1. Chapters 15/6 Ionic Bonding 15.1 Objectives – Use the periodic table to infer the number of valence electrons in an atom and draw its electron dot (Lewis dot) structure. – Describe formation of cations from metals and anions from nonmetals  California Standards  1d. Students know how to use the periodic table to determine the number of electrons available for bonding.  2e. Students know how to draw Lewis dot structures.  California Standards  1d. Students know how to use the periodic table to determine the number of electrons available for bonding.  2e. Students know how to draw Lewis dot structures.

2. Valence Electrons: ELECTRONS AVAILABLE FOR BONDING (the red ones)

3. Valence Electrons Valence electrons are electrons in the outmost shell (energy level). They are the electrons available for bonding. The number of valence electrons largely determines the chemical properties of that element. For Groups 1A-7A, the number 1-7 is the number of valence electrons for that atom. Group 0 is an exception – you can think of it as group 8A because all the noble gases (except He) have 8 valence electrons.

4. Group 1 (alkali metals) have 1 valence electron

5. Group 2 (alkaline earth metals) have 2 valence electrons

6. Group 13 elements have 3 valence electrons

7. Group 14 elements have 4 valence electrons

8. Group 15 elements have 5 valence electrons

9. Group 16 elements have 6 valence electrons

10. Group 17 (halogens) have 7 valence electrons

11. Group 18 (Noble gases) have 8 valence electrons, except helium, which has only 2

12. Transition metals (“d” block) have 1 or 2 valence e-. Why?

13. Lanthanides and actinides (“f” block) have 1 or 2 valence electrons Lanthanides and actinides (“f” block) have 1 or 2 valence electrons

14. Valence Electrons Valence electrons are usually the only e- used to bond to other atoms. – Therefore you usually only show the valence e- in electron dot structures. – Electron dot structures are diagrams that show valence e- as dots.

15. Generic Dot Notation An atom’s valence electrons can be represented by electron dot (AKA Lewis dot) notations. 1 valence e - X 2 valence e - X 3 valence e - X 4 valence e - X 5 valence e - X 6 valence e - X 7 valence e - X 8 valence e - X ?

16. Dot Notations – Period 2 Lewis dot notations for the valence electrons of the elements of Period 2. lithium Liberyllium Beboron Bcarbon C nitrogen Noxygen Ofluorine Fneon Ne

17. Electron Dot Structures Note how you draw two dots per side x four sides = 8 dots maximum. Note how each side gets one before any side gets two. See how the number of dots is the same for each element within a group (column).

18. Octet Rule The Octet Rule was created by Gilbert Lewis in 1916. That’s why these diagrams are sometimes called Lewis dot structures. In forming compounds, atoms tend to achieve the e- configuration of a noble gas, 8 valence e-. An octet is a set of 8. Each noble gas (except He) has 8 valence electrons in their highest principle energy level, and the general configuration is ns 2 np 6 (like 2s 2 2p 6 or 3s 2 3p 6 )

19. Metallic vs. Nonmetallic Elements Atoms of the metallic elements (including column 1A and 2A) tend to lose their outer shell valence e- so they can have a complete octet at the next energy level down. Atoms of nonmetallic elements tend to gain e- (steal e-) or share e- with another nonmetallic element to achieve their complete octet. There are exceptions but the octet rule usually applies to most atoms in compounds.

20. Cations and Anions If an atom loses a valence e- = cation If an atom gains a valence e- = anion Metals create cations because they start with 1 to 3 e- and usually get all of the valence e- stolen so they can get down to a full lower level octet. Example: Sodium loses 1 e- Before Na 1s 2 2s 2 2p 6 3s 1 After Na + 1s 2 2s 2 2p 6 – (note the 8 e- in the n=2 shell) Like Ne 1s 2 2s 2 2p 6 – (Neon has 8 e- in the n=2 shell) The change is written as follows: – Na· Na + + e -

21. Cations Cations of group 1A alkali metals +1 Cations of group 2A alkali metals +2 ·Mg· Mg 2+ + 2e - For transition metals, the charges on the cations may vary. Example: Fe has two: iron(II) or Fe 2+ iron(III) or Fe 3+ Some atoms formed by transition metals do not have noble-gas electron configurations and are therefore exceptions to the octet rule.

22. Exceptions Example: Ag Silver 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 4d 10 (oddball) Silver would have to lose 11 electrons to get down to noble gas Krypton’s configuration. To gain enough e- to get to Xenon’s configuration, it would have to gain 7 electrons. Neither one is likely. But if Ag loses its one 5s 1 electron, then it has an outer shell with 18 e- (the 4 shell), which is a full shell, and relatively favorable. Therefore Ag always forms the Ag + cation.

23. Exceptions Other elements that behave similarly are Copper(I) – Cu +, Gold(I) – Au +,Cadmium Cd 2+ Mercury(II) – Hg 2+ All have pseudo-noble gas configurations. Note: The column Cu and Ag and Au is in is labeled column 1B in your book on page 392. Cd is column 2B. Can you now see why?

24. Anions Anions are atoms or groups with a negative charge (extra electrons). Atoms of nonmetallic elements have relatively full valence shells and are looking to steal e- to make their shells full. Cl 1s 2 2s 2 2p 6 3s 2 3p 5 neutral atom Cl - 1s 2 2s 2 2p 6 3s 2 3p 6 anion Ar 1s 2 2s 2 2p 6 3s 2 3p 6 now Cl - has Ar config.

25. Anions The ions that are produced when atoms of chlorine and other halogens gain electrons are called halide ions. Here are some common anions:

26. Section 15.2 Ionic Bonding California Standards  Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.  Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction.

27. Ionic vs. Covalent Bonds Bonds: Forces that hold groups of atoms together and make them function as a unit.  Ionic bonds – transfer of electrons  Covalent bonds – sharing of electrons (this will be Ch. 16) (this will be Ch. 16)

28. Ionic Bonding Na: 1s 2 2s 2 2p 6 3s 1 now Na + 1s 2 2s 2 2p 6 Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 now Cl - 1s 2 2s 2 2p 6 3s 2 3p 6

29. Aluminum has three valence e- to steal, and the Bromine atoms would each like to steal one e-. So the Aluminum atom gives up three electrons and the Bromine atoms each receive one.

30. Examples of Ionic Compounds Mg 2+ Cl 2 1- Magnesium chloride: Magnesium loses two electrons and each chlorine gains one electron Al 2 3+ S 3 2- Aluminum sulfide: Each aluminum loses two electrons (six total) and each sulfur gains two electrons (six total)

32. Recall that anions end in –ide.

33. Sodium Chloride crystal lattice Ionic compounds form solid crystals at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions. All salts are ionic compounds and form crystals.

34. NaCl – Which is a cation? Anion? Why do they get bigger/smaller?

35. Question: What is a “formula unit”?

36. Properties of Ionic Compounds

37. Two K atoms lose 1 e- each => One O atom gains 2 e-

38. 3 Mg atoms lose x 2 e- each => 2 N atoms gain 3 e- each

39. Ch. 6 – Ionic Naming The Laws of Definite and Multiple Proportions The law of Definite Proportions states that in samples of any chemical compound, the masses of the elements are always in the same proportions. The law of Multiple Proportions states that whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.

40. Ions of Representative Elements Add – ide to anion name

41. Ions of metallic elements

43. Specific list of polyatomic ions you are accountable to memorize for the test -1 Ions-2 Ions-3 Ions NameFormulaNameFormulaNameFormula AcetateC 2 H 3 O 2 -1 SulfiteSO 3 -2 PhosphatePO 4 -3 HydroxideOH -1 SulfateSO 4 -2 NitrateNO 3 -1 CarbonateCO 3 -2 NitriteNO 2 -1 Bicarbonate (hydrogen carbonate) HCO 3 -1 + 1 Ions AmmoniumNH 4 + Also know diatomic molecules: I 2 Br 2 Cl 2 F 2 O 2 N 2 H 2 and also H 2 O, NH 3 (ammonia), and CH 4 (methane)

45. Ionic Bonding Practice The best way to learn ionic naming is a lot of practice. The sheet in your packet called “Extra Lecture Notes on Ionic Naming of Compounds” is the next thing we need to do to get that practice.

46. 15.3 Metallic Bonding Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

47. Properties of Metals  Metals are good conductors of heat and electricity  Metals are malleable, they can be hammered or forced into shapes  Metals are ductile – they can be drawn into wires  Metals have high tensile strength  Metals have luster

48. Metallic Bonding  Metallic bonding is the chemical bonding that results from the attraction between metal cations and the surrounding sea of electrons  Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move freely throughout the metal  Valence electrons do not belong to any one atom

49. Metallic Bonding Strong forces of attraction are responsible for the high melting point of most metals.

50. Crystalline Structure in Metals Body centered cubic - – Every atom has 8 neighbors – Na, K, Fe, Cr, W Face centered cubic – Every atom has 12 neighbors – Cu, Ag, Au, Al, Pb Hexagonal close packed – Every atom has 12 neighbors – Packed in hexagons – Mg, Zn, Cd

51. Packing in Metals Hexagonal close packed. Model: Packing uniform, hard spheres to best use available space. This is called Hexagonal close packed. Each atom has 12 nearest neighbors. Does this look like how oranges are stacked at the grocery store? It should.

52. Alloys Sterling silverAg 92.5%, Cu 7.5% Stainless steelFe 80.6%, Cr 18.0%, C 0.4%, Ni 1.0% Cast ironFe 96%, C 4% BrassCu and Zn (% varies) BronzeCu and Sn (% varies)

53. Metal Alloys  Substitutional Alloy  Substitutional Alloy: some metal atoms replaced by others of similar size.

54. Metal Alloys  Interstitial Alloy:  Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.