1. 1 Matter and Energy Chapter 3

2. Why does soda fizz when you open the bottle? Why is the sun hot? When iron rusts, what’s happening?

3. 3 Universe Classified Matter is the part of the universe that has mass and volume -There are three states of matter  Solid, liquid, and gas Chemistry is the study of matter –The properties of different types of matter –The way matter behaves when influenced by other matter and/or energy

4. 4 Elements and Compounds a pure substance homogeneous – same composition throughout Contains only one type of atom Can not be broken down into simpler substances by ordinary chemical reactions (not a nuclear reaction) Element

5. Chemical Symbols of Elements System started by Jons Berzelius (Sweden, 1779-1848) 90-91 naturally occurring elements One or two first letters of name of the element. Many elements names have roots from: Latin, Greek, mythology, geography, names of scientists.

6. Examples: Americium, Am Einsteinium, Es Bromine, Br Helium, He Lead(Plumbum), Pb Niobium, Nb Iron (Ferrum), Fe Mendelevium, Md Examples of other elements: O 2, H 2, I 2

7. Compound – chemical combination of two or more elements Have two properties in common with elements: 1. pure substance 2. Homogeneous – same chemical composition at all times Have two properties that differ from elements: 1. two or more elements that are chemically combined, in a definite ratio 2. Compounds can be broken down by chemical reactions using energy: a) decomposition - uses heat b) electrolysis - uses electricity 7

8. > 10 million compounds elements are represented by symbols compounds are represented by chemical formulas chemical formula 1. symbol --> tells which elements are present in compounds 2. subscript (little # lower right) --> tells the number of atoms of each element 8

9. When the elements sodium and chlorine combine chemically to form sodium chloride, there is a change in composition and a change in properties. Properties of Compounds Sodium chloride (commonly known as table salt) is a white solid. Distinguishing Elements and Compounds Compounds have different properties from their individual elements.

10. Breaking down Compounds Breaking down NaCl Sodium is a soft gray metal. Distinguishing Elements and Compounds

11. Breaking down NaCl Breaking down Compounds Chlorine is a pale yellow poisonous gas. Distinguishing Elements and Compounds

12. Classification of Matter (by composition)

13. 13 Classification of Matter Mixtures can be classified as: Homogeneous Mixtures Heterogeneous Mixtures

14. Homogeneous mixture = uniform throughout, appears to be one thing –Also called solutions –Examples: olive oil, salt water, lemonade, coffee, air 14 The substances in the olive oil are evenly distributed throughout the mixture

15. Example: Stainless Steel A homogeneous mixture of: -Iron (Fe) -Chromium (Cr) -Nickel (Ni)

16. Heterogeneous mixture = non-uniform, contains regions with different properties than other regions - Examples: oil and vinegar, salad, chicken soup 16

17. 17 Pure Substances vs. Mixtures Pure Substances –All samples have the same physical and chemical properties –Constant Composition  all samples have the same composition –Homogeneous –Separate into components based on chemical properties Mixtures –Different samples may show different properties –Variable composition –Homogeneous or Heterogeneous –Separate into components based on physical properties All mixtures are made of pure substances

18. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–18 Figure 3.4: When table salt is stirred into water (left), a homogeneous mixture called a solution forms (right).

19. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 19 Identity Each of the following as a Pure Substance, Homogeneous Mixture or Heterogeneous Mixture ¬Gasoline ­A stream with gravel on the bottom ®Copper metal

20. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 20 Identity Each of the following as a Pure Substance, Homogeneous Mixture or Heterogeneous Mixture ¬Gasoline –a homogenous mixture ­A stream with gravel on the bottom –a heterogeneous mixture ®Copper metal –A pure substance (all elements are pure substances)

21. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 21 Separation of Mixtures Separate mixtures based on different physical properties of the components EvaporationVolatility ChromatographyAdherence to a Surface FiltrationState of Matter (solid/liquid/gas) DistillationBoiling Point TechniqueDifferent Physical Property

22. Example: Separate iron filings from sulfur using a magnet.

23. Filtration: separates a solid from a liquid in a heterogeneous mixture

24. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–24 Figure 3.6: Distillation of a solution consisting of salt dissolved in water. - separate dissolved solids from a liquid in a homogeneous mixture -uses boiling and condensation.

25. Distillation of Crude Oil (Refining) Crude Oil is a mixture of Hydrocarbons

26. Distillation of Crude Oil

27. Paper Chromatography Chromatography separates mixtures of substances into their components. They all have: a stationary phase (a solid, or a liquid supported on a solid) and a mobile phase (a liquid or a gas). Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–27

28. Objectives Observe and explain the difference between states of matter. Explain the difference between physical and chemical changes. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 28

29. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.. What you observe when you look at a particular sample of matter is its properties. –Is a solid shiny or dull? –Does a liquid flow quickly or slowly? –Is a gas odorless, or does it have a smell? Describing Matter

30. Some Criteria for the Classification of Matter Composition (elements and ✔ compounds) State (solid, liquid, gas) Properties

31. 3 States of Matter Solid Liquid Gas Copyright©2004 by Houghton Mifflin Company. All rights reserved. 31

32. States of Matter

33. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–33

34. Figure 3.1: Liquid water takes the shape of its container.

35. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–35 Figure 3.11: In ice, the water molecules vibrate randomly about their positions in the solid. Their motions are represented by arrows.

36. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–36 Figure 3.12: Equal masses of hot water and cold water separated by a thin metal wall in an insulated box.

37. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–37 Figure 3.13: The H 2 O molecules in hot water have much greater random motions than the H 2 O molecules in cold water.

38. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–38 Figure 3.14: The water samples now have the same temperature (50°C) and have the same random motions.

39. Phase Changes Meltingsolid  liquid Condensationgas  liquid Freezingliquid  solid Deposition gas  solid Evaporationliquid  gas Sublimationsolid  gas *Boiling: Evaporation occurring beneath the liquid’s surface.

40. Gallium metal has such a low melting point (30°C) that it melts from the heat of a hand.

43. Video - https://www.youtube.com/watch?v=wln6WSv- cro

44. Properties of Matter Copyright©2004 by Houghton Mifflin Company. All rights reserved. 44

45. Extensive Properties mass - a measure of the amount of matter the object contains. volume of an object is a measure of the space occupied by the object. –The volume of a basketball is greater than the volume of a golf ball. 45

46. Who has a greater volume?

47. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.. Mass and volume are both examples of extensive properties. –An extensive property is a property that depends on the amount of matter in a sample. Examples: mass and volume Extensive Properties

48. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.. –An intensive property is a property that depends on the type of matter in a sample, not the amount of matter.  Examples include: - Hardness of an object -Color -Softness -Boiling point -Absorbency -Odor Intensive Properties

49. Is changing phase a physical or chemical change?

50. 50 Properties of Matter Physical Properties are the characteristics of matter that can be changed without changing its composition –Characteristics that are directly observable –Examples: Color, odor, hardness, density, melting point, boiling point, state, solubility.

51. SubstanceStateColor Melting Point (C°) Boiling Point (C°) Density (g/cm 3 ) OxygenO2O2 GasColorless-218-1830.0014 MercuryHgLiquidSilvery- white -3935713.5 BromineBr 2 LiquidRed-brown-7593.12 WaterH2OH2OLiquidColorless01001.00 Sodium Chloride NaClSolidWhite80114132.17 Example: Physical Properties

52. Properties of Matter Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter - Examples: burning, decompose, reactivity, corrode, tarnish, explode, ferment Copyright©2004 by Houghton Mifflin Company. All rights reserved. 52

53. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 53 Classify Each of the following as Physical or Chemical Properties ¬The boiling point of ethyl alcohol is 78°C. ­Diamond is very hard. ®Sugar ferments to form ethyl alcohol.

54. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 54 Classify Each of the following as Physical or Chemical Properties ¬The boiling point of ethyl alcohol is 78°C. –Physical property – describes inherent characteristic of alcohol – boiling point ­Diamond is very hard. –Physical property – describes inherent characteristic of diamond – hardness ®Sugar ferments to form ethyl alcohol. –Chemical property – describes behavior of sugar – forming a new substance (ethyl alcohol)

55. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 55 Changes in Matter Physical Changes are changes to matter that do not result in a change in the composition that make that substance –State Changes – boiling, melting, condensing –Breaking, splitting, grinding, cutting Chemical Changes involve a change in the composition of the substance –Produce a new substance –Chemical reaction –Reactants  Products

56. Indications of a Chemical Reaction Color change Solid forms (precipitate) Gas bubbles Odor Temperature change Fizzing Copyright©2004 by Houghton Mifflin Company. All rights reserved. 56

57. Formation of a Precipitate Cu(OH) 2 Precipitate

58. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 58 Classify Each of the following as Physical or Chemical Changes ¬Iron metal is melted. ­Iron combines with oxygen to form rust. ®Sugar ferments to form ethyl alcohol.

59. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 59 Classify Each of the following as Physical or Chemical Changes ¬Iron is melted. –Physical change – describes a state change, but the material is still iron ­Iron combines with oxygen to form rust.. –Chemical change – describes how iron and oxygen react to make a new substance, rust ®Sugar ferments to form ethyl alcohol. –Chemical change – describes how sugar forms a new substance (ethyl alcohol)

60. Oxygen combines with the chemicals in wood to produce flames. Is a physical or chemical change taking place? Source: Jim Pickerell/ Stone/Getty Images

61. Burning of Methane CH 4 +2O 2  CO 2 + 2H 2 O l_______________l l_________________l l l reactants products

62. Burning of Methane CH 4 + 2O 2  CO 2 + 2H 2 O

63. The Law of Conservation of Mass (Antoine Lavoisier) In any chemical or physical change, mass is neither created or destroyed Mass is CONSTANT

64. 64 Energy and Energy Changes Capacity to do work –chemical, mechanical, thermal, electrical, radiant, sound, nuclear Energy may affect matter –e.g. raise its temperature, eventually causing a state change –All physical changes and chemical changes involve energy changes

65. 65 Heat Heat: a flow of energy due to a temperature difference 1.Exothermic = A process that results in the evolution of heat. Example: when a match is struck, it is an exothermic process because energy is produced as heat. 2.Endothermic = A process that absorbs energy. Example: melting ice to form liquid water is an endothermic process.

66. A burning match releases energy. Source: ElektraVision/ PictureQuest

67. Video Endothermic & Exothermic Reactions Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–67

68. 68 Units of Energy One calorie is the amount of energy needed to raise the temperature of one gram of water by 1°C joule (J) –4.184 J = 1 cal In nutrition, calories are capitalized –1 Cal = 1,000 cal(1 kcal = 1,000 cal)

69. 69 Example - Converting Calories to Joules Convert 60.1 cal to joules

70. Convert 0.1 kcal to Joules Copyright©2004 by Houghton Mifflin Company. All rights reserved. 70

71. Video Chemmatters Calorie vs calorie https://www.youtube.com/watch?v=G0O87 gWv-Xkhttps://www.youtube.com/watch?v=G0O87 gWv-Xk Copyright©2004 by Houghton Mifflin Company. All rights reserved. 71

72. 72 Energy and the Temperature of Matter The amount the temperature of an object increases depends on the amount of heat added (Q). –If you double the added heat energy the temperature will increase twice as much. The amount the temperature of an object increases depends on its mass (m) –If you double the mass it will take twice as much heat energy to raise the temperature the same amount.

73. 73 Specific Heat Capacity Specific Heat (s) is the amount of energy required to raise the temperature of one gram of a substance by one Celsius degree

74. Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–74

75. Specific Heat Capacity Calculations Amount of Heat = Specific Heat x Mass x Temperature Change Where  T = (Final temperature – initial temperature) = (T f - T i ) Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. Q = s x m x  T

76. 76 Example 1 – Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C Mass = 7.40 g Temperature Change = 46.0°C – 29.0°C = 17.0°C Specific Heat of Water = 4.184 Q = s x m x  T

77. 77 Example 2 – A 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold? Table 3.2 lists the specific heat of gold as 0.13 Therefore the metal cannot be pure gold.

78. Calorimetry Calorimetry is the study of the heat flow that accompanies physical and/or chemical changes. The apparatus used to measure heat is a calorimeter. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 78

79. “ loses ” heat Calorimetry m = 75 g T = 25 o C SYSTEM Surroundings m = 30 g T = 100 o C Ag H2OH2O A hot piece of silver (Ag) at 100 o C is added to water that has an initial temperature of 25 o C. Silver will lose heat, and water will gain heat.

80. 240 g of water (initially at 20 o C) are mixed with an unknown mass of iron (initially at 500 o C). When thermal equilibrium is reached, the system has a final temperature of 42 o C. Find the mass of the iron. Calorimetry Problems 2 question #5 Fe T = 500 o C mass = ? grams T = 20 o C mass = 240 g -LOSE heat = GAIN heat - [(S, Fe ) (mass) (  T)] = (S, H2O ) (mass) (  T) - [(0.45 J/g o C) (X g) (42 o C - 500 o C)] = (4.184 J/g o C) (240 g) (42 o C - 20 o C)] Drop Units: - [(0.45 J/g o C) (X) (-458 o C)] = (4.184J/g o C) (240 g) (22 o C) (206.1 J/g) X = 22091 J X = 107.2 g Fe

81. A 97 g sample of gold at 785 o C is dropped into 323 g of water, which has an initial temperature of 15 o C. If the final temperature of the system has a temperature of 22.1 o C, what is the the specific heat of gold? Au T = 785 o C mass = 97 g T = 15 o C mass = 323 g LOSE heat = GAIN heat - [(S,Au) (mass) (  T)] = (S,H 2 O) (mass) (  T) - [(x) (97 g) (22.1 - 785 o C)] = (4.184 J/g o C) (323 g) (22.1 - 15 o C)] Drop Units: - [(x)(97 g) (-762.9 o C)] = (1351.4 J/ o C) (7.1 o C)] -(-74001.3 g o C) x = 9594.9 J x = 9594.9 J 74001.3 g o C x = S Au = 0.13 J/g o C

82. A Coffee Cup Calorimeter Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 302 Thermometer Styrofoam cover Styrofoam cups Stirrer