1 / 72

Chapters 1 & 2: Measurement and Calculations

Chapters 1 & 2: Measurement and Calculations. Learning Targets Identify a given substance as an element or compound Classify properties and changes as chemical or physical Explain the structure of the periodic including properties of elements based on their location (metal, nonmetal, etc.)

calvin-fischer
Download Presentation

Chapters 1 & 2: Measurement and Calculations

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapters 1 & 2:Measurement and Calculations Learning Targets Identify a given substance as an element or compound Classify properties and changes as chemical or physical Explain the structure of the periodic including properties of elements based on their location (metal, nonmetal, etc.) Determine the amount of heat transferred in a process Express data and results of calculations with appropriate significant figures, units, and in scientific notation Calculate percent error from lab data and use this to evaluate the quality of lab data Perform density calculations and apply density conceptually (e.g. identifying substances or determining if an object will float

  2. Key Words • atom • compound • element • pure substance • mixture • homogeneous/heterogeneous • chemical change/property • physical change/property • direct/inverse proportion • family/group • period • significant figure • scientific notation • conversion factor • percent error • metal • nonmetal • metalloid • noble gas • quantitative • qualitative

  3. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 • Matter is anything that has volume and mass • Mass is a measure of the amount of matter present • Volume is the amount of space an object occupies • Atoms are the fundamental building block of matter • Elements are pure substances made of one type of atom • Compounds are substances made from the atoms of two or more elements

  4. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 Properties of Matter • Matter has physical properties and can undergo physical changes • Physical properties are observed without changing the identity of a substance (examples include: density, color, melting point) • Physical changes don’t involve a change in a substance’s identity (examples include: phase change, breaking)

  5. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 Properties of Matter • Matter has chemical properties and can undergo chemical changes • Chemical properties relate to a substance’s ability to become a new substance (example: flammability, corrosiveness) • Chemical changes involve a change in a substance’s identity • Chemical Equation (Reactants to Products)

  6. Section 1: Classifying and Changing Matter Signs of a Chemical Change • Gas Production • Formation of a Precipitate • Dramatic Temperature Change • Unexpected Color Change

  7. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 Classification of Matter • Matter can be classified as either a pure substance or a mixture • Pure Substances are fixed ratios and can be either a element or a compound. • Elements are found on the periodic table, They can not be broken down by physical means. • Compounds are substances made of two or more elements.

  8. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 Classification of Matter • Mixtures are blends of two or more types of matter, each retaining its own identity and properties • Mixtures are combined physically and can be separated using physical means • There are two types of mixtures: • Homogenous –they are unified throughout • Heterogenous – not evenly mixed

  9. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23

  10. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 Periodic Table • The periodic table organizes elements into groups based on similar properties • The vertical columns are known as groups or families; elements in the same group have similar chemical properties • The horizontal rows are called periods; elements in the same period don’t necessarily have similar chemical properties

  11. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23

  12. Section 1: Classifying and Changing MatterPages 26-27 RBQs Pgs. 59-61 #9-11,14, 15, 18-21,23 • Metals: • Nonmetals: • Metalloids: • Noble Gases:

  13. Which of the following is a chemical change? • Sanding wood • Melting ice • Milk going sour • Vaporizing of gasoline

  14. Blood would be considered • Element • Compound • Homogenous mixture • Heterogenous mixture

  15. Section 2 : Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 Scientific notation is a ay of taking very large numbers and/or very small numbers and writing them more simply For example, an important number in chemistry is 602,000,000,000,000,000,000,000which suck to write…but in scientific notation it is6.02 x 1023

  16. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 Exercise: Use the examples below to come up with a set of rules for converting from scientific to regular notation.Sci. NotationReg. NotationSci. NotationReg. Notation 4.521 x 105 452,100 8.2 x 10-8 .0000000823.8862 x 102 388.62 6.447 x 10-4 .0006447

  17. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 Exercise: Use the examples below to come up with a set of rules for converting from regular to scientific notation.Sci. NotationReg. NotationSci. NotationReg. Notation 817 8.18 x 102 0.00456 4.56 x 10-3 0.000006 6 x 10-6 48260000 4.826 x 107

  18. Convert 506100 to scientific notation: • 5 x 105 • 5.1 x 10 -5 • 5.061 x 105 • 51 x 105

  19. 4.02 x 103 in standard notation is • .00402 • .000402 • 4020 • 40020

  20. Which of the following is not in scientific notation? • 2.31 x 108 • 2.31 x 10-3 • 231 x 107 • 2.31 x 1056 • 2.31 x 103

  21. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 • Measurements made in the lab are never perfect • Data can only include numbers that we are sureof • Record all numbers you aresure of, then estimate anadditional number

  22. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 • Numbers given this way are called significant figures • Sig figs also indicate how accurate a measuring device is

  23. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 • Since numbers are often given in a question and not gathered in a lab, it is important to be able to look at a number and determine how many significant figures it contains • Also, the number of significant figures in an answer will depend on the sig figs contained in the numbers in the question • Applying sig fig rules to math ensures that no uncertain numbers will be in your answers

  24. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 When are nonzero numbers significant? (circle one) Always Sometimes Never When are leading zeros significant? (circle one) Always Sometimes Never When are captive zeros significant? (circle one) Always Sometimes Never When are trailing zeros significant? (circle one) Always Sometimes Never

  25. How many significant figures are in 2.0900 ? • 1 • 2 • 3 • 4 • 5

  26. How many significant figures are in 0.0006042? • 7 • 3 • 8 • 4 • 0

  27. Round 14.859 to three significant figures • 15.0 • 14.9 • 14.81 • 14.809 • 14.8

  28. Round 2.00152 to four significant figures • 2.002 • 2.001 • 2.000 • 2.152 • None of these

  29. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 • Significant figure rules also exist and must be applied when performing calculations • There is one rule for addition and subtraction and a second rule for multiplication and division • The exercises on the following slides will illustrate these rules

  30. Section 1: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 X & ÷ 2.0 x 4 = 89.166 x 3.2 = 292.66543 x 0.0032 = .00850.02 ÷ 0.00606894 = 3 2.44 x 8.629 = 21.1199.2 ÷ 4.05 = 49.20.026 x 0.00449 = .00012(5.4 x 102)(6.39 x 10-6) = 3.5 x 10-3 Determine the number of significant figures in each answer above. Determine the number of significant figures in each number in the questions above. How is the number of significant figures in the answer determined? (write in space below)

  31. Section 2: Scientific Notation & Sig FigsPages 46-52 RBQs Pgs. 59-61 #21-23, 36-46 + & - 8.663 – 2.1 = 6.61.00036 + 0.2 = 1.28.365434534385 + 1 = 968.633 + 7.9343 = 76.567 14.2 + 2 = 169.887467 – 2.003 = 7.8846.22 + 2.1 = 8.34.0 + 12.98373 = 17.0 Determine the number of decimal places in each answer above. Determine the number of decimal places in each number in the questions above. How is the number of decimal places in the answer determined? (write in space below)

  32. Sample: What is the product of 4.56 and 1.4, reported with correct significant figures? • 6.384 • 6.38 • 6.3 • 6.4 • 6

  33. Sample: What is 4.56 - 1.4, reported with correct significant figures? • 3.16 • 3.2 • 3.1 • 3.160 • 3

  34. Sample: Do this! And with correct sig figs! 4.184 x 100.62 x (25.27 – 24.16) = ? • 470 • 467.3 • 467.30 • 460 • 467

  35. Sample: Do this! And with correct sig figs! (6.0 x 1023)(4.22) = ? • 2.532 x 1024 • 2.5 x 1024 • 2.53 x 1024 • 3 x 1024 • 2.5320 x 1024

  36. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 A key skill in chemistry is being able to convert from one unit of measurement to another For example, converting from one unit of distance to another such as feet to miles Using conversion factors is done using the same approach taken to multiplying fractions

  37. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 Determine the answers to following problems: 3 x 4 4 b) 2 x 5 5 7

  38. Sample: What does 3 x 2/3 equal? • 2 • 3 • 1/2 • 3/2 • 6 0 of 0

  39. Sample: What is the product of 3/4 and 2/3? • 3 • 4 • 2/3 • 3 • 1/2 0 of 2

  40. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 In each case, notice how a common numerator and denominator cancelled each other This same idea is the key idea to using conversion factors With conversion factors, the difference is that you select the fraction to the answer you want

  41. Section 2: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 To convert one unit to another, e.g. pounds to grams, the same principles as above are used Arrange units as needed to get the desired unit The fraction used to convert one unit to another is known as a conversion factor pounds x ---------------- = grams

  42. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 Set up the appropriate conversion factor for the following: Inches to centimeters Miles per hour to meters per minute

  43. Sample: What is the correct conversion factor for converting feet to inches? A. Feet Inches • Inches Feet

  44. Sample: What is the correct conversion factor to convert feet per second to inches per minute? • Feet x seconds inches minutes B. Inches x minutes feet seconds C. Feet x minutes inches seconds D. Inches x seconds minutes feet

  45. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 The numerical relationship between units must be taken into account as well For example, to convert feet to inches, you need to know that there are 12 inches in one foot Once the units are in place, the final step is to put each number with its unit Potentially Useful Information 1 ft3 = 28.32 L 1 mi = 1.609 km 1 in3 = 16.38 cm3 1 in = 2.54 cm 1 kg = 2.2 lbs 1 oz = 28.35 g 1 lb = 16 oz 1 gallon = 3.785 L 1 lb = 453.59 g 1 ft = 12 in 1 ft3 = 1728 in3 3 ft = 1 yd 1 m = 3.281 ft 1 mi = 5280 ft 1 cal = 4.184 J

  46. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 Determine the answers to following problems: a) How many inches are there in 2.0 feet? b)How many seconds are in 3 hours?

  47. Sample: How many grams are equal to a mass of 10.0 pounds? • .022 g • 4535.9 g • 45.359 g • 22 g

  48. Sample: How many inches are in 5 yards? • 1.25 • 20 • 180 • .139

  49. Section 3: Conversion FactorsPages 40-42 RBQs Pgs. 60-61 #28-32, 49 • Convert miles per hour to meters per minute

  50. Sample: What speed, in meters per minutes, is equivalent to 20.0 feet per second? • 3930 m/min • .101 m/min • 1.09 m/min • 366 m/min

More Related