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

2. The Modern Periodic Table

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

4. Sublevels – s, p, d, f, g

5. Put it all together…

6. A new way to write electron configurations

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

8. Organizing the Table Noble Gases – –Filled outermost s and p orbitals (s 2 p 6 ) Representative Elements- (s and p block) –Partially filled outermost s and p orbitals Transition Metals (d block) –Outermost s and nearby d orbitals contain electrons Inner Transition Metals –Outermost s and nearby f contain electrons

10. Representative Elements Sometimes called “group A elements” 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases The group number equals the number of electrons in its outermost energy level (this will be more important later on…)

11. Metals and Nonmetals From here on out you need to identify whether an element is a metal or nonmetal (FAST and accurate).

12. Other things you need to know… (Ch. 6) When atoms gain or lose electrons they form IONS (positively or negatively charged atoms). Charge determined by difference between p+ and e- Positive ions = cations (lose valence e-) –Metal atoms lose valence e- Negative ions = anions (gain valence e-) –Nonmetal atoms gain valence e-

13. Cations Formed when a metallic atom LOSES electrons. Examples: Calcium and Magnesium calcium ion magnesium ion You can determine how many electrons are lost based on location on the periodic table. These must be memorized. –Group 1A, 2A, Al, Ag, Zn

14. Naming Ions How to name: name of element + “ion” Examples: Cations –Aluminum (Al): Aluminum Ion (Al 3+ ) –Sodium (Na): Sodium Ion (Na + )

15. Transition Metals… Can form different charges (you can’t memorize) Here is how you know the charge: –They are metals, so are cations. –Roman Numerals (I, II, III, IV, V, VI, VII) indicate how many e- lost. –Copper (II) :two val e- : Cu 2+ –Copper (I) : one val e- : Cu +

16. Anions Non-metal atoms that gain electrons become anions. Examples: Bromine and Nitrogen bromide ionnitride ion You need to MEMORIZE the common charges of the anions to be successful for the rest of the school year…

17. Naming Anions Anions change the ending of the element – unlike cations Stem-ide ion Examples: Chlorine = chloride ion Oxygen = oxide ion Anions –Chlorine (Cl): chloride ion (Cl - ) –Oxygen (O): oxide ion (O 2- )

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

19. Polyatomic ions Polyatomic = many atoms Ions = charged 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).

20. Now on to trends… There is no secret to success this chapter other than to memorize the following trends…

21. Types of Trends Periodic Trends: Trends across a period (row) of the periodic table. Group Trends: Trends down a group (family) of the periodic table

22. Nuclear Charge (+ in nucleus) Periodic: Nuclear charge increases as you go left to right across a period. Group: Nuclear charge increases as you go down a group.

23. Atomic Radii and Size Periodic: Atomic Radii and size decrease as you go L to R across the table. –Same principle energy level (n) –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)

25. 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

26. 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.

27. Electronegativity Electronegativity: The tendency for atoms to attract electrons when they are chemically combined. –Stronger attraction.

28. 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.

29. Summary Decreasing Atomic Radius Increasing Ionization Energy Increasing Electronegativty Increasing Atomic RadiusDecreasing ElectronegativityDecreasing Ionization Energy