1. Periodic Trends and Bonding Chapters 5 & 6

2. Ions and valence electrons How many valence electrons are in the following elements? Na Mg H He Cl Al

3. Show your Understanding Draw the Lewis dot structures for each of those elements What ions will those elements form?

4. S- block Chemically reactive metals Group #1= alkali metals Slippery appearance and can be cut with a knife!  For real? They all have one valence electron Combine readily with the halogens to form salts Group #2- Alkaline Earth Metals Harder, denser and stronger than group #1 metals Have 2 valence electrons

5. d- block - Transition Metals Lowest quantum # = 3 Maximum # of electrons = 10 There are exception to the electron configuration rules Some metals may form several different ions They are all metals and good conductors of heat and electricity and have high luster. Properties vary greatly. Some metals are highly reactive Other metals not so much- Au, Pt, Pd

6. p-block – Groups 13-18 Group 13- 3 valence electrons Group 14 – 4 valence electrons Group 15- 5 valence electrons etc. etc. Contains metals, non-metals and metalloids. Important group- #17- Halogens Most are gases- most reactive with metals

7. Main Group elements Elements found in the s block and p block Only elements that can be used in Lewis Dot Structures

8. Key Concepts/ Terms Electrons have an attraction or pull towards the nucleus of the atom. Shielding/ Screening: the attraction of outer shell electrons is counterbalanced by the repulsion of the inner-shell electrons. The inner-shell electrons “screen” or “shield” the outer-shell electrons from full attraction

9. Atomic Radii TREND: Increases from top to bottom Decreases from left to right Arrange the following in increasing atomic radius: Rb, In, Sb, Sr, I

10. Ionization Energy The energy required to remove an electron (IE 2 ) = energy req. to remove a second electron The more electrons removed from an atom the greater the IE TREND: Increases from left to right Decreases from top to bottom Exceptions: ( within the same energy level) Group 3A – lower energy than 1A and 2A bc entering p orbital, slightly higher in energy than the s orbital for the same level Group 6A – electrons are paired up in the p-orbital and the IE slightly dips, making it easier to remove

11. Ionization Energy Arrange the following elements in order of decreasing IE. Na, Mg, Al, Si

12. Ionization Energy Arrange the following elements in order of decreasing IE. Na, Mg, Al, Si Si < Al < Mg< Na

13. Electron Affinity Energy absorbed when an electron is added TREND: Greater Attraction for electrons from left to right  ( Left to right = least negative to most negative) Decreases from top to bottom

14. Electron Affinity Which group would be the easiest to add an electron to? Arrange the following elements in increasing electron affinity: C, Si, Ge Cl, S, P

15. Electron Affinity - answer Which group would be the easiest to add an electron to? The halogens Arrange the following elements in the most negative to least negative electron affinity: C, Si, Ge – same Cl, S, P – same

16. Ionic Radii- DO NOT DO Term: Isoelectronic : species that have the same number of electrons. ( Na +, Mg 2+ ) TREND: Top to Bottom = increases Left to right = decreases

17. Ionic Radii Arrange the following ions in order of increasing ionic radii: Na +, Tl 3+, Mg 2+,

18. Ionic Radii - answer Arrange the following ions in order of increasing ionic radii: Mg 2+, Na +, Tl 3+,

19. Electronegativity Ability for an atom to attract electrons When it is chemically combined with another atom. Elements with high electronegativities (nonmetals) often gain electrons to form anions. Elements with low electronegativities (metals) often lose electrons to form cations.

20. Electronegativity TREND: Top to Bottom = decreases Left to right = increases