost e -, now nucleus pulling in more on remaining e - s Anion: negative ion –Always bigger than atom –Gaining e -, now e - are crowded and spread out (repulsion of like charges)

23. Ionic Radii – Group Trends Same as Atomic Radii More energy levels as go down  size increases

24. Ionic Radii – Periodic Trends 2 sections Metals on LEFT make CATIONs Nonmetals on RIGHT make ANIONS Cations (1A – 4A) Anions (5A – 8A) –Decrease as go across (L-R) due to increase nuclear charge

25. Electronegativity Valance e - are involved in forming bonds Some atoms in a chemical bond attract the valance e - more than the other (tug of war) Linus Pauling – electronegativity – measure of the ability of an atom in a chemical compound to attract e - from another atom in the compound

26. Electronegativity – Group Trends Tend to decrease down a group or remain about the same Noble gases are NOT assigned electronegativities

27. Electronegativity – Periodic Trends Tend to increase as you go across the table F – most electronegative Fr – least electronegative