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Chem 140 Section A Instructor: Ken Marr

Chem 140 Section A Instructor: Ken Marr. Weekly Schedule “Lecture” 9 -10, MWF in STB-2 “Lab” 8 -10 , Tu in STB-2 8 -10 , Th in STB-5. Chem 140 Section C Instructor: Ken Marr. Weekly Schedule “Lecture” 10 –11, MWF in STB-2 “Lab” 10-12 , Tu in STB-2

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Chem 140 Section A Instructor: Ken Marr

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  1. Chem 140 Section AInstructor: Ken Marr Weekly Schedule “Lecture” 9 -10, MWF in STB-2 “Lab” 8 -10 , Tu in STB-2 8 -10 , Th in STB-5

  2. Chem 140 Section CInstructor: Ken Marr Weekly Schedule “Lecture” 10 –11, MWF in STB-2 “Lab” 10-12 , Tu in STB-2 10-12 , Th in STB-5

  3. Chem 140 Section EInstructor: Ken Marr Weekly Schedule “Lecture” 1 - 2, MWF in STB-2 “Lab” 1 - 3 , Tu in STB-2 1 - 3 , Th in STB-5

  4. Day 1 Activities • Introduction to Course • Briefly Review Course Outline/Syllabus • Homework Assignments • Reading: See Chem 140 schedule • Lab: Do Prelab assignment for the “Measurement and Density” Lab • Stamped Assignment #1: Chapter 1 HW • due Tues. 10/01/02 ……But start now!!! • Note: 10/2 is a very special day for your instructor!! • Begin Chapter 1 • Alice 1 and 2

  5. CHEMISTRY The Study of Matter and the Changes that Matter Undergoes and The Energy Associated with The Changes

  6. Chemistry as the Central Science Engineering Physics Atmospheric Sciences Oceanography Medicine Economics Governments Chemistry People Geology Biology Politics Astronomy Anthropology

  7. Chapter# 1 : Keys to the Study of Chemistry 1.1 Some Fundamental Definitions 1.2 Chemical Arts and the Origins of Modern Chemistry 1.3 The Scientific Approach: Developing a Model 1.4 Chemical Problem Solving 1.5 Measurement in Scientific Study 1.6 Uncertainty in Measurement: Significant Figures

  8. Measurement and Significant Figures • Measured Numbers are Never Exact...Why? • Which Graduated Cylinder is the most precise? • How is precision indicated when we record a measurement?

  9. The Number of Significant Figures in a Measurement Depends Upon the Measuring Device Fig 1.14 3e

  10. Significant Figures • We use significant figures to indicate the maximum precision of a measurement • Significant Figures • The number of digits that are known with certainty, plus one that is uncertain • Significant figures are used only with measured quantities. • Some numbers are exact and do not have any uncertainty......e.g...’s??

  11. More Examples • Record the exact length in centimeters, cm (T2c) • Record the exact amounts for numbers 1-11 (T2d)

  12. Sig. Fig. Rules to memorize.....(See page 27-30 Silberberg 3e) • All nonzero numbers are significant e.g. 23.8 g, 2345 km, 11 mL, 5 inches • Zeros between nonzero digits are significant i.e. Sandwiched zeros are significant e.g. 509 m, 2001 mL, 2050.1 L • Zeros preceding the 1st nonzero digit are not significant.......they serve only to locate the decimal point e.g. 0.083 m, 0.000306 L • Try converting these numbers to Scientific Notation to prove this!

  13. More Sig. Fig. Rules Involving Zeros • Zeros at the end of number that include a decimal point are significant • 0.800, 11.40, 10.00, 400. • Zeros at the end of a number without a decimal point are not significant... The Greenwater Rule! • 40, 8800, 300, • Use of underlining and decimal points

  14. Examples of Significant Digits in Numbers Number - Sig digits Number - Sig digits 0.0050 L 1.34000 x 107 nm 0.00012 kg 87,000 L 83.0001 L 78,002.3 ng 0.006002 g 0.000007800 g 875,000 oz 1.089 x 10 -6L 30,000 kg 0.0000010048 oz 5.0000 m3 6.67000 kg 23,001.00 lbs 2.70879000 ml 0.000108 g 1.0008000 kg 1,470,000 L 1,000,000,000 g

  15. Rounding off Numbers • Rounding off is used to drop non-significant numbers • Rule 1 When the 1st digit after those you want to retain is 4 of less, that digit and all others to the right are dropped • Round off the following to 3 sig. figs. • 105.29, 189.49999, 1.003, 100.3, 1001

  16. Rounding off Numbers • Rule 2 When the 1st digit after those you want to retain is 5 or greater, that digit and all others to the right are dropped and the last digit retained is increased by one • Round off the following to 4 sig. figs. • 10.87519, 13.59800, 99.999, 1042.5

  17. Sig. Figs. in Calculations • The Central Idea..... • The result of a calculation based on measurements can not be more precise than the least precise measurement! • Some Rules to, yes, memorize......

  18. Sig. Figs. in Multiplication and Division • “The Chain Rule” • Your answer must contain the same number of sig figs as the measurement with the fewest sig figs.....Some e.g...’s... (3.04) x (2.2) = 6.688 = ??? (2.00) / (0.3 ) = 6.666... = ??? (18.4) x (4.0) = 1.1117824 = ??? (66.2)

  19. Sig. Figs. in Addition and Subtraction • “The Decimal Rule” • The answer must have the same precision as the least precise measurement...or... • Your answer must be expressed to the same number of decimal places as the measurement with the fewest decimal places. • The number of sig figs are not considered, only the number of decimal places are considered!!! • Some examples..

  20. Sig. Figs. in Addition and Subtraction • Examples..... • 12.89 + 12.1 + 11.803 + 19 = 55.793 = ? • 1786 - 130 = 1656 = ??? • 7331 + 0.495 = 7331.495 = ???

  21. Scientific Notation • Scientific Notation • Writing a number as a number between 1 and 10 times a power of 10 • WHY DO IT??? • The Rules...

  22. How to Write Numbers in Scientific Notation • Move the decimal point in the original number so that it is located after the first nonzero digit • e.g. 5682  ???? • Multiply this number by the proper power of 10 • The power of 10 is equal to the number of places the decimal point was moved. • POSITIVE IF MOVED TO THE LEFT • NEGATIVE IF MOVED TO THE RIGHT

  23. Examples.... • Express the following numbers in scientific notation... • 0.0421 • 150,000 • 5899 • Express the following in “longhand” • 5.30 x 10-4 • 8.000 x 106

  24. Meaning of Powers of 10 • 103 = 10-3 = • 102 = 10-2 = • 101 = 10-1 = • 100 =

  25. Metric System • System of measure built around standard or base units • Uses factors of 10 to express larger or smaller numbers of these units

  26. Table 1. 2 (p. 17, 3e) SI - Base Units Physical Quantity Unit Name Abbreviation Mass Kilogram kg Length meter m Time second s Temperature Kelvin K Electric current ampere A Amount of substance mole mol Luminous intensity candela cd

  27. Metric Base Units and their Abbreviations • Length • Mass • Volume • Temperature • Prefixes are added to these base units for quantities larger or smaller than the base unit • Prefixes are a multiple of 10

  28. Table 1.3Common Decimal Prefixes Used with SI Units. Prefix Prefix Number Word Exponential Symbol Notation tera T 1,000,000,000,000 trillion 1012 giga G 1,000,000,000 billion 109 Mega M 1,000,000 million 106 Kilo k 1,000 thousand 103 hecto h 100 hundred 102 deka da 10 ten 101 ----- ---- 1 one 100 deci d 0.1 tenth 10-1 centi c 0.01 hundredth 10-2 milli m 0.001 thousandth 10-3 micro millionth 10-6 nano n 0.000000001 billionth 10-9 pico p 0.000000000001 trillionth 10-12 femto f 0.000000000000001 quadrillionth 10-15

  29. Common Metric Prefixes • Memorize the Symbol, Numerical Value, and Power of 10 Equivalent for..... • kilo- • centi- • milli- • micro- • nano-

  30. Common Prefix Applications • Length: • km 1 km = ? m • cm 1 cm = ? m • mm 1 mm = ? m • µm 1 µm = ? m • nm 1 nm = ? m

  31. Common Prefix Applications • Mass • kg 1 kg = ? g • mg 1 mg = ? g • µg 1 µg = ? g

  32. Common Prefix Applications • Volume • mL 1 mL = ? L • µL 1 µL = ? L

  33. Important Relationships • Length • 1 m = ?? cm • 1 m = ?? mm • 1 m = ?? µm • 1 cm = ?? mm

  34. Important Relationships • Mass • 1 g = ?? mg • 1 kg = ?? g • 1 kg = ?? lb..

  35. Some Volume Relationships in SI Units Fig. 1.9

  36. Important Relationships • Volume • 1 L = ?? mL • 1 mL = ?? cm3 • 1 L = ?? cm3 • 1 L = 1.057 qt.

  37. Solving Chemistry ProblemsDevelop a Plan  Carryout Plan  Check Answer • Developing a Plan: Read the problem carefully! • Clarify the know and unknown: • What information is given? • What are you trying to find? • Think about how to solve the problem before you start to juggle numbers • Suggest steps from the known to unknown • Determine principles involved and the relationships needed • Use sample problems as a guides • Map out the strategy you will follow

  38. Solving Chemistry Problems (cont.) • Solve the problem: Carry out your plan • Set up problem in a neat, organized, and logical way! • Unwanted units should cancel to give the desired unit of measure • Make a rough estimate of the answer before using your calculator • Round off to correct number of sig. figs. • Answer must have correct units

  39. Solving Chemistry Problems (cont.) • Check your answer • Is it reasonable? • Correct nits? • Same “ballpark” as a rough estimate? • Makes chemical sense?

  40. Problem Solving: Some Examples • How many hours would it take a pump to remove the water from a flooded basement that is about 30 feet wide and 50 feet long with water at a depth of about 2 feet? The pump has a capacity of 80 liters per minute. See Table 1.4, Common SI-English Equivalent Quantities, page 18 Silberberg 3e. 1062 min = 17.7 hours = 20 hours

  41. Metric Conversion Factors • Be able to do conversions within the metric system involving the common metric prefixes • kilo- • centi- • milli- • micro- • nano- • e.g. #32 on page 43

  42. Metric - English Conversions • Given metric - English conversion factors, be able to convert between these two systems • You do not have to memorize metric to English conversions factors

  43. Measurement of Temperature • Heat vs. Temperature • Temperature (SI unit: Kelvin, K) • A measure of how hot or cold an object is relative to another object • Also measured in degrees Celsius, oC • Heat (SI unit: joule, J) • The energy transferred between objects at different temperatures • A form of energy associated with the motion of atoms and molecules (the small particles of matter) • Also measured in calories, cal

  44. Application: Heat vs. Temperature • Which contains more heat... • 1 mL of water at 90 oC or 1 liter of water at 90 oC ? • 1 burning match or 10 burning matches?

  45. Temperature Conversions The boiling point of Liquid Nitrogen is - 195.8 oC, what is the temperature in Kelvin and degrees Fahrenheit? T (in K) = T (in oC) + 273.15 T (in K) = -195.8 + 273.15 = 77.35 K = 77.4 K T (in oF) = 9/5 T (in oC) + 32 T (in oF) = 9/5 ( -195.8oC) +32 = -320.4 oF The normal body temperature is 98.6oF, what is it in Kelvin and degrees Celsius? T (in oC) = [ T (in oF) - 32] 5/9 T (in oC) = [ 98.6oF - 32] 5/9 = 37.0 oC T (in K) = T (in oC) + 273.15 T (in K) = 37.0 oC + 273.15 = 310.2

  46. Density • Density = mass (g) / Volume(mL or cm3 or L) • Physical characteristic of a substance • Aids in identification of a substance • Calculated by..... • divide the mass of a substance by the volume occupied by that mass • Units • mass in grams • volume • Solids and Liquids: mL or cm3 • Gasses: L

  47. Density • Densities vary with temperature! • Why?? • Would you expect densities to increase or decrease as the temperature increases?

  48. Density • Immiscible liquids and solids separate into layers according to their densities • List the order from top to bottom when the following are mixed • Hg (13.5525 g/mL) • Carbon Tetrachloride (1.59525 g/mL) • Mg (1.7425 g/mL) • Water (1.004 g/mL) • What do the superscripts mean next to each density listed above?

  49. Calculations Involving Density • Be able to calculate the density, mass, or volume of a substance • Use the plug and chug method or use density as a conversion factor • Practice makes perfect....

  50. Specific Gravity • Compares the density of a liquid or solid to that of water... Units??? • Sp. Gravity = dsolid or liquid / dwater • Usually use dwater @ 4oC = 1.000g/mL • Compares the density of a gas to that of air...... Units??? • Sp. Gravity = dgas/ dair

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