1. C HAPTER 1 The Foundations of Chemistry

2. What is Chemistry?  Physical science that studies the composition, structure, and properties of matter and the changes it undergoes  Includes many different branches of study (focuses on a particular area, they do overlap)  Organic  Inorganic  Physical  Analytical  Biochemistry  Theoretical

3. Chemistry is……  A natural science  A language with its own vocabulary  A way of thinking

4. Click below to watch the Visual Concept. Visual Concept Chapter 1 Chemistry

5. Prediction Experiment Modify Observations Hypothesis Experiment Law Theory (Model) Scientific Method

6. What is Matter?  Matter is anything that takes up space and has mass  Mass is the amount of matter in an object  Mass is resistance to change in motion along a smooth and level surface

7. Matter  Atoms are the building blocks of all matter  An atom is the smallest particle of an element that maintains its chemical identity through all chemical and physical changes. 7

8. Properties of Matter  Physical properties- characteristic that can be observed without changing the identity  Observed with the senses  Changes of state  Density, color solubility  Chemical properties- indicates how a substance reacts with something else  When observed original substance is changed  Rusting or oxidation  Chemical Rxn

9. Properties of Matter

10. Click below to watch the Visual Concept. Visual Concept Chapter 1 Comparing Physical and Chemical Properties

11. Characteristic Properties  Quality of a substance that never changes, used to identify substance  Extensive- depends on amount of matter present  Mass  volume  Intensive- does not depend on amount of matter present  Melting point  Boiling point

12. Click below to watch the Visual Concept. Visual Concept Chapter 1 Comparing Extensive and Intensive Properties

13. Changes in Matter  Physical change- alter form of a substance but NOT identity  Original substance continues to exist  Chemical change- substances combine/break apart to form new substance  Original substance no longer exists

14. Chemical Changes, continued The products are the substances that are formed by the chemical change. reactants product chemical reactionA change in which one or more substances are converted into different substances is called a chemical change or chemical reaction. The reactants are the substances that react in a chemical change. Carbon plus oxygen yields (or forms) carbon dioxide. carbon + oxygen carbon dioxide

15. Evidence of a Chemical Change Chapter 1

16. Click below to watch the Visual Concept. Chapter 1 Chemical Reaction Visual Concept

17. Click below to watch the Visual Concept. Visual Concept Chapter 1 Comparing Chemical and Physical Changes

18. Physical vs. Chemical  Examples:  rusting iron  dissolving in water  burning a log  melting ice  grinding spices chemical physical chemical physical

20. 20 Mixtures, Substances, Compounds, & Elements  Substance  matter in which all samples have identical composition and properties  Elements  substances that cannot be decomposed into simpler substances via chemical reactions  Elemental symbols  found on periodic table

21. Click below to watch the Visual Concept. Visual Concept Chapter 1 Element

22. 22 Mixtures, Substances, Compounds, & Elements  Mixtures  composed of two or more substances can be separated by physical means  homogeneous mixtures  heterogeneous mixtures  Compounds  substances composed of two or more elements in a definite ratio by mass  can be decomposed into the constituent elements by chemical means  Water is a compound that can be decomposed into simpler substances – hydrogen and oxygen

23. Click below to watch the Visual Concept. Visual Concept Chapter 1 Compounds

24. Click below to watch the Visual Concept. Visual Concept Chapter 1 Classification Scheme for Matter

25. States of Matter  Solids  Particles packed tightly together  Have definite volume and definite shape  Particles vibrate  Gases  Change volume very easily  Particles spread apart, filling all space available to them  No definite shape nor volume  Liquids  No shape of its own, takes shape of its container  Has definite volume

26. © 2009, Prentice-Hall, Inc. States of Matter

27. Click below to watch the Visual Concept. Chapter 1 Liquid Section 2 Matter and Its Properties Visual Concept

28. 28 States of Matter  Changes in state require changes in energy.  heating  cooling

29. States of Matter Solid Liquid Gas Definite Volume? YES NO Definite Shape? YES NO Temp. increase Small Expans. Large Expans. Compressible? NO YES

30. 30 A Molecular View Dalton’s Atomic Theory Dalton’s atomic theory summarized the nature of matter as known in 1808 1)An element is composed of extremely small indivisible particles called atoms. 2)All atoms of a given element have identical properties, which differ from those of other elements. 3)Atoms cannot be created, destroyed, or transformed into atoms of another element. 4)Compounds are formed when atoms of different elements combine with each other in small whole- number ratios. 5)The relative numbers and kinds of atoms are constant in a given compound.

31. Natural Laws  Scientific (natural) law  A general statement based the observed behavior of matter to which no exceptions are known.  Law of Conservation of Mass  Law of Conservation of Energy  Law of Conservation of Mass and Energy  Einstein’s Theory of Relativity  E=mc 2 31

32. Number vs. Quantity  Quantity = number + unit UNITS MATTER!!

33. 33 Measurements in Chemistry QuantityUnitSymbol  lengthmeter m  masskilogram kg  timesecond s  currentampere A  temperatureKelvin K  amt. substancemole mol

34. Measurements in Chemistry Metric Prefixes mega-M10 6 deci-d10 -1 centi-c10 -2 milli-m10 -3 PrefixSymbolFactor micro-  10 -6 nano-n10 -9 pico-p10 -12 kilo-k10 3 BASE UNIT---10 0

35. 35 Units of Measurement Definitions  Mass  measure of the quantity of matter in a body  Weight  measure of the gravitational attraction for a body

36. 36 Units of Measurement Common Conversion Factors  Length  1 m = 39.37 inches  2.54 cm = 1 inch  Volume  1 liter = 1.06 qt  1 qt = 0.946 liter  See Table 1-8 for more conversion factors

37. Use of Numbers  Exact numbers  1 dozen = 12 things  Accuracy  how closely measured values agree with the correct value  Precision  how closely individual measurements agree with each other 37

38. Click below to watch the Visual Concept. Visual Concept Accuracy and Precision

39. Significant Figures  Indicate precision of a measurement.  Consists of all the digits known with certainty plus one final digit, which is somewhat uncertain or estimated 2.35 cm

40. Significant Figures Rules  Counting Sig Figs Count all numbers EXCEPT:  Leading zeros -- 0.0025  Trailing zeros without a decimal point -- 2,500

41. Calculating with Significant Figures  Exact Numbers do not limit the # of sig figs in the answer.  Counting numbers: 12 students  Exact conversions: 1 m = 100 cm  “1” in any conversion: 1 in = 2.54 cm

42. Significant Figures, continued Rounding

43. Counting Sig Fig Examples 4. 0.080 3. 5,280 2. 402 1. 23.50 2. 402 3. 5,280 4. 0.080 4 sig figs 3 sig figs 2 sig figs

44. Calculating with Significant Figures  Multiply/Divide - The # with the fewest sig figs determines the # of sig figs in the answer. (13.91g/cm 3 )(23.3cm 3 ) = 324.103g 324 g 4 SF3 SF

45. Calculating with Significant Figures  Add/Subtract - The # with the lowest decimal value determines the place of the last sig fig in the answer. 3.75 mL + 4.1 mL 7.85 mL 224 g + 130 g 354 g  7.9 mL  350 g 3.75 mL + 4.1 mL 7.85 mL 224 g + 130 g 354 g

46. Practice Problems (15.30 g) ÷ (6.4 mL) = 2.390625 g/mL  18.1 g 18.9g - 0.84 g 18.06 g 4 SF2 SF  2.4 g/mL 2 SF

47. Introduction to the Periodic Table vertical columns groups, or familiesThe vertical columns of the periodic table are called groups, or families. Each group contains elements with similar chemical properties. horizontal rows periodsThe horizontal rows of elements in the periodic table are called periods. Physical and chemical properties change somewhat regularly across a period.

48. 48 The Unit Factor Method  Simple but important method to get correct answers in word problems.  Method to change from one set of units to another.  Visual illustration of the idea.

49. B. Dimensional Analysis  Steps: 1. Identify starting & ending units. 2. Line up conversion factors so units cancel. 3. Multiply all top numbers & divide by each bottom number. 4. Check units & answer.

50. B. Dimensional Analysis  How many milliliters are in 1.00 quart of milk? 1.00 qt 1 L 1.057 qt = 946 mL qtmL 1000 mL 1 L 

51. 51 The Unit Factor Method Example 1-2: Express 627 milliliters in gallons. You do it!

52. 52 The Unit Factor Method Example 1-2: Express 627 milliliters in gallons.

53. 53 The Unit Factor Method Example 1-3: Express 2.61 x 10 4 cm 2 in ft 2. Area is two dimensional, thus units must be in squared terms.

54. 54 The Unit Factor Method Example 1-3: Express 2.61 x 10 4 cm 2 in ft 2. Area is two dimensional, thus units must be in squared terms.

55. Derived Units  Combination of units.  Volume amount of space occupied by an object  length  length  length  (m 3 or cm 3 ) D = MVMV 1 cm 3 = 1 mL 1 dm 3 = 1 L Density (kg/m 3 or g/cm 3 )  mass per volume

56. 56 Density and Specific Gravity  density = mass/volume  D = M / V  How heavy something is for its size.  The ratio of mass to volume for a substance.  Independent of how much of it you have  gold - high density  air low density.

57. 57 Density and Specific Gravity Example 1-6: Calculate the density of a substance if 742 grams of it occupies 97.3 cm 3.

58. 58 Density and Specific Gravity Example 1-6: Calculate the density of a substance if 742 grams of it occupies 97.3 cm 3.

59. 59 Density and Specific Gravity Example 1-7: Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need? (liquid’s density = 1.32 g/mL) You do it!

60. 60 Density and Specific Gravity Example 1-7 Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need? (liquid’s density = 1.32 g/mL)

61. 61 Density and Specific Gravity  Water’s density is essentially 1.00 at room T.  Thus the specific gravity of a substance is very nearly equal to its density.  Specific gravity has no units.

62. 62 Density and Specific Gravity Example 1-8: A 31.0 gram piece of chromium is dropped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium? You do it

63. 63 Density and Specific Gravity Example 1-8: A 31.0 gram piece of chromium is dropped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium?

64. 64 Density and Specific Gravity Example 1-9: A concentrated hydrochloric acid solution is 36.31% HCl and 63.69% water by mass. The specific gravity of the solution is 1.185. What mass of pure HCl is contained in 175 mL of this solution? You do it!

65. 65 Density and Specific Gravity Example 1-9: A concentrated hydrochloric acid solution is 36.31% HCl and 63.69% water by mass. The specific gravity of the solution is 1.185. What mass of pure HCl is contained in 175 mL of this solution?

66. Heat and Temperature  Heat and Temperature are not the same thing  Temperature- measure of the average kinetic energy  Temperature is which way heat will flow. (from hot to cold)  3 common temperature scales - all use water as a reference 66

67. 67 Heat and Temperature MP waterBP water  Fahrenheit 32 o F 212 o F  Celsius 0.0 o C 100 o C  Kelvin 273 K 373 K

68. 68 Relationships of the Three Temperature Scales

69. 100ºC212ºF 0ºC 32ºF 100ºC = 212ºF 0ºC = 32ºF 100ºC = 180ºF How much it changes

70. 70 Relationships of the Three Temperature Scales

71. 71 Relationships of the Three Temperature Scales

72. 72 Relationships of the Three Temperature Scales Easy method to remember how to convert from Centigrade to Fahrenheit. 1. Double the Centigrade temperature. 2. Subtract 10% of the doubled number. 3. Add 32.

73. 73 Heat and Temperature Example 1-10: Convert 211 o F to degrees Celsius.

74. 74 Heat and Temperature Example 1-10: Convert 211 o F to degrees Celsius.

75. 75 Heat and Temperature Example 1-11: Express 548 K in Celsius degrees.

76. 76 Heat and Temperature Example 1-11: Express 548 K in Celsius degrees.

77. 77 Heat Transfer and the Measurement of Heat  Heat is energy, ability to do work.  SI unit J (Joule)  calorie Amount of heat required to heat 1 g of water 1 o C 1 calorie = 4.184 J  Calorie Large calorie, kilocalorie, dietetic calories Amount of heat required to heat 1 kg of water 1 o C  English unit = BTU  Specific Heat amount of heat required to raise the T of 1g of a substance by 1 o C unit = J/g o C

78. 78 Heat Transfer and the Measurement of Heat  Heat capacity amount of heat required to raise the T of 1 mole of a substance by 1 o C  unit = J/mol o C  Heat transfer equation necessary to calculate amounts of heat amount of heat = amount of substance x specific heat x  T

79. 79 Heat Transfer and the Measurement of Heat Example 1-12: Calculate the amt. of heat to raise T of 200.0 g of water from 10.0 o C to 55.0 o C

80. 80 Heat Transfer and the Measurement of Heat Example 1-12: Calculate the amt. of heat to raise T of 200.0 g of water from 10.0 o C to 55.0 o C

81. 81 Heat Transfer and the Measurement of Heat Example 1-13: Calculate the amount of heat to raise the temperature of 200.0 grams of mercury from 10.0 o C to 55.0 o C. Specific heat for Hg is 0.138 J/g o C. You do it!

82. 82 Heat Transfer and the Measurement of Heat Example 1-13: Calculate the amount of heat to raise the temperature of 200.0 grams of mercury from 10.0 o C to 55.0 o C. Specific heat for Hg is 0.138 J/g o C.

83. 83 Heating Curve for 3 Substances Which substance has the largest specific heat? Which substance’s T will decrease the most after the heat has been removed?

84. 84 Heating Curve for 3 Substances

85. 85 Synthesis Question  It has been estimated that 1.0 g of seawater contains 4.0 pg of Au. The total mass of seawater in the oceans is 1.6x10 12 Tg, If all of the gold in the oceans were extracted and spread evenly across the state of Georgia, which has a land area of 58,910 mile 2, how tall, in feet, would the pile of Au be? Density of Au is 19.3 g/cm 3. 1.0 Tg = 10 12 g.

86. 86 Synthesis Question

87. 87 Synthesis Question

88. 88 Synthesis Question