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CHEM 1405

CHEM1405. 2. Assignments and Reminders. Reading AssignmentChapter 2 by TuesdayHomework Problems due Tuesday Chapter 1 problems 6, 8, 14, 16, 24, 26, 28, 30, 34, 36, 44, 46, 52, 56, 58For those in my laboratory section Reminder to have Safety goggles and appropriate clothing for lab class on TuesdayAlso don't forget to do prelab questions due at beginning of lab class.

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CHEM 1405

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    1. CHEM1405 1 CHEM 1405 Class Meeting 2

    2. CHEM1405 2 Assignments and Reminders Reading Assignment Chapter 2 by Tuesday Homework Problems due Tuesday Chapter 1 problems 6, 8, 14, 16, 24, 26, 28, 30, 34, 36, 44, 46, 52, 56, 58 For those in my laboratory section Reminder to have Safety goggles and appropriate clothing for lab class on Tuesday Also dont forget to do prelab questions due at beginning of lab class

    3. CHEM1405 3 Metric System, Calculations, Conversions, Density 7. What is the metric system of measurement? How does one convert between metric units and the units commonly used in the United States? 8. What is the difference between precision and accuracy in measurements? 9. Why is an understanding of significant figures important in chemistry? How do we determine the number of significant figures to report? 10. What is density? 11. What is the difference between temperature and heat? 12. How are the different temperature scales related to one another?

    4. CHEM1405 4

    5. CHEM1405 5 Measurement Consists of two parts a number AND a unit HAVE to have both E.g. I have a cat that is 3

    6. CHEM1405 6 Review of Scientific Notation Used to deal with very small or very large numbers as powers of 10 Examples: 0.00002 is written as 2 x 10-5 4,000,000 is written as 4 x 106 Note: a negative exponent just means its a number less than 1

    7. CHEM1405 7 Scientific Notation Review (Continued) 100 = 1 101 = 10 102 = 100 (= 10 x 10) 103 = 1000 ( = 10 x 10 x 10) 104 = 10000 105 = 100000 106 = 1000000 100 = 1 10-1 = 0.1 (= 1/10) 10-2 = 0.01 (= 1/(10 x 10)) 10-3 = 0.001 10-4 = 0.0001 10-5 = 0.00001 10-6 = 0.000001

    8. CHEM1405 8 Modern Metric System International System of Units (SI) SI comes from French Systeme International Based on decimal system All units related by factors of 10 Prefixes denote magnitude All measured quantities based on 7 base units

    9. CHEM1405 9 The Seven SI Base Units

    10. CHEM1405 10 Length Originally the meter was intended to equal 10-7 or one ten-millionth of the length of the meridian through Paris from pole to the equator. They were about 0.2 millimeter short. oops!

    11. CHEM1405 11 Metric Prefixes

    12. CHEM1405 12 Metric Lengths Distance from Campus to Scotty Ps Hamburgers 2560 meters 2.56 x 103 meters Basketball Diameter 0.146 meters 1.46 x 10-1 meter Dime Diameter 0.0179 meter 1.79 x 10-2 meter Dime Thickness 0.00135 meter 1.35 x 10-3 meter

    13. CHEM1405 13 Volume Derived from length

    14. CHEM1405 14 Measurements Numbers obtained from measurements are not exact Measurements are subject to error Calibration of the equipment may be off May not be able to read value accurately

    15. CHEM1405 15 Precision and Accuracy Precision is how closely members of a set of measurements agree with one another. It reflects the degree of reproducibility of the measurements. Accuracy the closeness of the average of the set to the "correct" or most probable value

    16. CHEM1405 16 Accuracy vs. Precision

    17. CHEM1405 17 Ability To Read a Scale

    18. CHEM1405 18 Significant Figures Not every number your calculator gives you can be believed The measuring device determines the number of significant figures a measurement has.

    19. CHEM1405 19 For example, if you measured the length, width, and height of a block you could calculate the volume of a block: Length: 0.11 cm Width: 3.47 cm Height: 22.70 cm Volume = 0.11cm x 3.47cm x 22.70cm = 8.66459 cm3 Where do you round off? = 8.66? = 8.6? = 8.7? 8.66459?

    20. CHEM1405 20 Rules for Significant Figures All nonzero digits are significant. 3.51 has 3 sig figs The number of significant digits is independent of the position of the decimal point 0.00000000034 and 56. Both have 2 sig figs Zeros located between nonzero digits are significant 4055 has 4 sig figs

    21. CHEM1405 21 Rules for Significant Figures (cont.) Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point. 5.7000 has 5 sig figs Trailing zeros are ambiguous if the number does not contain a decimal point 2000. versus 2000 Zeros to the left of the first nonzero integer are not significant. 0.00045 (note: 4.5 x 10-4)

    22. CHEM1405 22 Examples of Significant Figures How many significant figures are in the following? 7.500 2009 600. 0.003050 80.0330

    23. CHEM1405 23 Examples of Significant Figures cont

    24. CHEM1405 24 Scientific Notation and Significant Figures Often used to clarify the number of significant figures in a number. Example: 4,300 = 4.3 x 1,000 = 4.3 x 103 0.070 = 7.0 x 0.01 = 7.0 x 10-2

    25. CHEM1405 25 Sig Figs in Calculations Rules for Addition and Subtraction The answer in a calculation cannot have greater significance than any of the quantities that produced the answer. example: 54.4 cm + 2.02 cm 54.4 cm 2.02 cm 56.42 cm

    26. CHEM1405 26 Sig Figs in Calculation Rules for Multiplication and Division The answer can be no more precise than the least precise number from which the answer is derived. The least precise number is the one with the fewest significant figures.

    27. CHEM1405 27 For example, if you measured the length, width, and height of a block you could calculate the volume of a block: Length: 0.11 cm Width: 3.47 cm Height: 22.70 cm Volume = 0.11cm x 3.47cm x 22.70cm = 8.66459 cm3 Where do you round off? = 8.66? = 8.6? = 8.7? 8.66459?

    28. CHEM1405 28 Rules for Rounding Off Numbers If the leftmost digit to be dropped is less than 5, leave the final digit unchanged. If the leftmost digit to be dropped is greater than 5, increase the final digit by one. If the leftmost to be dropped is exactly 5, we round up if the preceding digit is odd and down if the preceding digit is even.

    29. CHEM1405 29 Examples of Rounding Rules Round following numbers to 3 significant figures

    30. CHEM1405 30 Examples of Rounding Rules Round following numbers to 3 significant figures

    31. CHEM1405 31 For example, if you measured the length, width, and height of a block you could calculate the volume of a block: Length: 0.11 cm Width: 3.47 cm Height: 22.70 cm Volume = 0.11cm x 3.47cm x 22.70cm = 8.66459 cm3 Where do you round off? = 8.66? = 8.6? = 8.7? 8.66459?

    32. CHEM1405 32 You dont always have the units you want or need The method used for conversion is called the Dimensional Analysis Unit Conversion

    33. CHEM1405 33 Unit Conversion Need to be able to convert between units We use these two mathematical facts to do the dimensional analysis a number divided by itself = 1 any number times one gives that number back

    34. CHEM1405 34 Dimensional Analysis Example: I have 3 dozen doughnuts how many doughnuts do I have We know 1 dozen = 12

    35. CHEM1405 35 Dimensional Analysis Example convert 1.47 miles to inches

    36. CHEM1405 36 Dimensional Analysis Example convert 67.34 kilometers to millimeters

    37. CHEM1405 37 Conversion Factors Some Conversions Between Common (U.S) and Metric Units Metric Common Mass 1 kg = 2.205 lb 453.6 g = 1 lb 28.35 g = 1 ounce (oz) Length 1 m = 39.37 in. 1 km = 0.6214 mile 2.54 cm = 1 in. a aThe U.S. inch is defined as exactly 2.54 cm. The other equivalencies are rounded off. Metric Common Volume 1 L = 1.057 qt 3.785 L = 1 gal 29.57 mL = 1 fluid ounce (fl oz)

    38. CHEM1405 38 How to remember length conversion Remember at least one of length conversions Use dimensional analysis to find others

    39. CHEM1405 39 Volume Example Convert 4832 cm3 to liters

    40. CHEM1405 40 Density: the ratio of mass to volume Most commonly used units are g/mL for liquids and solids g/L for gases Density

    41. CHEM1405 41 How is density useful? Allows us to relate how much stuff is in a volume Determines what materials will float

    42. CHEM1405 42 Density Example If 73.2 mL of a liquid has a mass of 61.5 g, what is its density in g/mL?

    43. CHEM1405 43 Density Example How much volume does 130.4 g of gold (density = 19.30 g/mL) occupy?

    44. CHEM1405 44 Specific Gravity Specific gravity: the density of a substance compared to water as a standard Often the health industry uses specific gravity to test urine and blood samples Also used by brewers to measure alcohol content of beer

    45. CHEM1405 45 Specific gravity - the ratio of the density of the object in question to the density of pure water at 4oC. 1.00 g/mL Specific Gravity

    46. CHEM1405 46 Example of Specific Gravity The density of copper at 20C is 8.92 g/mL. The density of water at the reference temperature 4oC is 1.00 g/mL. What is the specific gravity of copper?

    47. CHEM1405 47 Hint for Specific Gravity Specific gravity is really just the density (in g/mL) but without the units If the density of an object is 2.3 g/mL, what is the specific gravity of the object? 2.3

    48. CHEM1405 48 Energy: Heat and Temperature Heat flows from warmer objects to cooler objects Temperature is a property that tells us in what direction heat will flow Temperature is the degree of hotness or coldness of a body or environment (corresponding to its molecular kinetic energy)

    49. CHEM1405 49 Fahrenheit (F): defined by setting 0F at the coldest temperature he could achieve (ice/salt bath) and 100F at his body temperature This led to freezing point of water at 32F and the boiling point of water at 212F Temperature Scales

    50. CHEM1405 50 Celsius (C): defined by setting freezing point of water at 0C and boiling point of water at 100C Temperature Scales

    51. CHEM1405 51 Kelvin (K): defined by setting absolute zero as 0 Kelvin and and using the Celsius degree interval 0 K is the temperature where all molecular motion stops Temperature in Kelvin is proportional to average kinetic energy Temperature Scales

    52. CHEM1405 52 Comparing Temperature Scales Fahrenheit and Celsius

    53. CHEM1405 53 Comparing Temperature Scales Celsius and Kelvin

    54. CHEM1405 54 Comparing Temperature Scales

    55. CHEM1405 55 Heat Energy SI unit of heat is the joule(J) calorie(cal) is another unit of heat energy A calorie is defined as the is the amount of heat required to raise the temperature of 1 g of water 1 C

    56. CHEM1405 56 Little c calories and Big C Calories 1000 cal = 1 kilocalorie The Calories on a food label are kilocalories

    57. CHEM1405 57 Specific Heat The specific heat of a substance is the quantity of heat required to raise the temperature of one gram of substance by 1 C (or 1 K). From definition of the calorie the specific heat of water is 1.00 cal/(gC)

    58. CHEM1405 58 Specific Heat Substance cal/(gC) J/(gC) Aluminum 0.216 0.902 Copper (Cu) 0.0921 0.385 Ethyl alcohol 0.588 2.46 Iron (Fe) 0.106 0.443 Ethylene glycol 0.561 2.35 Magnesium (Mg) 0.245 1.025 Mercury (Hg) 0.0332 0.139 Sulfur 0.169 0.706 Water (H2O) 1.000 4.182

    59. CHEM1405 59 Using Specific Heat

    60. CHEM1405 60 Using Specific Heat

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