1. Hydrogen and Helium Hydrogen does not share the same properties as the elements of group 1. Helium has the electron configuration of group 2 elements however it behaves like group 18 (Noble Gases)

2. f-block elements Lanthanides & Actinides Lanthanides – shiny metals Actinides – radioactive; only the 1st four are found naturally

3. The Octet Rule All atoms want their outer shell (highest energy level) to have 8 electrons. If the outer shell has eight, the atom is stable (non-reactive)

5. Valence Electrons e- available to be lost, gained, or shared in forming chemical compounds these e- are in highest energy level # of valence electrons? Li C Ca Ne As At

6. Ions Ions are formed when: a metal gives up valence electrons a nonmetal gains valence electrons They do this so their highest energy level has 8 valence electrons = octet

7. IONS This transfer of electrons creates an element (ion) that is no longer neutral Now the protons ≠ electrons Cations are ions with a positive charge Anions are ions with a negative charge

8. Cations (The “t” in “cation” looks like a plus sign) All metals form positive cations when they give electrons from their valence shell. Alkali metals (1 valence e- to give away), so they become a ion with +1 charge b/c they have 1 more p+ than e-) ex. Li +, Na +, K + Alkaline Earth Metals (2 valence e- to give away), so… ex. Ca +2, Ba +2

9. Common Ionic Charges

10. Anions (“A-negative-ion”) All non-metals form anions when they gain electrons from metals to fill their outer shell with a total of eight. More e- than p+ means they have a - charge All halogens (7 valence electrons); so they form an ion with a -1 charge b/c gained 1e- ex. Cl -, F -, Br -

11. Common Ionic Charges

12. Your Turn! potassium ion? K +1 oxygen ion? O -2 magnesium ion? Mg +2 noble gases? Already stable!

13. Ionic Radii (Trend) Cations (+ ions) are smaller b/c they lose e- Anions (- ions) are bigger b/c they gain e- Metals (left side p.table) form cations Nonmetals (right side p.table) form anions Ionic Radii INCREASES as you go across and down the p.table

14. Atomic Size - (Trend) As we increase the atomic number (or go down a group)... each atom has another energy level, so the atoms get bigger. H Li Na K Rb

15. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. Electrons are in the same energy level. But, there is more nuclear charge. Outermost electrons are pulled closer. NaMgAlSiPSClAr

16. Atomic Number Atomic Radius (pm) H Li Ne Ar 10 Na K Kr Rb 3 Period 2

17. Ionization Energy (Trend ) Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom). Removing one electron makes a 1+ ion. The energy required to remove only the first electron is called the first ionization energy.

18. Ionization Energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE.

19. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

20. Ionization Energy - Group trends As you go down a group, the first IE decreases because... The electron is further away from the attraction of the nucleus, and There is more shielding.

21. Ionization Energy - Period trends All the atoms in the same period have the same energy level. Same shielding. But, increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 full orbitals.

22. First Ionization energy Atomic number He He has a greater IE than H. Both elements have the same shielding since electrons are only in the first level But He has a greater nuclear charge H

23. First Ionization energy Atomic number H He l Li has lower IE than H l more shielding l further away l These outweigh the greater nuclear charge Li

24. First Ionization energy Atomic number H He l Be has higher IE than Li l same shielding l greater nuclear charge Li Be

25. First Ionization energy Atomic number H He l B has lower IE than Be l same shielding l greater nuclear charge l By removing an electron we make s orbital filled Li Be B

26. First Ionization energy Atomic number H He Li Be B C

27. First Ionization energy Atomic number H He Li Be B C N

28. First Ionization energy Atomic number H He Li Be B C N O Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital

29. First Ionization energy Atomic number H He Li Be B C N O F

30. First Ionization energy Atomic number H He Li Be B C N O F Ne Ne has a lower IE than He Both are full, Ne has more shielding Greater distance

31. First Ionization energy Atomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

32. First Ionization energy Atomic number

33. Electronegativity Electronegativity refers to the ability of an atom to attract the electrons from another atom bonded together in a compound. EA (e- affinity) deals with an isolated atom and Electronegativity deals with 2 atoms bonded together

34. Electronegativity Trends (same trend as EA) Electronegativity generally increases as you move across the periodic table Electronegativity decreases as you move down a group because of the shielding effect

35. Electronegativity

36. Sample Problems Which has the greatest electronegativity? Cl, S, P, Al Ra, Ba, Ca, Be Fr, Ba, Cs, Al

37. Element with Highest Electronegativity? Fluorine, has the highest electronegativity.

38. Electron Affinity EA The change in energy (always negative) when you add an e- to an atom. Affinity = “you like it” EA increases (more negative #) as you move across because those atoms want to gain e- to gain an octet EA decreases as you move down because atoms want to gain e- less because of the shielding effect Noble gases have an EA of ZERO, b/c they already have an octet

39. Atomic Radii Decrease as You Move Across a Period Why?????? The atomic radius decreases because the greater positive charge (more p+) in the nucleus pulls the electrons closer.. (opposite charges attract)

40. Atomic Radii Increase as you Move Down a Group WHY????? As you move down you add energy levels which increases the size and radius of the atom.

42. Sample Problems Ca, Be, Ba, and Sr: largest atomic radius? Al, Mg, Si, and Na: smallest radius? Largest atomic radius? Li O C F

43. Ionization Energy The IE is the minimum amount of energy required to remove an electron from the outer shell of a neutral atom. energy to remove the first e- is called the “first ionization energy” IE 1

44. Ionization Energy Trend IE (energy to take an e-) increases as you move across because non-metals want to gain e–, not lose their e- nuclear charge increases due to adding more # p+, so it takes more energy to remove an electron. IE decreases as you move down the periodic table because the electrons are farther away from the nucleus and easier to remove. Why? all the electrons in the many energy levels between repel the valence e- = shielding effect