Ionic Bonding

1. Ionic Bonding Valence Electrons, Lewis Dot Structures, and Electronegativity

4. Valence Electrons • valence electrons – the outermost electrons.

5. 6 protons = C = carbon Outer electrons = 4 e–available for bonding Inner electrons = 2 e– not available for bonding

6. 4 electrons in valence shell

7. Lewis Dot Structures H • Lewis dot structures are a convenient way to show how many valence electrons an atom has. • Example: Draw the Lewis dot structure for hydrogen.

8. Lewis Dot Structures • dots = number of valence electrons. • The maximum number of dots is 8. • Look for the number at the top of the column (e.g. 5A). • Exception: Helium only has 2 dots. He Ne

9. More Lewis Structure Practice • Draw the Lewis structure for oxygen. • Draw the Lewis structure for magnesium. • Draw the Lewis structure for chlorine. Mg O Cl

10. Even More Lewis Structure Practice K • Draw the Lewis structure for carbon. • Draw the Lewis structure for potassium. • Draw the Lewis structure for phosphorus. P C

11. Electronegativity • electronegativity – how much an atom wants to keep hold of its electrons. • ionization energy – the energy required to remove an electron from an atom.

13. Lower electronegativity Greater electronegativity

14. Metals

15. Nonmetals

16. Role Models: The Noble Gases • An atom’s electrons are at their most stable when they reorganize their electrons to more closely resemble the electron configuration of a noble gas. • All atoms want to have stable electron configurations.

17. Role Models: The Noble Gases • All atoms wish their electrons were like the noble gases’ electrons.

18. Example: Beryllium and Oxygen

19. The 2 valence e– Now beryllium’s matches that of helium 4 protons = Be = beryllium

20. Now oxygen’s matches that of neon The 6 valence e– 8 protons = O = oxygen

21. Same Thing, but in Lewis Dot Structure Be O

23. Semi-metals or metalloids Nonmetals Metals

24. Three General Bonding Types • Metal with Nonmetal - form ionic compounds • Metal with Metal - form metallic compounds • Nonmetal with Nonmetal - form covalent compounds

25. Metal with Nonmetal Bonding • Ionic compounds – the metal gives all of its valence e– to the nonmetal. • Known as – Salts, ions

27. Ions • ion – an atom that gained or lostelectrons to become more like a noble gas. + Metals lose electrons to become positively charged ions. We call them cations (cat-ions) the “t” looks like a “+”. [e.g. 2A  +2] – Nonmetals gain electrons to become negatively charged ions. We call them anions (an-ions) “n” for negative “–”. [e.g. (8 – 6A) × -1 -2]

28. Writing the Charges • Write out the ion that sodium forms. Na+ • Write out the ion that chlorine forms. Cl– • Write out the ion that magnesium forms. Mg2+ • Write out the ion that oxygen forms. O2–

29. So where do the electrons go? • Usually atoms that become cations give their electrons to anions. • Now both the cations and anions resemble noble gases, however now both have net charges. O 2– 2+ Be

30. Basic Electrical Charge Laws + and– : Attract (pull together)

31. Naming (aka nomenclature) • Metals keep their names unchanged. (e.g. sodium, aluminum, calcium) • Transition metals have their charge shown as roman numerals in parenthesis after the name. Fe2+ iron (II) Cu1+ copper (I) Fe3+ iron (III) Cu2+ copper (II) • Nonmetals have the last one or two syllables of their names altered with an –ide ending. fluorine  fluoride nitrogen  nitride chlorine  chloride oxygen  oxide

32. Naming (aka nomenclature) • Metals keep their names unchanged. (e.g. sodium, aluminum, calcium) • If there are more than one possible charge for a metal (the transition metals), the charge will be specified in roman numerals after the name. Fe2+ iron (II) Cu1+ copper (I) Fe3+ iron (III) Cu2+ copper (II)

33. Naming Continued • Nonmetals have the last one or two syllables of their names altered with an –ide ending. • Examples: carbon  carbide fluorine  fluoride nitrogen  nitride chlorine  chloride oxygen  oxide bromine  bromide sulfur  sulfide iodine  iodide phosphorus  phosphide

34. Naming Continued • Now put the metal and nonmetal ion names together and you get the name for the ionic compound. • Examples: LiF  lithium fluoride NaCl  sodium chloride KBr  potassium bromide MgS  magnesium sulfide CuI  copper (I) iodide CuO  copper (II) oxide FeN  iron (III) nitride

35. Sodium Chloride – NaCl + 6 Na+ Na+ Na+ Na+ Na+ Na+ Cl– Cl– Cl– Cl– Cl– Cl– – 6 0

36. Magnesium Chloride + 12 Cl– Cl– Cl– Cl– Cl– Cl– Cl– Cl– Cl– Cl– Cl– Cl– – 12 – 6 Mg2+ Mg2+ Mg2+ Mg2+ Mg2+ Mg2+ + 6 0

37. Magnesium Chloride 6 x Cl– 12 x Mg6Cl12 Mg2+ 61 12 2 MgCl2 =

38. Magnesium Chloride Empirical Formula Cl– Cl– MgCl2 Mg2+ Formula Unit (f.u.) – the smallest amount of an ionic compound that still has the same ratio of ions as in the formula.

39. Iron (III) Oxide Fe3+ O2– 2 x (+3) = +6 3 x (–2) = –6 criss-cross 0 Fe2O3

40. Sodium Cloride Na+ Cl– criss-cross Na1Cl1 NaCl

41. Magnesium Sulfide Mg2+ S2– criss-cross Mg2S2 21 2 1 = MgS

42. Sodium Oxide Na+ O2– criss-cross Na2O

43. REMEMBER Fe3+ O2– Top right corner: Charge Fe2O3 Bottom right corner: How many atoms/ions

44. Polyatomic Ions • Poly – many • Atomic – having to do with atoms • Polyatomic ions – ions made from multiple atoms • List on p.257

45. Polyatomic Ions (List on p.257) Ion nameFormula Nitrate NO3– Nitrite NO2– Permanganate MnO4– Peroxide O22– Phosphate PO43– Sulfate SO42– Sulfite SO32– Thiosulfate S2O32– Ion nameFormula Acetate CH3COO– Ammonium NH4+ Carbonate CO32– Chromate CrO42– Cyanide CN– Dichromate Cr2O72– Hydroxide OH–

46. Magnesium Nitrate Mg2+ NO3– criss-cross Mg(NO3)2 1 x Mg 2 x N 6 x O

47. REMEMBER Mg2+ NO3– Top right corner: Charge Mg(NO3)2 Bottom right corner: How many atoms/ions

48. Sodium Chloride – NaCl Na+ Na+ Na+ Na+ Na+ Na+ Cl– Cl– Cl– Cl– Cl– Cl–

49. Sodium Chloride – NaCl

50. Sodium Chloride – NaCl