Classification of Matter

1. Classification of Matter

2. Chapter Outline 3.1 Matter Defined 3.7Symbols of the Elements 3.8Metals, Nonmetals andMetalloids 3.2Physical States of Matter 3.3Substances and Mixtures 3.9Compounds 3.4Elements 3.10Elements that Exist asDiatomic Molecules 3.5Distribution of Elements 3.6Names of the Elements 3.11Chemical Formulas

3. Matter Defined

4. Matter is anything that has mass and occupies space. • 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.

5. Example: 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. 3.1

6. Physical States of Matter

7. SOLIDS Shape • Definite - does not change. It is independent of its container. Volume • Definite Particles • Particles are close together. Theycohere rigidly to each other. Compressibility • Very slight–less than liquidsand gases.

8. A solid can be either crystalline or amorphous. Which one it is depends on the internal arrangement of the particles that constitute the solid.

9. LIQUIDS Shape • Not definite - assumes the shape of its container. Volume • Definite Particles • Particles are close together. • Particles are held together by strong attractive forces. • They can move freely throughout the volume of the liquid. Compressibility • Very slight–greater than solids,less than gases.

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

11. 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.

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

13. 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.

16. Substances and Mixtures

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

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

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

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

21. Homogeneous Heterogeneous

22. 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.

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

24. 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.

25. Heterogeneous Mixture A heterogeneous mixture consists of 2 or more phases. 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.

26. Heterogeneous Mixture A heterogeneous mixture consists of 2 or more phases. 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.

27. solid phase1 solid phase2 Heterogeneous Mixture liquid phase

28. 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.

29. 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.

30. 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. 3.2

31. Elements

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

33. 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.

34. 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.

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

36. Distribution of Elements

37. 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%). • Elements are not distributed equally by nature.

38. Distribution of the common elements in nature. 3.3

39. Names of theElements

40. Sources of Element Names 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. German- Color • Bismuth: from the German weisse mass which means white mass. Location • Germanium: discovered in 1866 by a German chemist. Famous- Scientists • Einsteinium: named for Albert Einstein.

41. Symbols of the Elements

42. A symbol stands for • the element itself • one atom of the element • a particular quantity of the element

43. Rules governing symbols of the elements are: • Symbols have either one or two letters. • If one letter is used it is capitalized. H hydrogen C carbon • If two letters are used, only the first is capitalized. Ne neon Ba barium

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

46. Metals, Nonmetals and Metalloids

47. Metals

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

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

50. Examples of Metals gold iron lead