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Basic Concepts

1. Basic Concepts. Chapter Outline. States of Matter Chemical and Physical Properties Chemical and Physical Changes Mixtures, Substances, Compounds, and Elements Measurements in Chemistry Units of Measurement. Chapter Outline. Use of Numbers

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Basic Concepts

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  1. 1 Basic Concepts

  2. Chapter Outline • States of Matter • Chemical and Physical Properties • Chemical and Physical Changes • Mixtures, Substances, Compounds, and Elements • Measurements in Chemistry • Units of Measurement

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

  4. Mixtures, Substances, Compounds, and Elements • Matter-Anything that occupies space and has mass. • Pure Substances or Substances-Cannot be separated by physical processes. -Elements-A substance which cannot be broken down into simpler substances. e.g. Na, He, C, (atoms) or N2, Cl2 (molecules) -Compounds-A pure substance made up of two or more elements. e.g. NaCl, H2O. • Mixtures-Can be separated by physical processes. • composed of two or more substances • homogeneous mixtures-A mixture that is uniform throughout-e.g. white wine, grape juice. Clear. Solutions. • heterogeneous mixtures-A mixture that is not uniform throughout-e.g. oil and water, orange juice. Cloudy.

  5. Mixtures, Substances, Compounds, and Elements

  6. Substances and Mixtures Matter Physical process Mixtures Pure Substances Chemical Reaction Homogeneous Mixture Heterogeneous Mixture Compounds Elements

  7. Fig. 1-7, p. 10

  8. States of Matter • Change States • heating • cooling Vaporization: Evaporation Boiling Freezing: Solidification Crystallization Melting: Fusion

  9. States of Matter Solid Liquid Gas heat heat cool cool Less attractive force and more disordered.

  10. States of Matter • Illustration of changes in state • requires energy

  11. Mixtures, Substances, Compounds, and Elements

  12. Elements you Need to Know:

  13. Types of Solutions: • Liquid solutions are the most common, but there are also gas and solid solutions. • Solutions have two components: • Solute - Solution component(s) present in lesser amounts. • Solvent - Solution component present in the greatest amount.

  14. Types of Solutions:

  15. Characteristics of Solutions • Uniform distribution • Components do not separate upon standing. • Components cannot be separated by filtration. • Within certain limits its composition can vary. • Almost always transparent. (i.e. one can see through it). • An alloy is a homogeneous mixture of metals. i.e. brass, bronze, sterling silver.

  16. Chemical and Physical Properties • Physical Properties – A property that can be observed in the absence of any change in composition. e.g. color, odor, taste, melting point, boiling point, freezing point, density, length, specific heat, density, solubility. • Physical Changes-Changes observed without a change in composition. i.e. cutting wood, melting of solids and boiling of liquids. • water, ice water, liquid water, steam • changes of state • dissolving • polishing

  17. Chemical and Physical Properties • Chemical Properties-A property that matter exhibits as it undergoes changes in composition. e.g. coal and gasoline can burn in air to form carbon dioxide and water; iron can react with oxygen in the air to form rust; bleach can turn blond hair blonde. • Chemical Changes-Changes observed only when a change in composition is occurring. e.g. reaction of sodium with chlorine, rusting of iron, dying of hair, burning of wood, cooking an egg, rotting food. • Extensive Properties - depends on the amount of material present. e.g. volume and mass. • Intensive Properties – does not depend on the amount of material present. e.g. melting point, boiling point, freezing point, color, density.

  18. Natural Laws • Law of Conservation of Mass-Mass is neither created nor destroyed. • Law of Conservation of Energy-Energy is neither created nor destroyed, only converted from one form to another. • Law of Definite Proportions-Different samples of any pure compound contain the same element in the same proportion by mass. e.g. water (H2O) contains 11.1 % H and 88.9% O by mass. Thus, a 25.0 sample of water would contain 2.78 g of H and 22.2 g of O.

  19. Law of Definite Proportions 11.1 % H and 88.9% O by mass, 25.0 g sample of water:

  20. Use of Numbers • Exact numbers • 1 dozen = 12 things for example

  21. Rounding off Numbers Previous digit 1.29 4 Next digit • If the next digit is less than 5 the previous .9946 .99 • digit remains the same. 1.294 1.29 • 2. If the next digit is greater than 5 or 5 • followed by non zeros then the previous digit .999 1.00 • is increased by one. 1.2951 1.30 • 3. If the next digit is 5 or 5 followed by all zeros 1.285 1.28 • then the previous digit remains the same if it 1.295 1.30 • is even or increased by one if it is odd. 1.22500 1.22

  22. Scientific Notation • Used to handle very large and very small numbers. Any number that is from + or – 1 to 9 N. X 10x For example: 3.21 x 103 -9.9 x 10-4 1.0 x 100 (Note that 100 is 1) Exponent-Power of 10

  23. Scientific Notation • To convert numbers to scientific notation use the following guidelines: 1750.0 = 1750.0 x 100= 1.7500 x 103 Exponent increases by 3 powers of 10 Number decreases by 3 powers of 10 A you move the decimal place to the left (i.e. make the number smaller), the power of ten (i.e., exponent) must increase by the same amount.

  24. Scientific Notation 0.050 = 0.050 x 100 = 5.0 x 10-2 The number gets larger by 2 powers of 10 The exponent gets smaller by 2 powers of 10. As you move the decimal place to the right (i.e. make the number larger), the power of ten (i.e. exponent) must decrease by the same amount.

  25. 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 12.000000000000000 = 1 dozen because it is an exact number

  26. Use of Numbers Significant Figures - Rules • Leading zeroes are never significant 0.000357 has three significant figures • Trailing zeroes may be significant must specify significance by how the number is written • Use scientific notation to remove doubt 2.40 x 103 has ? significant figures

  27. Use of Numbers • 3,380 ? significant figures 3.38 x 103 • 3,380. has ? significant figures 3.380 x 103 • Imbedded zeroes are always significant 3.0604 has ? significant figures

  28. Use of Numbers • Piece of Paper Side B – enlarged • How long is the paper to the best of your ability to measure it? 13.36 in. The second decimal place is estimated

  29. Use of Numbers • Piece of Paper Side A – enlarged • How wide is the paper to the best of your ability to measure it? 8.3 in The first decimal place is estimated

  30. Manipulating Powers of 10 • a) When multiplying powers of ten, the exponents are added. For example: 105 x 10-4 = 105+(-4)=101 b) When dividing powers of ten, the exponents are subtracted. For example: 104 = 104-(-4) = 108 10-4 c) When raising powers of ten to an exponent, the exponents are multiplied. For example: (104)3 = 10(4 x 3) = 1012

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

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

  34. Multiplying and Dividing Numbers with Powers of Ten • When using scientific notation: a.) Place the powers of ten together. (1.76 x 10200) x (2.650 x 10200)= (1.76 x 2.650) x (10200 + 200)= b.) The final answer has the same number of significant figures as the number with the least number of significant figures. 4.66 x 10400 c.) You must round off correctly. d.) Preferably report the answer in scientific notation.

  35. Multiplying and Dividing Numbers with Powers of Ten (1.760 x 102) /(2.65 x 10-2)= (1.760 / 2.65) x (102 – (-2))= 0.664 x 104 = 6.64 x 103

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

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

  38. Addition and Subtraction with Powers of Ten a.) All numbers must have the same power or ten before addition or subtraction is performed. b.) Once the powers of ten are the same, the coefficients can then be added or subtracted while the power of ten remains the same. c.) After adding or subtracting the coefficients, the answer must have the same number of decimal places as the coefficient with the fewest decimal places at the time of the operation. d.) You must round off correctly. e.) Preferably report the answerin scientific notation.

  39. Addition and Subtraction with Powers of Ten • 4.76 x 10200 + 9.6 x 10201 = ? 0.4 76 x 10201 + 9.6 x 10201 10.0 76 x 10201 1.01 x 10202 (written in scientific notation and rounded off to the correct number of significant figures)

  40. Addition and Subtraction with Powers of Ten • 2.95 x 10-15 – 1.00 x 10-14 = ? -1.00 x 10-14 0.29 5 x 10-14 -0.70 5 x 10-14 -7.0 x 10-15 (written in scientific notation and rounded off to the correct number of significant figures)

  41. Mixing Addition/Subtraction with Multiplication/Division 7.54 x 10-5 (99. x 10200 + 1.25 x 10201) (1.75 x 10-3)3 7.54 x 10-5 (9.9 x 10201 + 1.25 x 10201) = 1.75 x 10-3 x 1.75 x 10-3 x 1.75 x 10-3 7.54 x 10-5 [(9.9 + 1.25) x 10201) = 1.75 x 1.75 x 1.75 x 10-3 x 10-3 x 10-3 7.54 x 10-5 (11.2 x 10201) = 5.36 x 10-9 7.54 x 11.2 x 10-5 x 10201 = 1.58 x 10206 5.36 10-9

  42. Measurements in Chemistry QuantityUnitSymbol • length meter m • mass kilogram kg • time second s • current ampere A • temperature Kelvin K • amt. substance mole mol

  43. Measurements in ChemistryMetric Prefixes NameSymbolMultiplier • mega M 106 • kilo k 103 • deka da 10 • deci d 10-1 • centi c 10-2

  44. Measurements in ChemistryMetric Prefixes NameSymbolMultiplier • milli m 10-3 • micro  10-6 • nano n 10-9 • pico p 10-12 • femto f 10-15

  45. Metric Conversions • 1 km = 103 m • 1 dL = 10-1 L • 1 msec = 10-3 sec • 1 m = 10-6 m

  46. Fig. 1-20, p. 24

  47. Metric English Conversions Common Conversion Factors • Length • 2.54 cm = 1 inch (exact conversion) • Volume • 1 qt = 0.946 liter (Rounded off) • Mass _ 1 lb = 454 g (Rounded off)

  48. Use of Conversion Factors in Calculations • Commonly known relationship (i.e. equality): • 1 ft = 12 in • Respective conversion factors to above equality: 1 ft or 12 in 12 in 1 ft Use the conversion factor that allows for the cancellation of units. Convert 24 in to ft: ? ft = 24 in x

  49. Conversion Factors • Example 1-1: Express 9.32 yards in millimeters.

  50. Conversion Factors

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