1 The Foundations of Chemistry

2 Matter and Energy Chemistry – A Molecular View of Matter States of Matter Chemical and Physical Properties Chemical and Physical Changes Mixtures, Substances, Compounds, and Elements Measurements in Chemistry Units of Measurement Chapter Outline

3 Use of Numbers The Unit Factor Method (Dimensional Analysis) Percentage Density and Specific Gravity Heat and Temperature Heat Transfer and the Measurement of Heat Chapter Outline

4 Chemistry –Science that describes matter – its properties, the changes it undergoes, and the energy changes that accompany those processes Matter –Anything that has mass and occupies space. Energy –The capacity to do work or transfer heat. Scientific (natural) law –A general statement based the observed behavior of matter to which no exceptions are known. Matter and Energy - Vocabulary

5 Natural Laws Law of Conservation of Mass Law of Conservation of Energy Law of Conservation of Mass-Energy –Einstein’s Relativity –E=mc 2

6 Scientific Method Observation Hypothesis Observation or experiment Theory Observation or experiment Law

7 A Molecular View Dalton’s Atomic Theory Dalton’s atomic theory summarized the nature of matter as known in )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.

8 A Molecular View Some Definitions An atom is the smallest particle of an element that maintains its chemical identity through all chemical and physical changes. Fundamental particles are the basic building blocks of atoms, they consist of electrons, protons, and neutrons. A molecule is the smallest particle of an element or compound that can have a stable independent existence.

9 States of Matter Solids

10 States of Matter Solids Liquids

11 States of Matter Solids Liquids Gases

12 States of Matter Change States –heating –cooling

13 States of Matter Illustration of changes in state –requires energy

14 Chemical and Physical Properties Chemical Properties - chemical changes –rusting or oxidation –chemical reactions Physical Properties - physical changes –changes of state –density, color, solubility Extensive Properties - depend on quantity Intensive Properties - do not depend on quantity

15 Mixtures, Substances, Compounds, and 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 chart

16 Mixtures, Substances, Compounds, and Elements

17 Mixtures, Substances, Compounds, and Elements Compounds –substances composed of two or more elements in a definite ratio by mass –can be decomposed into the constituent elements Water is a compound that can be decomposed into simpler substances – hydrogen and oxygen

18 Mixtures, Substances, Compounds, and Elements

19 Mixtures, Substances, Compounds, and Elements Mixtures –composed of two or more substances –homogeneous mixtures –heterogeneous mixtures

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

21 Measurements in Chemistry Metric Prefixes NameSymbolMultiplier mega M 10 6 kilo k 10 3 deka da 10 deci d centi c 10 -2

22 Measurements in Chemistry Metric Prefixes NameSymbolMultiplier milli m micro  nano n pico p femto f

23 Units of Measurement Definitions Mass –measure of the quantity of matter in a body Weight –measure of the gravitational attraction for a body

24 Units of Measurement Common Conversion Factors Length –1 m = inches –2.54 cm = 1 inch Volume –1 liter = 1.06 qt –1 qt = liter See Table 1-8 for more conversion factors

25 Use of Numbers Exact numbers – 1 dozen = 12 things for example Accuracy –how closely measured values agree with the correct value Precision –how closely individual measurements agree with each other

26 Use of Numbers Significant figures –digits believed to be correct by the person making the measurement Measure a mile with a 6 inch ruler vs. surveying equipment Exact numbers have an infinite number of significant figures = 1 dozen because it is an exact number

27 Use of Numbers Significant Figures - Rules Leading zeroes are never significant has three significant figures Trailing zeroes may be significant must specify significance by how the number is written 1300 nails - counted or weighed? Use scientific notation to remove doubt 2.40 x 10 3 has ? significant figures

28 Use of Numbers Scientific notation for logarithms take the log of 2.40 x 10 3 log(2.40 x 10 3 ) = How many significant figures? Imbedded zeroes are always significant has five significant figures

29 Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

30 Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

31 Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

32 Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends

33 Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends

34 Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends

35 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.

36 The Unit Factor Method Change from a to a by obeying the following rules.

37 The Unit Factor Method Change from a to a by obeying the following rules. 1.Must use colored fractions.

38 The Unit Factor Method Change from a to a by obeying the following rules. 1.Must use colored fractions. 2.The box on top of the fraction must be the same color as the next fraction’s bottom box.

39 The Unit Factor Method R Fractions to choose from R OR O O B B O B BB B

40 The Unit Factor Method Fractions to choose from R OR O O B B O B BB B R O R

41 The Unit Factor Method Fractions to choose from R OR O O B B O B BB B R O R O B

42 The Unit Factor Method Fractions to choose from R OR O O B B O B BB B R O R O B B B B

43 The Unit Factor Method Fractions to choose from R OR O O B B O B BB B R O R O B B B B

44 The Unit Factor Method Fractions to choose from R OR O O B B O B BB B R O R O B B B B

45 The Unit Factor Method Fractions to choose from R OR O O B B O B BB B R O R O B B B B

46 The Unit Factor Method colored fractions represent unit factors 1 ft = 12 in becomes or Example 1-1: Express 9.43 yards in millimeters.

47 The Unit Factor Method

48 The Unit Factor Method

49 The Unit Factor Method

50 The Unit Factor Method

51 The Unit Factor Method R O R O B B B T B T

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

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

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

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

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

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

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

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

60 Percentage Percentage is the parts per hundred of a sample. Example 1-5: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore? You do it!

61 Percentage Percentage is the parts per hundred of a sample. Example 1-5: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore?

62 Density and Specific Gravity density = mass/volume What is density? Why does ice float in liquid water?

63 Density and Specific Gravity density = mass/volume What is density? Why does ice float in liquid water?

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

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

66 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!

67 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

68 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

69 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.

70 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

71 Density and Specific Gravity Example1-8: A 31.0 gram piece of chromium is dipped 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?

72 Density and Specific Gravity Example1-8: A 31.0 gram piece of chromium is dipped 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?

73 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 What mass of pure HCl is contained in 175 mL of this solution? You do it!

74 Density and Specific Gravity

75 Density and Specific Gravity

76 Density and Specific Gravity

77 Heat and Temperature Heat and Temperature are not the same thing T is a measure of the intensity of heat in a body 3 common temperature scales - all use water as a reference

78 Heat and Temperature Heat and Temperature are not the same thing T is a measure of the intensity of heat in a body 3 common temperature scales - all use water as a reference

79 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

80 Relationships of the Three Temperature Scales

81 Relationships of the Three Temperature Scales

82 Relationships of the Three Temperature Scales

83 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.

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

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

86 Heat Transfer and the Measurement of Heat SI unit J (Joule) calorie Amount of heat required to heat 1 g of water 1 o C 1 calorie = 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 units = J/g o C

87 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 units = J/mol o C Example 1-12: Calculate the amt. of heat to raise T of g of water from 10.0 o C to 55.0 o C

88 Heat Transfer and the Measurement of Heat Heat transfer equation necessary to calculate amounts of heat amount of heat = amount of substance x specific heat x   T

89 Heat Transfer and the Measurement of Heat Heat transfer equation necessary to calculate amounts of heat amount of heat = amount substance x specific heat x   T

90 Heat Transfer and the Measurement of Heat Heat transfer equation necessary to calculate amounts of heat amount of heat = amount substance x specific heat x   T

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

92 Heat Transfer and the Measurement of Heat Example 1-13: Calculate the amount of heat to raise T of g of Hg from 10.0 o C to 55.0 o C. Specific heat for Hg is J/g o C. Requires 30.3 times more heat for water is 30.3 times greater than 0.138

93 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?

94 Heating Curve for 3 Substances

1 The Foundations of Chemistry