1. 1 Classification of Matter Chapter 3 Hein and Arena Eugene Passer Chemistry Department Bronx Community College © John Wiley and Sons, Inc Eugene Passer Chemistry Department Bronx Community College © John Wiley and Sons, Inc Version 1.0

2. 2 Chapter Outline 3.1 Matter DefinedMatter Defined 3.2 Physical States of MatterPhysical States of Matter 3.3 Substances and MixturesSubstances and Mixtures 3.7 Symbols of the ElementsSymbols of the Elements 3.8 Metals, Nonmetals and MetalloidsMetals, Nonmetals andMetalloids 3.4 ElementsElements 3.5 Distribution of ElementsDistribution of Elements 3.6 Names of the ElementsNames of the Elements3.11 Chemical FormulasChemical Formulas 3.9 CompoundsCompounds 3.10 Elements that Exist as Diatomic MoleculesElements that Exist asDiatomic Molecules

3. 3 Matter Defined

4. 4 Matter can be invisible. –Air is matter, but it cannot be seen. Matter appears to be continuous and unbroken. –Matter is actually discontinuous. It is made up of tiny particles call atoms. Matter is anything that has mass and occupies space.

5. 5 3.1 An apparently empty test tube is submerged, mouth downward in water. Only a small volume of water rises into the tube, which is actually filled with invisible matter–air.

6. 6 Physical States of Matter

7. 7 Shape Definite - does not change. It is independent of its container. Volume Definite ParticlesParticles are close together. They cohere rigidly to each other. SOLIDS CompressibilityVery slight–less than liquids and gases.

8. 8 Solid Amorphous Solid Particles lack a regular internal arrangement Glass, plastics, gels Crystalline Solid Particles exist in regular, repeating three-dimensional geometric patterns. Diamond, metals, salts A solid can be either crystalline or amorphous. Which one it is depends on the internal arrangement of the particles that constitute the solid.amorphous

9. 9 Amorphous: without shape or form.

10. 10 ShapeNot definite - assumes the shape of its container. Volume Definite ParticlesParticles are close together. Particles are held together by strong attractive forces. They can move freely throughout the volume of the liquid. LIQUIDS CompressibilityVery slight–greater than solids, less than gases.

11. 11 GASES Shape No fixed shape. Volume Indefinite. ParticlesParticles are far apart compared to liquids and solids. Particles move independently of each other.

12. 12 GASES Compressibility The actual volume of the gas particles is small compared to the volume of space occupied by the gas. –Because of this a gas can be compressed into a very small volume or expanded almost indefinitely.

13. 13 Attractive forces are strongest in a solid. –These give a solid rigidity. ATTRACTIVE FORCES Solid LiquidAttractive forces are weaker in liquids than in solids. –They are sufficiently strong so that a liquid has a definite volume.

14. 14 ATTRACTIVE FORCES Gas Attractive forces in a gas are extremely weak. Particles in the gaseous state have enough energy to overcome the weak attractive forces that hold them together in liquids or solids. –Because of this the gas particles move almost independently of each other.

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16. 16

17. 17 Substances and Mixtures

18. 18 Matter refers to all of the materials that make up the universe.

19. 19 Substance A particular kind of matter that has a fixed composition and distinct properties. Examples ammonia, water, and oxygen.

20. 20 Homogeneous Matter Matter that is uniform in appearance and with uniform properties throughout. Examples ice, soda, pure gold

21. 21 Heterogeneous Matter Matter with two or more physically distinct phases present. Examples ice and water, wood, blood

22. 22 Homogeneous Heterogeneous

23. 23 Phase A homogenous part of a system separated from other parts by physical boundaries. Examples In an ice water mixture ice is the solid phase and water is the liquid phase.

24. 24 Mixture Matter containing 2 or more substances that are present in variable amounts. Mixtures are variable in composition. They can be homogeneous or heterogeneous.

25. 25 Homogeneous Mixture (Solution) A homogeneous mixture of 2 or more substances. It has one phase. Example Sugar and water. Before the sugar and water are mixed each is a separate phase. After mixing the sugar is evenly dispersed throughout the volume of the water.

26. 26 Example Sugar and fine white sand. The amount of sugar relative to sand can be varied. The sugar and sand each retain their own properties. Heterogeneous Mixture A heterogeneous mixture consists of 2 or more phases.

27. 27 Example Iron (II) sulfide (FeS) is 63.5% Fe and 36.5% S by mass. Mixing iron and sulfur in these proportions does not form iron (II) sulfide. Two phases are present: a sulfur phase and an iron phase. If the mixture is heated strongly a chemical reaction occurs and iron (II) sulfide is formed. FeS is a compound of iron and sulfur and has none of the properties of iron or sulfur. Heterogeneous Mixture A heterogeneous mixture consists of 2 or more phases.

28. 28 solid phase 2 liquid phase solid phase 1 Heterogeneous Mixture

29. 29 Mixture of iron and sulfur Compound of iron and sulfur FormulaHas no definite formula: consists of Fe and S. FeS CompositionContains Fe and S in any proportion by mass. 63.5% Fe and 36.5% S by mass. SeparationFe and S can be separated by physical means. Fe and S can be separated only by chemical change.

30. 30 Heterogeneous Mixture of One Substance A pure substance can exist as different phases in a heterogeneous system. Example Ice floating in water consists of two phases and one substance. Ice is one phase, and water is the other phase. The substance in both cases is the same.

31. 31 System The body of matter under consideration. Examples In an ice water mixture ice is the solid phase and water is the liquid phase. The system is the ice and water together.

32. 32 3.2 Classification of matter: A pure substance is always homogeneous in composition, whereas a mixture always contains two or more substances and may be either homogeneous or heterogeneous.

33. 33

34. 34Elements

35. 35 An element is a fundamental or elementary substance that cannot be broken down into simpler substances by chemical means.

36. 36 All known substances on Earth and probably the universe are formed by combinations of more than 100 elements. Each element has a number. –Beginning with hydrogen, as 1 the elements are numbered in order of increasing complexity.

37. 37 Most substances can be decomposed into two or more simpler substances. –Water can be decomposed into hydrogen and oxygen. –Table salt can be decomposed into sodium and chlorine. An element cannot be decomposed into a simpler substance.

38. 38 ATOM The smallest particle of an element that can exist. The smallest unit of an element that can enter into a chemical reaction.

39. 39 Distribution of Elements

40. 40 Elements are not distributed equally by nature. –Oxygen is the most abundant element in the human body (65%). –Oxygen is the most abundant element in the crust of the earth (49.2%).. In the universe the most abundant element is hydrogen (91%) and the second most abundant element is helium (8.75%).

41. 41 3.3 Distribution of the common elements in nature.

42. 42 Names of the Elements

43. 43 Sources of Element Names Famous- Scientists Einsteinium: named for Albert Einstein. LocationGermanium: discovered in 1866 by a German chemist. German- Color Bismuth: from the German weisse mass which means white mass. Greek- Color Iodine: from the Greek iodes meaning violet. Latin- Property Fluorine: from the Latin fluere meaning to flow. The fluorine containing ore fluorospar is low melting.

44. 44 Symbols of the Elements

45. 45 A symbol stands for –the element itself –one atom of the element –a particular quantity of the element

46. 46 Ne neon Rules governing symbols of the elements are: 1.Symbols have either one or two letters. 2.If one letter is used it is capitalized. 3.If two letters are used, only the first is capitalized. H hydrogenC carbon Ba barium

47. 47 Most symbols start with the same letter as the element.A number of symbols appear to have no connection with the element. These symbols have carried over from the earlier names of the elements (usually Latin).

48. 48

49. 49 Metals, Nonmetals and Metalloids

50. 50 Metals

51. 51 Most elements are metals Metals are solid at room temperature. –Mercury is an exception. At room temperature it is a liquid. Metals have high luster (they are shiny). Metals are good conductors of heat and electricity. Metals are malleable (they can be rolled or hammered into sheets). physical properties of metals

52. 52 Most elements are metals Metals are ductile (they can be drawn into wires). Most metals have a high melting point. Metals have high densities

53. 53 Examples of Metals goldiron lead

54. 54 Many metals readily combine with nonmetals to form ionic compounds. –They can combine with sulfur. Metals have little tendency to combine with each other to form compounds. chlorine. –In nature, minerals are formed by combinations of the more reactive metals with other elements. Chemical Properties of Metals oxygen.

55. 55 –A few of the less reactive metals such as copper, silver and gold are found in the free state. –Metals can mix with each other to form alloys.  Brass is a mixture of copper and zinc.  Bronze is a mixture of copper and tin.  Steel is a mixture of carbon and iron. Chemical Properties of Metals

56. 56 Nonmetals

57. 57 Have relatively low melting points Have low densities. Poor conductors of heat and electricity At room temperature carbon, phosphorous, sulfur, selenium, and iodine are solids. Physical Properties of Nonmetals Lack luster (they are dull)

58. 58 Solid sulfur selenium Physical State at Room Temperature phosphorouscarbon iodine

59. 59 liquid Physical State at Room Temperature bromine

60. 60 gas Physical State at Room Temperature helium, neon, argon, krypton, xenon, radon nitrogen, oxygen fluorine, chlorine

61. 61 Metalloids

62. 62 Metalloids have properties that are intermediate between metals and nonmetals

63. 63 The Metalloids 1.boron 2.silicon 3.germanium 4.arsenic 5.antimony 6.tellurium 7.polonium

64. 64 Metals are found to the left of the metalloids Nonmetals are found to the right of the metalloids.

65. 65 Compounds

66. 66 A compound is a distinct substance that contains two or more elements combined in a definite proportion by weight.

67. 67 Compounds can be decomposed chemically into simpler substances– that is, into simpler compounds or elements. Elements cannot be decomposed into simpler substances. Atoms of the elements that constitute a compound are always present in simple whole number ratios. They are never present as fractional parts.

68. 68 There are two types of compounds: molecular and ionic.

69. 69 Molecules

70. 70 A molecule is the smallest uncharged individual unit of a compound formed by the union of two or more atoms.

71. 71 A water molecule consists of two hydrogen atoms and one oxygen atom. 3.5 If it is subdivided the water molecule will be destroyed and hydrogen and oxygen will be formed.

72. 72 Ionic Compounds

73. 73 An ion is a positively or negatively charged atom or group of atoms.

74. 74 A cation is a positively charged ion. 3.5

75. 75 An anion is a negatively charged ion. 3.5

76. 76 Ionic compounds are held together by attractive forces between positively and negatively charged ions.

77. 77 Sodium chloride is a colorless crystalline ionic substance. Sodium Chloride It is 39.3% sodium and 60.7% chlorine by mass. The solid does not conduct electricity. Passing an electric current through the molten salt produces solid sodium and gaseous chlorine.

78. 78 The ultimate particles of sodium chloride are positively charged sodium ions and negatively charged chloride ions. Sodium Chloride 3.5

79. 79 The crystalline structure of sodium chloride is held together by the attractive forces between the positive sodium ions and the negative chloride ions. Sodium Chloride

80. 80 Ionic Compound Formulas

81. 81 Sodium chloride and other ionic compounds consist of large aggregates of cations and anions. The actual chemical formulas of ionic compounds express the smallest whole number ratio that exists between the cations and the anions.

82. 82 The formula NaCl does not mean that a molecule of NaCl exists. The formula NaCl means that the ratio of sodium to chlorine in a sodium chloride crystal is one to one.

83. 83 The ratio of Na + to Cl - is 1:1

84. 84 3.4 Compounds can be classified as molecular or ionic. Ionic compounds are held together by attractive forces between their positive and negative charges. Molecular compounds are held together by covalent bonds.

85. 85 Elements That Exist as Diatomic Molecules

86. 86 A diatomic molecule contains exactly two atoms of the same or different elements.

87. 87

88. 88 Occurrence of Diatomic Molecules HydrogenH Not found in nature. HydrogenH2H2 Found in nature. NitrogenN Not found in nature. NitrogenN2N2 Found in nature.

89. 89 Hydrogen gas is found in volcanoes and it can be prepared in the laboratory. In both cases it is diatomic hydrogen, H 2. Air is about 21% oxygen by volume. Oxygen can also be prepared in the laboratory. In both cases it is diatomic oxygen, O 2.

90. 90 Water has the formula H 2 O. H2H2 O2O2 H2OH2O It does not contain free hydrogen, H 2 or free oxygen, O 2. The H 2 part of H 2 O means that 2 atoms of hydrogen are combined with one atom of oxygen in the water molecule.

91. 91 Chemical Formulas

92. 92 Serve as abbreviations of the names of compounds. CaCl 2 calcium chloride chemical formulas

93. 93 chemical formulas Tell which elements the compound is composed of and how many atoms of each element are present in a formula unit. CaCl 2 calcium chlorine

94. 94 chemical formulas Show the symbols of the atoms of the elements present in a compound. CaCl 2 Ca calcium Cl chlorine

95. 95 chemical formulas Show the ratio of the atoms of the elements present in a compound. CaCl 2

96. 96 Rules for Writing Chemical Formulas

97. 97 When a formula contains one atom of an element, the symbol of that element represents the one atom. The number one (1) is not used as a subscript.

98. 98 NaCl indicates the element sodium (one atom) indicates the element chlorine (one atom)

99. 99 When the formula contains more than one atom of an element, the number of atoms is indicated by a subscript written to the right of the symbol of that atom.

100. 100 indicates 3 H atoms indicates the element phosphorous (P) indicates the element hydrogen (H) indicates the element oxygen (O) H 3 PO 4 indicates 4 O atoms

101. 101 When the formula contains more than one of a group of atoms that occurs as a unit, parentheses are placed around the group, and the number of units of the group is indicated by a subscript placed to the right of the parentheses.

102. 102 Ba 3 (PO 4 ) 2 indicates three Ba atoms indicates the element barium indicates the phosphate group composed of one phosphorous atom and four oxygen atoms

103. 103 Formulas written as H 2 O, H 2 SO 4, Ca(NO 3 ) 2 and C 12 H 22 O 11 show only the number and kind of each atom contained in the compound; they do not show the arrangements of the atoms in the compound or how they are chemically bonded to each other.

104. 104 H2OH2O

105. 105 Mixtures

106. 106