1. Review… How do you tell metals from nonmetals on the periodic table?

2. Characteristics of Metals? Good conductors of heat and electricity Shiny (have luster) Malleable (bendable)

3. Characteristic of Nonmetals Poor (bad) conductors of heat and electricity Solids are brittle (break instead of bend) Generally the opposite of metals

4. Where are the metalloids?

5. What are metalloids? Sometimes they have characteristics of metals Sometimes they have characteristics of nonmetals

6. Groups names… 1A Alkali Metals 2A Alkaline Earth Metals 7A Halogens 8A Nobel Gases

7. Middle Group are the….

8. Two rows on the bottom are… Inner Transition Metals

9. What are the representative elements? Representative elements: elements in groups 1A – 7A * Wide range of physical and chemical properties

10. Electrons are the Key! The reason elements in the same family have the same chemical and physical properties is because they have the same number of electrons in their outer shell

11. The group number equals the number of electrons in its outermost energy level (this will be more important later on…) So all Group 1A elements have 1 electron in the outer shell

12. Another way of organizing the periodic table…

13. Groups 1A and 2A (including He) are in the s block Groups 3A – 8A are in the p block (excluding He) Transition metals are d block Inner transition metals are f block This all has to do with electron configurations which we will learn later! Just know the blocks and what groups are in each block

14. Atoms can gain or lose electrons

15. When atoms gain or lose electrons they form IONS (positively or negatively charged atoms). REMEMBER electrons are negatively charged! If an atom gains electrons it will have a negative charge If an atom loses electrons it will have a positive charge

16. Cations Positive ions = cations Formed when a metallic atom LOSES electrons. Examples: Sodium  Na +1 Magnesium  Mg +2 You can determine how many electrons are lost based on location on the periodic table. These must be memorized.

17. Anions Negative ions = anions Non-metals gain electrons become anions. Examples: Bromine  Br -1 Oxygen  O -2 You need to MEMORIZE the common charges of the anions to be successful for the rest of the school year…

18. Group 1A = +1 Group 2A = +2 Group 6A = -2 Group 7A = -1 Now onto PERIODIC TRENDS…..

19. Atomic Radii and Size Periodic: Atomic Radii and size decrease as you go L to R across the table. –Same principle energy level –Add p+ and e-, increase nuclear charge, pulls in orbitals closer to the nucleus Group: Atomic Radii and size increase as you go down a group –Electrons being added to outer orbital (increasing principle energy level)

21. Ion Size Cations are smaller in size than the neutral element. Anions are larger in size than the neutral element.

22. Ionization Energy Ionization Energy (IE): The energy required to remove an electron from a gaseous atom. –Remove 1 st electron = 1 st IE –Remove 2 nd electron = 2 nd IE

23. Trends in Ionization Energy Periodic: IE generally increases as you move L to R across the period. –Harder to remove an electron as you go L to R because of greater attraction to nucleus –Shielding effect - Group: 1 st IE generally decreases as you go down a group. –Atom gets bigger, outermost e- farthest from nucleus, easy to be removed.

25. Electronegativity Electronegativity: The tendency for atoms to attract electrons when they are chemically combined. –Do they share electrons equally?

26. Electronegativity Trends Periodic: Electronegativity increases as you go L to R across the period –Elements want to be like noble gases! Group: Electronegativity decreases as you go down a group. The most electronegative element is Fluorine.

28. Summary of Trends Increasing Electronegativty Increasing Ionization Energy Decreasing Atomic Radius Decreasing Ionization Energy Decreasing ElectronegativityIncreasing Atomic Radius