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Welcome to Chemistry 1: A college preparatory course.

Welcome to Chemistry 1: A college preparatory course. Cellphone = NO!!!! Webpage Need to Buy Scientific Calculator Notebook or Binder Safety Contract – Signed. Course Expectations HW & Review Make up missed work Notes Tests Flex: Tuesday-Friday Flex Shoes in Locker

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Welcome to Chemistry 1: A college preparatory course.

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  1. Welcome to Chemistry 1: A college preparatory course.

  2. Cellphone = NO!!!! • Webpage • Need to Buy • Scientific Calculator • Notebook or Binder • Safety Contract – Signed

  3. Course Expectations • HW & Review • Make up missed work • Notes • Tests • Flex: Tuesday-Friday Flex • Shoes in Locker • Letters of recommendation • “Are you a student, or just a kid who comes to school?”

  4. What do chemists study/do? • What professions/college majors require a chemistry course? • Where is chemistry important in business/industry? • What household products are “chemicals”? • Where in history was chemistry important?

  5. 1. Three specific things you must do to be successful in this course. 2. One thing you must never do (academically) in this course.

  6. Scientific Method A. Hypothesis – untested, educated guess B. Theory – successfully tested hypothesis C. Law – Theory with NO known exceptions

  7. Hypothesis Theory Law

  8. Science changes!!!!!

  9. Measuring • Estimated place – every measurement must have ONE estimated place. • One place past the smallest marking

  10. Measuring

  11. Measuring

  12. Measuring

  13. Measuring

  14. Measuring

  15. Graphing 1. Label x and y axis including units 2. Mark Axis using a convenient scale 3. Title your graph “The Dependence of Y on X” 4. Mark dots with a small circle 5. Draw “Best Fit” line or curve

  16. Graphing • Best Fit Line a. Used ONLY for linear relationships. b. Fits y = mx + b m = slope b = y-intercept c. If graph is almost perfect line, same # dots above and below X = independent variable (you can control) Y = dependent variable (can’t control)

  17. Graphing • Best Fit Curve a. Used if points are clearly not linear. b. Can be fit to higher order eqns: y = mx2 + b

  18. Graphing Rectangle A = L X W Triangle A = ½ B X H Circle A = r2 Irregular Shape?

  19. Graphing Lab • Use centimeters • TWO decimal places, last one is the estimated place • Write down the letter of your shape • See me for the actual value

  20. Scientific Notation • Descartes (1637) • Powers of 10 100 = 1 101 = 10 102 = 10 X 10 = 100 103 = 10 X 10 X 10 =1000

  21. Scientific Notation 200,000,000,000 stars (Andromeda): 2 X 100,000,000,000 2 X 1011 stars

  22. Scientific Notation 3. A Helium atom masses 0.000,000,000,000,000,000,000,006,645g 6.645 X 10-24 g

  23. Scientific Notation 4. Who created Sci. Notation? (“I think, therefore I am”)

  24. Scientific Notation 340 378,400 0.00234 0.000 000 000 0918 5.6 X 105 6.12 X 10-3 2.6 X 10-7 4 x 102

  25. Scientific Notation 43 575,400 0.000723 0.000 000 0014 6.5 X 10-5 2.16 X 103 6.2 X 107 8 x 10-2

  26. Scientific Notation There are ~900 students at Dallas 9 X 102 = 90 X 101 = 0.9 X 103 =

  27. Scientific Notation Write 4500 in scientific notation with the following exponents: X 103 X 102 X 105 X 104

  28. Scientific Notation Write 4500 in scientific notation with the following exponents: 4.5 X 103 45 X 102 0.045 X 105 0.45 X 104

  29. Scientific Notation Examples: (2.0 x 102) + (3.0 x 103) = 3.2 X 103 (6.0 X 103) ÷ (3.0 x 10-5)=2.0X108 (2.0 x 107) - (6.3 x 105) = 1.9X107

  30. Scientific Notation (4.0 x 105) x (3.0 x 10-1)= (6.0 x 108) ÷ (3.0 x 105)= (8.4x 1012) ÷ (8.4 x 109)= NOTE: 103 = 1 X 103

  31. Scientific Notation (4.0 x 105) x (3.0 x 10-1)=1.2 X105 (6.0 x 108) ÷ (3.0 x 105)= 2 X 103 (8.4x 1012) ÷ (8.4 x 109)= 1 X 103 NOTE: 103 = 1 X 103

  32. Accuracy and Precision • Accuracy – how close the average of a set of measurements is to the accepted value (AAA) • Precision – How close a set of measured values are to one another (reproducibility) • Always compare to a textbook value

  33. X X X X X X X X X X X X X X X X

  34. Percent Error Percent Error – Measure of accuracy % Error = Experimental – Accepted X 100 Accepted NOTE: “Experimental” =average of all trials

  35. A student measures the density of a sample of copper at 8.75 g/mL. The accepted value is 8.96 g/mL. Calculate the percent error.

  36. Error Analysis: Range Range - Measure of precision Range = highest trial – lowest trial

  37. Example 1 A student measures the density of a sample of lead and does four trials (11.3, 10.5, 11.9, 10.8 g/cm3). Calculate the range and comment on precision.

  38. Accuracy and Precision Students did trials to measure the density of a metal. The accepted density is 7.2 g/cm3. Were they accurate or precise? Set 1 7.21 7.25 7.18 Set 2 6.40 7.90 7.30 Set 3 6.45 6.52 6.48

  39. Significant Figures 1. Def - All of the measured values plus one estimated place • Examples 6 cm 6.0 cm 6.01 cm 0.005 mm 0.0050 mm 0.00500 mm 1340 kg 1340. kg 1340.0 kg

  40. Numbers with a Decimal How many sig figs? Also, write in sci. notation: 3.44 cm 60.001 cm 430.0 cm 0.0032 cm 0.00320 cm

  41. Numbers without a Decimal 1. Often poor measurements 2. Examples: “Not left” 18,500 kg 120 ft

  42. Numbers without a Decimal How many sig figs? Also, write in scientific notation: 10,500 cm 240 cm 120,000 cm 4 cm 45 cm

  43. Significant Figures How many significant figures are in the following? Also, write the numbers in proper scientific notation. 1508 cm 20.003 lb 300 ft 300.0 ft 0.00705 m 0.007050 m 1250 1250. 1250.0

  44. Significant Figures Round the following to three sig figs: 32.45 32.449 0.0067530 0.003904 11,980

  45. How many significant figures? 0.00200 0.0020 100. 7450 144.0 200 8.40 X 1010 9.000 X 10-5 Round to three significant figures: 54.649999 1.456 X 10-4 300.847 8.605 X 107 200.49 0.00056732 0.0045282

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