1. Mr. Matthew Totaro Legacy High School Honors Chemistry Chemical Formulas and Chemical Nomenclature 1

2. 2 Molecular View of Elements and Compounds

3. 3 Classifying Materials Atomic elements = Elements whose particles are single atoms. Molecular elements = Elements whose particles are multi-atom molecules. Molecular compounds = Compounds whose particles are molecules made of only nonmetals. Ionic compounds = Compounds whose particles are cations and anions.

4. Atomic Elements 4 Atomic Elements = elements whose smallest units are single atoms Most elements are atomic elements Elemental Mercury

5. 5 Molecular Elements Certain elements occur as diatomic molecules. 7 diatomic elements H O F Br I N Cl H2H2 Cl 2 Br 2 I2I2 7 7A N 2 O 2 F 2

6. 6 Molecular Compounds Two or more nonmetals. Smallest unit is a molecule.

7. 7 Ionic Compounds Metals + nonmetals. No individual molecule units, instead have a 3- dimensional array of cations and anions made of formula units.

8. 8 Classify Each of the Following as Either an Atomic Element, Molecular Element, Molecular Compound, or Ionic Compound. Aluminum, Al. Aluminum chloride, AlCl 3. Chlorine, Cl 2. Acetone, C 3 H 6 O. Carbon monoxide, CO. Cobalt, Co.

9. 9 Classify Each of the Following as Either an Atomic Element, Molecular Element, Molecular Compound, or Ionic Compound, Continued. Aluminum, Al = Atomic element. Aluminum chloride, AlCl 3 = Ionic compound. Chlorine, Cl 2 = Molecular element. Acetone, C 3 H 6 O = Molecular compound. Carbon monoxide, CO = Molecular compound. Cobalt, Co = Atomic element.

10. Law of Constant Composition 10

11. Law of Constant Composition 11

12. 12 Law of Constant Composition All pure substances have constant composition. All samples of a pure substance contain the same elements in the same mass ratios. Mixtures have variable composition.

13. 13 Compounds Display Constant Composition If we decompose 18.0 g of water by electrolysis, we find 16.0 grams of oxygen to every 2.00 grams of hydrogen. Water has a constant mass ratio of oxygen to hydrogen of 8.0.

14. Example—Show that Two Samples of Carbon Dioxide, CO 2 Are Consistent with the Law of Constant Composition. Sample 1: 4.8 g O, 1.8 g C; Sample 2: 17.1 g O, 6.4 g C 14

15. 15 Why Do Compounds Show Constant Composition? The smallest piece of a compound is called a molecule. Every molecule of a compound has the same number and type of atoms. Since all the molecules of a compound are identical, every sample will have the same ratio of the elements. Since all molecules of a compound are identical, every sample of the compound will have the same properties.

16. Chemical Formulas 16

17. 17 Formulas Describe Compounds A chemical formula describes the compound by describing the number and type of each atom in the simplest unit of the compound. Each element is represented by its letter symbol. The number of atoms of each element is written to the right of the element as a subscript. If there is only one atom, the 1 subscript is not written. C2H2O4C2H2O4

18. Types of Chemical Formulas 18

19. 19 Formulas Describe Compounds, Continued Water = H 2 O  two atoms of hydrogen and 1 atom of oxygen Table sugar = C 12 H 22 O 11  12 atoms of C, 22 atoms of H and 11 atoms O

20. 20 Order of Elements in a Formula Metals are written first. NaCl Nonmetals are written in order from Table 5.1. CO 2 There are occasional exceptions for historical or informational reasons.  H 2 O, but NaOH. Table 5.1 Order of Listing Nonmetals in Chemical Formulas CPNHSIBrClOF

21. Example: Write a formula for ionic compound that forms between calcium and oxygen

22. Example: Write the formula of a compound made from aluminum ions and oxygen ions

23. Practice — What are the formulas for compounds made from the following ions? K + with N 3 − Ca 2+ with Br − Al 3+ with S 2 −

24. 24 Ion Charge and the Periodic Table The charge on an ion can often be determined from an elements position on the periodic table. Metals are always positive ions, nonmetals are negative ions. For many main group metals, the cation charge = the group number. For nonmetals, the anion charge = the group number – 8.

25. Li + Na + K+K+ Rb + Cs + Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+ Al 3+ Ga 3+ In 3+ O2O2 S2S2 Se 2  Te 2  FF Cl  Br  II N3N3 P3P3 As 3  1A 2A3A7A6A5A

26. 26 Formulas containing Polyatomic Ions

27. Compounds Containing Polyatomic Ions Polyatomic ions are single ions that contain more than one atom Often end in –ate or -ite Name and charge of polyatomic ion do not change Nitrate Polyatomic Ion (NO 3 - )

28. Some Common Polyatomic Ions

29. 29 Polyatomic Ions The polyatomic ions are attracted to opposite ions by ionic bonds. Atoms in the polyatomic ion are held together by covalent bonds.

30. Example – Write the formula for the ionic compound containing sodium and phosphate

31. Practice — What are the formulas for compounds made from the following ions? aluminum ion with a sulfate ion barium ion with hydrogen carbonate ion

32. Molecules with Polyatomic Ions, Continued Mg(NO 3 ) 2 Compound called magnesium nitrate. CaSO 4 Compound called calcium sulfate. Subscript indicating two NO 3 groups. No subscript indicating one SO 4 group. Implied “1” subscript on nitrogen, total 2 N. Implied “1” subscript on sulfur, total 1 S. Stated “3” subscript on oxygen, total 6 O. Stated “4” subscript on oxygen, total 4 O.

33. 33 Practice—Determine the Total Number of Atoms or Ions in One Formula Unit of Each of the Following. Mg(C 2 H 3 O 2 ) 2 (Hg 2 ) 3 (PO 4 ) 2 1 Mg + 4 C + 6 H + 4 O = 15 6 Hg + 2 P + 8 O = 16

34. Chemical Nomenclature 34 Mg(NO 3 ) 2 Compound called Magnesium Nitrate

35. 35 Naming Ions - Cations Metals form cations For each positive charge, the ion has 1 less electron than the neutral atom Na atom = 11 p + and 11 e -, Na + ion = 11 p + and 10 e - Ca atom = 20 p + and 20 e -, Ca 2+ ion = 20 p + and 18 e - Cations are named the same as the metal sodiumNa  Na + + 1e - sodium ion calciumCa  Ca 2+ + 2e - calcium ion

36. 36 Naming Ions - Anions Nonmetals form anions For each negative charge, the ion has 1 more electron than the neutral atom F = 9 p + and 9 e -, F ─ = 9 p + and 10 e - O = 8 p + and 8 e -, O 2 ─ = 8 p + and 10 e - Anions are named by changing the ending of the name to -ide fluorineF + 1e -  F ─ fluoride ion oxygenO + 2e -  O 2─ oxide ion

37. 37 s block Metals + a Nonmetal Binary Nomenclature

38. 38 Example—Naming Binary Ionic, Type I Metal, CsF

39. 39 Practice—Name the Following Compounds. KCl MgBr 2 Al 2 S 3

40. 40 Practice—Name the Following Compounds, Continued. KClpotassium chloride. MgBr 2 magnesium bromide. Al 2 S 3 aluminum sulfide.

41. 41 d block Metals + a Nonmetal Binary Nomenclature

42. - typically have more than one charge Transition Metal Cations Copper (I)Cu + Silver (I)Ag + Gold (I)Au + mercury (I)Hg 2 +2 zincZn +2 copper (II)Cu +2 cadmium (II)Cd +2 mercury (II)Hg +2 tin (II)Sn +2 manganese (II)Mn +2 iron (II)Fe +2 cobalt (II)Co +2 nickel (II)Ni +2 lead (II)Pb +2 chromium (III)Cr +3 gold (III)Au +3 bismuth (III)Bi +3 colbalt (III)Co +3 iron (III)Fe +3 lead (IV)Pb +4 tin (IV)Sn +3

43. 43 Determining the Name of a Type II Compound— Au 2 S 3

44. 44 Example—Naming Binary Ionic, Type II Metal, CuCl

45. 45 Practice ─ Name the Following Compounds. TiCl 4 PbBr 2 Fe 2 S 3

46. 46 Practice ─ Name the Following Compounds, Continued. TiCl 4 Titanium(IV) chloride. PbBr 2 Lead(II) bromide. Fe 2 S 3 Iron(III) sulfide. Cl = 4(−1) = −4 Ti = +4 = 1(4+) Br = 2(−1) = −2 Pb = +2 = 1(2+) S = 3(−2) = −6 Pb = +6 = 2(3+)

47. 47 Example—Naming Ionic with Polyatomic Ion, Na 2 SO 4

48. 48 Example—Naming Ionic with Polyatomic Ion, Fe(NO 3 ) 3

49. 49 Practice ─ Name the Following 1.NH 4 Cl 2.Ca(C 2 H 3 O 2 ) 2 3.Cu(NO 3 ) 2

50. 50 Practice ─ Name the Following, Continued 1.NH 4 ClAmmonium chloride. 2.Ca(C 2 H 3 O 2 ) 2 Calcium acetate. 3.Cu(NO 3 ) 2 Copper(II) nitrate.

51. 51 Nonmetal (p block)+ a Nonmetal (p block) Binary Nomenclature

52. 52 Subscript—Prefixes 1 = mono- Not used on first nonmetal. 2 = di- 3 = tri- 4 = tetra- 5 = penta- 6 = hexa- 7 = hepta- 8 = octa- Drop last “a” if name begins with vowel.

53. 53 Example—Naming Binary Molecular, BF 3

54. 54 Example—Naming Binary Molecular, N 2 O 4

55. 55 Practice ─ Name the Following NO 2 PCl 5 I 2 F 7

56. 56 Practice ─ Name the Following Continued NO 2 Nitrogen dioxide PCl 5 Phosphorus pentachloride I 2 F 7 Diiodine heptafluoride.

57. 57 Common Names—Exceptions H 2 O = Water, steam, ice. NH 3 = Ammonia. CH 4 = Methane. NaCl = Table salt. C 12 H 22 O 11 = Table sugar.

58. 58 Acids Acids are molecular compounds that form H + when dissolved in water. To indicate the compound is dissolved in water, (aq) is written after the formula.  Not named as acid if not dissolved in water. Sour taste. Dissolve many metals. Like Zn, Fe, Mg, but not Au, Ag, Pt. Formula generally starts with H. E.g., HCl, H 2 SO 4.

59. 59 Acids, Continued Contain H +1 cation and anion. In aqueous solution. Binary acids have H +1 cation and nonmetal anion. Oxyacids have H +1 cation and polyatomic anion.

60. 60 Naming Binary Acids Write a hydro- prefix. Follow with the nonmetal name. Change ending on nonmetal name to –ic. Write the word acid at the end of the name.

61. 61 Example—Naming Binary Acids, HCl

62. 62 Naming Oxyacids If polyatomic ion name ends in –ate, then change ending to –ic suffix. If polyatomic ion name ends in –ite, then change ending to –ous suffix. Write word acid at end of all names.

63. 63 Example—Naming Oxyacids, H 2 SO 4

64. 64 Example—Naming Oxyacids, H 2 SO 3

65. 65 Practice ─ Name the Following 1.H 2 S 2.HClO 3 3.HNO 2

66. 66 Practice ─ Name the Following Continued 1.H 2 S hydrosulfuric acid 2.HClO 3 chloric acid 3.HNO 2 nitrous acid

67. 67 Writing Formulas for Acids When name ends in acid, formulas starts with H. Write formulas as if ionic, even though it is molecular. Hydro- prefix means it is binary acid, no prefix means it is an oxyacid. For an oxyacid, if ending is –ic, polyatomic ion ends in –ate; if ending is –ous, polyatomic ion ends in –ous.

68. 68 Example—Binary Acids, Hydrosulfuric Acid

69. 69 Example—Oxyacids, Carbonic Acid

70. 70 Example—Oxyacids, Sulfurous Acid

71. 71 Practice—What Are the Formulas for the Following Acids? 1.Chlorous acid 2.Phosphoric acid 3.Hydrobromic acid

72. 72 Example—Writing Formula for a Binary Ionic Compound Containing Variable Charge Metal, Manganese(IV) Sulfide

73. 73 Example—Writing Formula for an Ionic Compound Containing Polyatomic Ion, Iron(III) phosphate

74. 74 Practice—What Are the Formulas for Compounds Made from the Following Ions? 1.Copper(II) Nitride 2.Iron(III) Bromide

75. 75 Practice—What Are the Formulas for Compounds Made from the Following Ions? 1.Aluminum Sulfate 2.Chromium(II) Hydrogen Carbonate