1. Chapter 7 Ionic compounds and metals

2. 7.1 Ion Formation Ions are formed when atoms gain or lose valence electrons to achieve a stable octet electron configuration

3. Valence Electrons and Chemical Bonds Recall: –All elements in the same group have the same number of valence electrons and therefore…

4. Valence electrons are involved in the formation of chemical bonds –The force that holds two atoms together –Attraction between positive and negative ions

5. Recall: Dot structures –Show only valence electrons –Carbon: 1s 2 2s 2 2p 2 –Bromine: [Ar] 4s 2 3d 10 4p 5

6. Recall: octet rule –Atoms will gain, lose, or share electrons to acquire the stable electron configuration of a noble gas –Metals – –Nonmetals -

7. Positive Ion formation Name: Formed by: Group 1: Group 2: Group13:

8. Transition metal ions Form cations only Charges may vary in some atoms –Fe can lose 2 or 3 electrons –Fe 2+ or Fe 3+ –Periodic table

9. Negative Ion Formation Name: Formed by: Group 15 Group 16 Group 17

10. 7.2 Ionic Bonds & Ionic Compounds Oppositely charged ions attract each other, forming electrically neutral compounds

11. Formation of an Ionic Bond Ionic bond – the force of attraction that holds oppositely charged ions together

12. Ionic compounds – compounds that contain ionic bonds –Cation + anion –Metal + nonmetal –Also called salts

13. Binary ionic compounds – contain two different elements –NaCl –MgO –K 2 S –CaI 2

14. Compound formation and charge –Ionic compounds are electrically neutral –Total positive charge must = total negative charge –Net charge of all ionic compounds = 0

15. Formation of Sodium Chloride Na: [Ne]3s 1 + Cl: [Ne]3s 2 3p 5 Dot Structure: Na + Cl  [Na] + +[ Cl ] -

16. Na 2 O Na + O 2- Na + Total positive charge =Total negative charge =

17. Al 2 O 3 Al 3+ O 2- Al 3+ O 2- Total positive charge =Total negative charge =

18. How would an ionic compound form from each of the following: Na + N Li + O Sr + F Group 1 + group 15

19. Formulas for ionic compounds Formula unit = the chemical formula for an ionic compound –Simplest ratio of ions involved –Mg 6 Cl 12 –MgCl 2 –Overall charge = 0

20. Monatomic ions – one atom ions –Ex: Oxidation number - the charge of a monatomic ion –Most transition metals have more than one oxidation number

21. What is the oxidation number of the ions in the following compounds? FeO MgCl 2 Cu 3 N Cu 3 N 2 Fe 2 O 3

22. Formulas of binary ionic compounds Cation always written 1 st Anion always written 2 nd Subscripts = # of ions of each element in the compound –No subscript = 1 ion of that element Remember: net charge of compound = 0

23. Write formulas for the following Potassium + oxygen Magnesium + chlorine Aluminum + bromine Barium + phosphorus

24. Crisscross method 1.Write the cation and anion as symbols with their charges. 1.Sn 4+ O 2- 2.Cross the number so they become subscripts 1.Sn 2 O 4 3.Reduce if needed. 1.SnO 2

25. Use the crisscross method to write formulas for the following: Calcium + Oxygen Aluminum + Sulfur

26. Ratio of atoms Na & Cl = –Ratio Mg & Cl = –Ratio Ca & N = –Ratio

27. What is the ratio when an element from group 2 combines with an element from group 17?

28. Formulas for polyatomic ionic compounds Polyatomic ions – ions made up of more than one atom –P. 221 in book –Some on periodic table Acts as an individual ion in a compound Charge applies to entire group of atoms NEVER CHAGE SUBSCRIPTS OF POLYATOMIC IONS

29. Write formulas for the following Calcium + phosphate Sodium + nitrate Calcium + carbonate Aluminum + sulfate

30. Naming ionic compounds Name cation 1 st Change ending of monatomic anion to –ide NaCl KF

31. Name the following NaBr CaCl 2 K 2 O Mg 3 N 2

32. When formula includes a polyatomic ion: Name cation 1 st Name polyatomic ion – DO NOT CHANGE ENDING!!!

33. Name the following KOH Ca(NO 3 ) 2 Na 2 SO 4 (NH 4 ) 3 PO 4

34. Naming with multiple oxidation numbers Previous naming rules apply only to metals in the s block & transition metals with only one oxidation number –Ni 2+, Zn 2+, Al 3+, Ag 1+, Cd 2+ Most transition metals & metals in the p block can have more than one oxidation number

35. Which charge the ion has is indicated by roman numerals –Iron (II) Oxide –Copper (I) Chloride

36. What is the oxidation number of the ions in the following? FeS CuCl Hg 2 O SnS 2

37. Name the following Fe 2 O 3 CuCl 2 CoN Hg 3 P 2

38. Writing formulas from names Figure out cation & charge Figure out anion & charge –Ends in –ide = monatomic –Ends in –ate or –ite = polyatomic Exceptions: hydroxide & cyanide Crisscross to get formula

39. Aluminum Sulfide Sodium Phosphate Calcium Nitrate Barium Oxide

40. Iron (III) Iodide Titanium (III) Oxide Chromium (III) Sulfate Copper (I) Chlorite

41. Properties of ionic compounds Physical structure – ions are packed into a regular repeating pattern

42. Crystal lattice – 3D geometric arrangement of particles in an ionic compound –Formed by the strong attractions among positive and negative ions –Each positive ion is surrounded by negative ions & each negative ion is surrounded by positive ions

44. Physical properties – ionic bonds are very strong, take a lot of energy to be broken apart –High melting point –High boiling point –Hard, rigid, brittle solids

45. More physical properties –Brilliant colors – due to transition metals in crystal lattices –Electrolytes when dissolved or melted Conducts electricity IONIC SOLIDS DO NOT CONDUCT ELECTRICITY

46. Energy and the Ionic Bond Exothermic reactions Endothermic reactions Formation of ionic compounds always releases energy & therefore is…

47. Lattice energy – the amount of energy required to separate 1 mol of ions in an ionic compound –Greater lattice energy = stronger force of attraction

48. Lattice energy is directly related to size of ions bonded –Smaller ions = stronger bond Which is stronger KCl or LiCl?

49. Lattice energy is also related to the charge of the ions –Bond formed from attraction of ions with larger charges = stronger Which is stronger MgO or NaF?

50. 7.4 Metallic Bonds & Properties of Metals Metals form crystal lattices and can be modeled as cations surrounded by a sea of freely moving valence electrons Illustration:

51. Metallic bonds METALS ARE NOT IONIC Share several properties with ionic compounds –Bonding based on attraction of particles with unlike charges –Often form lattices

52. A sea of electrons Metals do not lose or share valence electrons Within lattice, the outer energy levels of metal ions overlap Electron sea model – metal atoms contribute their valence electrons to from a sea of electrons that surround the metal cations

53. Delocalized electrons – electrons that are free to move easily from one atom to the next throughout the metal and are not attached to a particular atom Metallic bond – attraction of a metallic cation for delocalized electrons

54. Properties of metals Melting & boiling points: –Moderately high melting points –High boiling points

55. Malleability, ductility, & durability –Malleable –Ductile –Generally durable

56. Conductivity –Good conductors of both heat & electricity Hardness & strength –More electrons = stronger metal

57. Metal Alloys Alloy – mixture of elements that has metallic properties Stainless steel Brass Cast iron

58. Commercial alloys Brass90% Cu, 10% Zn 10-carat gold42% Au, 20% Ag,38% Cu Stainless steel79% Fe, 18% Cr, 9% Ni Sterling Silver92.5% Ag, 7.5% Cu

59. Substitutional alloys Some of the atoms in the original metallic solid are replaced by other metals of similar atomic size –Sterling silver: Cu replaces some Ag atoms

60. Interstitial alloys Formed when small holes in a metallic crystal are filled with smaller atoms –Steel: holes in iron crystal are filled with carbon atoms to strengthen it