1. Chapter 14: Periodic Trends …and naming ions (chapter 6)

2. Review: Energy Levels Principle quantum numbers (n) = energy level – Lower the number, lower the energy

3. Put it all together…

4. A new way to write electron configurations

5. Use the Noble gas abbreviation, write the electron configurations for: Na: Cl: W: Sn:

6. Organizing the Table Noble Gases – – Representative Elements- (s and p block) – Transition Metals (d block) – Inner Transition Metals –

7. Representative Elements Sometimes called “group A elements” = alkali metals = alkaline earth metals = halogens = noble gases The group number equals the number of electrons in its outermost energy level (this will be more important later on…)

8. Other things you need to know… (Ch. 6) When atoms gain or lose electrons they form IONS (positively or negatively charged atoms). Positive ions = cations Negative ions = anions

9. Cations Examples: Calcium and Magnesium You can determine how many electrons are lost based on location on the periodic table. These must be memorized. – Group

10. Naming Ions How to name: name of element + “ion” Examples:

11. Transition Metals… Can form different charges (you can’t memorize) Here is how you know the charge:

12. Anions Examples: Bromine and Nitrogen You need to MEMORIZE the common charges of the anions to be successful for the rest of the school year…

13. Naming Anions Anions change the ending of the element – unlike cations Stem-ide ion Examples: Chlorine = Oxygen =

14. Ion Size _______________ are smaller in size than the neutral element. _______________ are larger in size than the neutral element.

15. Polyatomic ions Polyatomic = Ions = You will get a list of 10 polyatomic ions. You must memorize the name and formula and be able to recall them at any time after the first test (i.e. I won’t feel guilty if there is a pop quiz).

16. Types of Trends _____________________________: Trends across a period (row) of the periodic table. _____________________________: Trends down a group (family) of the periodic table

17. Nuclear Charge (p+ in nucleus) Periodic: Group:

18. Atomic Radii and Size Periodic: Atomic Radii and size _________________ 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 ___________________ as you go down a group – (increasing principle energy level)

20. Ionization Energy Ionization Energy (IE): – Remove 1 st electron = 1 st IE – Remove 2 nd electron = 2 nd IE

21. Trends in Ionization Energy Periodic: IE generally ___________________ 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 – __________________________ Group: 1 st IE generally __________________ as you go down a group. – Atom gets bigger, outermost e- farthest from nucleus, easy to be removed.

22. Electronegativity Electronegativity:

23. Electronegativity Trends Periodic: ________________________________ as you go L to R across the period – Group: _________________________________ as you go down a group. The most electronegative element is Fluorine.

24. Summary Decreasing Increasing Atomic Radius Electronegativity Ionization Energy