2. Chapter 5 Molecules and Compounds 2006, Prentice Hall Sucrose molecule (sugar), contains C, H, and O atoms. However, the properties of sucrose are very different from those of C, H, and O alone. The properties of a compound are, in general, different from the properties of the elements that compose it. molecular formula C 12 H 22 O 11

3. 2 CHAPTER OUTLINE  Law of Constant Composition  Octet Rule and Ions  Ionic Charges  Definitions of Molecules and Compounds  Types of Compounds  Binary Ionic Compounds (Type I and Type II)  Polyatomic Ions  Binary Molecular Compounds  Naming Acids and Oxyacids  Formula Mass

4. 3 LAW OF DEFINITE COMPOSITION  In a pure compound, the elements are always present in the same definite proportion by mass. Mass of sample Mass of NMass of H Sample 11.840 g1.513 g0.327 g Sample 22.000 g1.644 g0.356 g % N in Sample 1 = % N in Sample 2 =  Based on this law, the mass of an element can be determined from its mass percent in a compound.

5. 4 Barium iodide, BaI 2, contains 35.1% barium by mass. How many grams of barium does an 8.50 g sample of barium iodide contain? Example 1: 35.1 100 2.98 3 significant figures

6. 5 When 12.66 g of calcium are heated in air, 17.73 g of calcium oxide is formed. What is the percent of oxygen in this compound? Example 2: 3 significant figures Mass of oxygen =17.73 g – 12.66 g = 5.07 g Percent oxygen =

7. 6 OCTET RULE & IONS  Most elements, except noble gases, combine to form compounds. Compounds are the result of the formation of chemical bonds between two or more different elements.  In the formation of a chemical bond, atoms lose, gain or share valence electrons to complete their outer shell and attain a noble gas configuration.  This tendency of atoms to have eight electrons in their outer shell is known as the octet rule.

8. 7 5.1 FORMATION OF IONS  An ion (charged particle) can be produced when an atom gains or loses one or more electrons. A cation (+ ion) is formed when a neutral atom loses an electron Metals form cations

9. 8 5.1 FORMATION OF IONS Non-metals form anions  An anion (- ion) is formed when a neutral atom gains an electron.

10. 9 IONIC CHARGES  The ionic charge of an ion is dependent on the number of electrons lost or gained to attain a noble gas configuration.  For most main group elements, the ionic charges can be determined from their group number, as shown below:

11. molecule - consist of at least two different atoms in a definite arrangement held together by very strong chemical bonds. compound is a substance consisting of two or more elements in fixed and definite proportions Definition of Molecule and Compound

12. Formulas Describe Compounds a compound is a distinct substance that is composed of atoms of two or more elements a compound can be described by the number and type of each atom in the simplest unit of the compound –molecules or ions 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

13. Formulas Describe Compounds 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

14. Tro's Introductory Chemistry, Chapter 5 13 Order of Elements in a Formula metals written first –NaCl nonmetals written in order from Table 5.1 –CO 2 –there are occasional exceptions for historical or informational reasons H 2 O fits rule, but does NaOH? Table 5.1 Order of Listing Nonmetals in Chemical Formulas CPNHSIBrClOF

15. Types of Chemical Formulas An empirical formula gives the relative number of atoms of each element in a compound. A molecular formula gives the actual number of atoms of each element in a molecule of the compound. For example, the molecular formula for hydrogen peroxide is H 2 O 2, and its empirical formula is HO. The molecular formula is always a whole number multiple of the empirical formula. For many compounds, such as H 2 O, the molecular formula is the same as the empirical formula. A structural formula uses lines to represent chemical bonds and shows how the atoms in a molecule are connected to each other.

16. Comparison of Formulas and Models for Methane, CH 4 The molecular formula of methane indicates that methane has 1 carbon atom and 4 hydrogen atoms. The structural formula shows how the atoms are connected: each hydrogen atom is bonded to the central carbon atom. The ball-and-stick model and the space-filling model illustrate the geometry of the molecule: how the atoms are arranged in three dimensions.

17. Tro's Introductory Chemistry, Chapter 5 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

18. 17 Molecular Elements Certain elements occur as 2 atom molecules Rule of 7’s –there are 7 common diatomic elements –find the element with atomic number 7, N –make a figure 7 by going over to Group 7A, then down –don’t forget to include H 2 H2H2 Cl 2 Br 2 I2I2 7 VIIA N 2 O 2 F 2

19. Molecular Compounds two or more nonmetals smallest unit is a molecule

20. Ionic Compounds metals + nonmetals no individual molecule, instead these have 3- dimensional array of cations and anions made of formula units (NaCl)

21. 20 IONIC COMPOUNDS  Ionic compounds contain ionic bonds, which occur when electrons are transferred between two atoms.  After bonding, each atom achieves a complete shell (noble gas configuration).  Ionic bonds occur between metals and non-metals. MetalNonmetal

22. 21 IONIC COMPOUNDS  Atoms that lose electrons (metals) form positive ions (cations).  Atoms that gain electrons (non-metals) form negative ions (anions).  The smallest particles of ionic compounds are ions (not atoms). Cation Anion

23. 22 IONIC CHARGES AND FORMULAS  The formula of an ionic compound indicates the number and kinds of ions that make up the ionic compound.  The sum of the ionic charges in the formula is always zero, which indicates that the total number of positive charges is equal to the total number of negative charges.  For example, the +1 charge on the sodium ion is cancelled by the –1 charge on the chloride ion, to form a net zero charge. loses 1 e  gains 1 e  (1+) + (1  ) = 0

24. 23 IONIC CHARGES AND FORMULAS  When charges between the two ions do not balance, subscripts are used to balance the charges.  For example, since each magnesium loses 2 electrons, and each chloride gains one electron, 2 chlorides are needed to balance the charge of the magnesium ion.  Therefore magnesium chloride is written as MgCl 2. (2+) + 2(1  ) = 0 loses 2 e  Each gains 1 e  MgCl 2

25. 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 = atomic element = ionic compound = molecular element = molecular compound = atomic element

26. 25 TYPES OF COMPOUNDS  Compounds are pure substances that contain 2 or more elements combined in a definite proportion by mass.  Compounds can be classified as one of two types: Metals and non-metals Two non- metals

27. Nomenclature of Compounds Common Names – Are Exceptions (like nicknames) H 2 O = water, steam, ice NH 3 = ammonia CH 4 = methane NaCl = table salt C 12 H 22 O 11 = table sugar

28. Metal Cations Type I –metals whose ions can only have one possible charge IA, IIA, (Al, Ga, In) –determine charge by position on the Periodic Table IA = +1, IIA = +2, (Al, Ga, In = +3) Type II –metals whose ions can have more than one possible charge –determine charge by charge on anion How do you know a metal cation is Type II? its not Type I !!!

29. 28 BINARY IONIC COMPOUNDS (TYPE I)  Binary compounds contain only two elements.  Ionic compounds are formed by combination of a metal and a non-metal.  Type I ions are those cations that form only one ion.  In these compounds, charges of the cations must equal the charges of the anions since the net charge is zero.

30. 29 BINARY IONIC COMPOUNDS (TYPE I)  Subscripts are used to balance the charges between cations and anions. sodium bromide Na + Br  +1 -1 = 0 NaBr No subscripts needed potassium sulfide K+K+ S2S2 +1 -2  0 K2+K2+ +2 -2 = 0 K2SK2S

31. 30 Write formulas for the following ionic compounds: calcium chloride Ca 2+ Cl  +2 -1  0 CaCl 2 sodium sulfide Na + S2S2 +1 -2  0 +2 -2 = 0 Na 2 S Example 1: Cl 2  +2 -2 = 0 Na 2 +

32. 31 Type I Binary Ionic Compounds Contain Metal Cation + Nonmetal Anion Metal listed first in formula & name 1.name metal cation first, name nonmetal anion second 2.cation name is the metal name 3.nonmetal anion named by changing the ending on the nonmetal name to -ide

33. 32 BINARY IONIC COMPOUNDS (TYPE I) MgCl 2 magnesium chloride NaI sodium iodide AlF 3 aluminum fluoride

34. 33 Name the following ionic compounds: barium chloride Na 3 Psodium phosphide BaCl 2 Example 2:

35. 34 BINARY IONIC COMPOUNDS (TYPE II)  Type II ions are those cations that form more than one ion.  When naming compounds formed from these ions, include the ionic charge as Roman numeral, in parentheses, after the metal’s name.  This method of nomenclature is called the “stock” system.

36. 35 NAMING AND WRITING IONIC FORMULAS  When writing ionic formula, knowing the charge of the ions are important since the net charge on the compound must be zero.  Some elements produce only one ion (Type I) while others produce two or more ions (Type II).  Differentiating between type I and II ions is important, since the naming system is different for each. Shown below are the common ions of each type: Type I Type II

37. 36 1.name metal cation first, name nonmetal anion second 2.metal cation name is the metal name followed by a Roman Numeral in parentheses to indicate its charge (only difference from type II!) –determine charge from anion charge 3.nonmetal anion named by changing the ending on the nonmetal name to -ide BINARY IONIC COMPOUNDS (TYPE II)

38. 37 BINARY IONIC COMPOUNDS (TYPE II) FeCl 2 ? -1 ? -2 = 0 +2 -1 iron (II) chloride FeCl 3 ? -1 ? -3 = 0 +3 -1 iron (III) chloride

39. 38 BINARY IONIC COMPOUNDS (TYPE II) Cu 2 O ? -2 2? -2 = 0 +1 -2 copper (I) oxide Roman numeral DOES NOT represent the subscript CuO ? -2 ? -2 = 0 +2 -2 copper (II) oxide

40. 39 BINARY IONIC COMPOUNDS (TYPE II)  Type II cations can also be named by an older method (classical).  In this system, cations with the higher charge end in –ic, while cations with the lower charge end in –ous.  In this system, some cations are named based on their Latin roots.

41. 40 CLASSICAL SYSTEM (DERIVED FROM LATIN) gold (aurum for Aurora the Roman goddess of the dawn ) silver (argentum for 'bright' ) copper (cuprum for 'Cyprus' where the Romans first obtained copper ) tin (stannum for alloys containing lead ) lead (plumbum for 'lead' ) mercury (hydrargyrum for 'liquid silver' or quick silver ) antimony (stibium for 'not alone' ) iron (ferrum for 'firmness‘ ) potassium (kalium via the Arabic qali for alkali ) sodium (natrium for soda )

42. 41 BINARY IONIC COMPOUNDS (TYPE II) FeCl 2 +2 -1 ferrous chloride FeCl 3 +3 -1 ferric chloride Lower charge Higher charge

43. 42 BINARY IONIC COMPOUNDS (TYPE II) Cu 2 O +1 -2 cuprous oxide CuO +2 -2 cupric oxide Lower charge Higher charge

44. 43 SnCl 2 ? -1 ? -2 = 0 +2 -1 tin (II) chloride Example 1: Name each of the following compounds using the stock and classical nomenclature system: Stock Classical Lower charge stannous chloride

45. 44 Cu 2 S ? -2 2? -2 = 0 +1 -2 copper (I) sulfide Example 1: Name each of the following compounds using the stock and classical nomenclature system: Stock Classical Lower charge cuprous sulfide

46. 45 Sn Example 2: Write formulas for each of the following compounds: Tin (II) bromideBr +2 -1 +2 -1  0 Br 2 SnBr 2 +2 -2 = 0

47. 46 Sn Example 2: Write formulas for each of the following compounds: Stannic oxideO +4 -2 +4 -2  0 SnO 2 +4 - 4 = 0 O2O2

48. Examples What, if anything, should go into the parenthesis LiCl = lithium ( ) chloride AlCl 3 = aluminum ( ) chloride PbO = lead ( ) oxide PbO 2 = lead ( ) oxide Mn 2 O 3 = manganese ( ) oxide X X II IV III

49. 48 POLYATOMIC IONS  Some ionic compounds contain polyatomic ions, an ion composed of several atoms bound together.

50. Patterns for Polyatomic Ions 1.elements in the same column form similar polyatomic ions –same number of O’s and same charge ClO 3 - = chlorate  BrO 3 - = bromate 2. if the polyatomic ion starts with H, the name adds hydrogen- prefix before name and add 1 to the charge CO 3 2- = carbonate  HCO 3 -1 = hydrogencarbonate

51. Periodic Pattern of Polyatomic Ions -ate groups BO 3 -3 NO 3 SiO 3 -2 PO 4 -3 SO 4 -2 ClO 3 AsO 4 -3 SeO 4 -2 BrO 3 TeO 4 -2 IO 3 CO 3 -2 IIIAIVA VA VIA VIIA

52. Patterns for Polyatomic Ions -ate ion –chlorate = ClO 3 -1 -ate ion + 1 O  same charge, per- prefix –perchlorate = ClO 4 -1 -ate ion – 1 O  same charge, -ite suffix –chlorite = ClO 2 -1 -ate ion – 2 O  same charge, hypo- prefix, -ite suffix –hypochlorite = ClO -1

53. © 2012 Pearson Education, Inc. Naming Ionic Compounds Containing a Polyatomic Ion Some examples of more than two ions in the series. ClO − hypochloriteBrO − hypobromiteIO − hypoiodite ClO 2 − chloriteBrO 2 − bromiteIO 2 − iodite ClO 3 − chlorateBrO 3 − bromateIO 3 − iodate ClO 4 − perchlorateBrO 4 − perbromateIO 4 − periodate When naming these ions in the homework, be sure to include the word ion in your answer, as in “perchlorate ion.”

54. 53 POLYATOMIC COMPOUNDS  When writing formulas for polyatomic compounds, treat the polyatomic ion as one group. potassium nitrate K+K+ NO 3 – +1 -1 = 0 KNO 3 calcium hydroxide Ca 2+ OH – +2 -1  0 (OH – ) 2 Ca(OH) 2 +2 -2 = 0

55. 54 POLYATOMIC COMPOUNDS ammonium acetate NH 4 + C2H3O2–C2H3O2– +1 -1 = 0 NH 4 C 2 H 3 O 2 sodium sulfate Na + SO 4 2– +1 -2  0 Na 2 + Na 2 SO 4 +2 -2 = 0

56. 55 POLYATOMIC COMPOUNDS copper (II) nitrate Cu 2+ NO 3 – +2 -1  0 Cu(NO 3 ) 2 (NO 3 – ) 2 +2 -2 = 0 Type II Roman numeral represents charge of ion Alternate name =cupric nitrate

57. 56 Mg(OH) 2 magnesium hydroxide NaCN Example 2: Name the following polyatomic compounds: Type I ion (Does not require roman numeral) sodium cyanide Fe 2 (SO 4 ) 3 Type II ion (Requires roman numeral) +3  2 iron (III) sulfate ferric sulfate

58. Subclasses Compounds containing a metal and a nonmetal = binary ionic –Type I and II Compounds containing a polyatomic ion = ionic with polyatomic ion

59. 58 BINARY MOLECULAR COMPOUNDS  Molecular compounds are formed by combination of 2 or more non-metals.  The smallest particles of molecular compounds are molecules.  These compounds are named similar to ionic compounds, with the second element named based on its root and suffix “-ide”.

60. 59 BINARY MOLECULAR COMPOUNDS  Greek prefixes are used to indicate the number of atoms in these compounds. NumberPrefixNumberPrefix 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10deca-

61. 60 Examples: Name the following binary molecular compounds: CS 2 carbon disulfide First atom uses a prefix only when more than one atom is present Second atom always uses a prefix indicates 1 carbon atom indicates 2 sulfur atoms

62. 61 Examples: Name the following binary molecular compounds: IF 7 iodine heptafluoride indicates 1 iodine atom indicates 7 fluorine atoms

63. 62 SUMMARY OF BINARY NOMENCLATURE Ionic Molecular Type I Type II

64. Naming Acids Contain H + cation and anion –in aqueous solution Binary acids have H + cation and nonmetal anion Oxyacids have H + cation and polyatomic anion

65. Naming Acids Binary Acids = hydro prefix + stem of the name of the nonmetal + ic suffix Example: HCl  hydrochloric acid Oxyacids –if polyatomic ion ends in –ate = name of polyatomic ion with –ic suffix Example: HNO 3  nitric acid –if polyatomic ion ends in –ite = name of polyatomic ion with –ous suffix Example: H 2 SO 3  sulfurous acid base name of oxyanion + -icacid base name of oxyanion + -ous acid

66. 65 ACIDS  Acids are molecular compounds that form ions when dissolved in water. Binary Acids  Formulas are written similar to binary ionic compounds, assigning a +1 charge to hydrogen. HCl +1 -1 H2SH2S +1 -2

67. 66 NAMING BINARY ACIDS  When naming the acids, use hydro- prefix, followed by the name of the non-metal with an –ic ending, followed with the word acid. HCl H2SH2S acidhydrochloric hydrosulfuric acid HFhydrofluoric acid

68. 67 POLYATOMIC ACIDS  Several polyatomic acids are important in the study of chemistry, and their names must be learned.  These acids and the polyatomic ions that form from their ionization are as follows:  Most of these are oxyacids

69. 68 POLYATOMIC ACIDS sulfuric acidH 2 SO 4 SO 4 2  sulfate phosphoric acidH 3 PO 4 PO 4 3  phosphate nitric acidHNO 3 NO 3  nitrate

70. 69 POLYATOMIC ACIDS carbonic acidH 2 CO 3 CO 3 2  carbonate HCO 3  bicarbonate acetic acid HC 2 H 3 O 2 C 2 H 3 O 2  acetate

71. The names of acids containing oxyanions ending with -ite take this form: The names of acids containing oxyanions ending with -ate take this form: base name of oxyanion + ous base name of oxyanion + ic

72. Names of Some Common Oxyacids and Their Oxyanions

73. Formula Mass: The Mass of a Molecule or Formula Unit also known as molecular mass For any compound, the formula mass is the sum of the atomic masses of all the atoms in its chemical formula: Mass of 1 molecule of H 2 O = 2(1.01 amu H) + 16.00 amu O = 18.02 amu

74. 73 THE END