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Warm-up Complete the following conversions:

Warm-up Complete the following conversions:. 5days = _______ s 2.5 m = ______ cm. Chemistry Chapter 2 Scientific Measurement. WARNING: Learn it now...it will be used all year in Chemistry!!!. 3.1 Qualitative measurements

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Warm-up Complete the following conversions:

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  1. Warm-upComplete the following conversions: • 5days = _______ s • 2.5 m = ______ cm

  2. ChemistryChapter 2Scientific Measurement WARNING: Learn it now...it will be used all year in Chemistry!!!

  3. 3.1 Qualitative measurements • measurements that give results in a descriptive, non-numerical form. Examples: He is tall Electrons are tiny

  4. Quantitative measurements • measurement that gives results in a definite form, usually as numbers and units. Examples: He is 2.2 m tall Electrons are 1/1840 times the mass of a proton

  5. Scientific Notation • a number is written as the product of two numbers: a coefficient and 10 raised to a power. Examples: 567000 = 5.67 X 105 0.00231 = 2.31 X 10-3

  6. Examples: Convert to or from Scientific Notation: 2.41 x 102 6.015 x 103 1.62 x 10-2 5.12 x 10-1 662 .0034 241 = 6015 = 0.0162 = 0.512 = 6.62 x 102 = 3.4 x 10-3 =

  7. Tuesday, September 4Bellwork • Jenna burns a piece of firewood that has a mass of 1.29 kg. Later, she sweeps up the ashes and takes their mass. They only weigh 0.99 kg. • What happened? • Does this break the law of conservation of mass? Why or why not? • How much mass was lost?

  8. Homework Check Give 3 examples of pure substances and a physical property of that example. 2. List the Phase(s) (S, L, G) below then match to the example on the left. _____ Heterogeneous Mixture Copper Wire _____ Homogenous Mixture Iced Tea _____ Element Sugar _____ Compound Soda 3. (circle one )When you did the fish lab you made qualitative and quantitative observations that were based on the (physical / chemical) properties.

  9. CALCULATOR PRACTICE (This is important to master!!!) 6.25 x 103 - 2.01 x 102 = (2.15 x 103)(6.1 x 105)(5.0 x 10-6) = 3.25 x 108 = 3.6 x 107 FYI: “EE” button on calc= typing “X10^” 6.05 x 103 6.6 x 103 9.03

  10. 3.2 Accuracy • the measure of how close a measurement comes to the actual or true value of whatever is measured. • how close a measured value is to the accepted value. Precision • the measure of how close a series of measurements are to one another.

  11. think about a TARGET Accurate and Precise Precise Neither

  12. Percent Error Formula: % Error = accepted value- experimental value x 100 accepted value *always a positive number- indicated by the absolute value sign* You will use this formula when checking the accuracy of your experiment.

  13. Significant Figures – includes all of the digits that are known plus a last digit that is estimated. ! FYI: These rules are IMPORTANT and they will save you many points in the future if you learn them NOW!

  14. Rules for determining Significant Figures1. All non-zero digits are significant. 1, 2, 3, 4, 5, 6, 7, 8, 9

  15. 2. Zeros between non-zero digits are significant. (AKA zero sandwich) 102 7002

  16. 3. Leading zeros (zeros at the beginning of a measurement) are NEVER significant. 00542 0.0152

  17. 4. Trailing zeros (zeros after last integer) are significant only if the number contains a decimal point. 210.0 0.860 210

  18. 5. All digits are significant in scientific notation. 2.1 x 10-5 6.02 x 1023 Time to practice!!

  19. Examples:How many significant digits do each of the following numbers contain: a) 1.2 d) 4600b) 2.0 e) 23.450c) 3.002 f) 6.02 x 1023 2 2 2 5 3 4

  20. BellworkTuesday, September 5th • Please take a reading of the volume in the graduated cylinder and the mass from the triple beam balance.

  21. Exact numbers have unlimited Significant Figures Do not use these when you are figuring out sig figs… Examples: 1 dozen = exactly 12 29 people in this room

  22. Rounding Rules:  5 round up < 5 round down (don’t change) Examples: Round 42.63 to 1 significant digit = Round 61.57 to 3 sig. digs.= Round 0.01621 to 2 = Round 65,002 to 2 sig. digs. = 40 61.6 0.016 65,000

  23. When Adding & Subtracting The measurement with the fewest significant figures to the right of the decimal point determines the number of significant figures in the answer. 107.9 m 45.756 m + 62.1 m = 75.263 m + 1123.93 m = 1199.19 m

  24. When Multiplying & Dividing The measurement with the fewest significant figures determines the number of significant figures in the answer. 3.43 m X 6.4253 m = 45.756 m X 1.2 m = 45.01 m / 2.2 m = 22.0 m 55 m 20. m

  25. Measurement in Lab In lab, you record all numbers you know for sure plus the first uncertain digit. The last digit is estimated and is said to be uncertain but still considered significant. This graduated cylinder has markings to the nearest mL (milliliter) and you will determine volume to the nearest 0.1 mL because that is ONE DIGIT OF UNCERTAINTY.

  26. International System of Units • revised version of the metric system • abbreviated SI All units, their meanings and values can be found on pgs. 63,64,65. Meter (m) – Liter (L) – Gram (g) – SI unit for length SI unit for volume SI unit for mass

  27. Density • Mass – (g) amount of matter in an object • Volume – (mL) amount of space occupied by an object • Density – (g/mL) a ratio of mass to volume

  28. Formula: m v D = Rewrite this formula to solve for m & v! What is the unit for Density???? Remember: A material has the same density no matter how big or small it is!

  29. Example: • A piece of metal has a volume of 4.70 mL and a mass of 57.3 g. What is the density? M = 57.3 g V = 4.70 mL D = M / V D = 57.3 g / 4.70 mL D = 12.19148936 g/mL D = 12.2 g/mL

  30. Trevor performed a lab about density in his chemistry class. He took the following measurements. Calculate Trevor’s density using correct significant figures. 2.9 g/mL 2.8 g/mL 2.8 g/mL

  31. Graphs of density and volume can be used to find the density of a substance. The slope of the line formed when mass and volume are plotted is the density. Remember “rise over run”.

  32. Temperature – measurement of the average kinetic energy of a system.

  33. Temperature Scales Celsius • Sets the freezing point of water at 0C and the boiling point at 100C • Kelvin • Absolute zero is set as the zero on the • Kelvin scale. It is the temperature at • which all motion theoretically ceases.

  34. To convert: K = ºC + 273 (Kelvin does not use “degrees”.) -273 º C = 0 K = absolute zero

  35. Examples: Convert 25 º C to Kelvin. K = ºC + 273 K = 25ºC + 273 = 298 K

  36. Three-step Problem Solving Approach • Analyze – determine how you will find the solution • Calculate – perform the calculation, this may involve measurements • Evaluate– does the answer make sense, and did you use correct units and significant digits

  37. Ex. What is the mass, in grams, of a piece of lead that has a volume of 19.84 cm3? • Analyze: list the knowns and the unknown. Volume = 19.84 cm3Density = mass/ volume Density = 11.4 g/cm3 Mass = ? • Calculate: solve for the unknown. D = m / V so… m = D x V Mass = 19.84 cm3 x 11.4 g/cm3 = 226 g • Evaluate: does the result make sense? Would a piece of lead that is about the size of an eraser have a mass of 226 grams? Yes!

  38. Practice: Solve the following using correct significant figures. 7.823 x 15.76 = 892 + 173 + 56 = 123.3 28459

  39. Practice What is the mass, in grams of a piece of lead that has a volume of 8.73 cm3? • Knowns • V = 8.73 cm3 • D = 11.4 g/cm3 • m = VD • Unknowns • m = ? • To solve: • m = (8.73 cm3)(11.4 g/cm3) • m = • 99.522 g = 99.5 g

  40. Practice: What is the volume, in cm3, of a sample of cough syrup that has a mass of 20.0g and a density of 11.4 g/cm3? • Knowns • M = 20.0 g • D = 11.4 g/cm3 • V = M / D • Unknowns • V = ? • To Solve • V = 20.0 g • 11.4 g/cm3 • V = 1.75 cm3

  41. Dimensional Analysis • To convert from one unit to another, we will use a problem solving method called dimensional analysis. • This method uses equalities or conversion factors to change one unit into another. • Example: If someone gives you 32 quarters, how many dollars do you have? How did you know this?

  42. How many inches are there in 4 ft? • 48 inches • What did you have to know in order to figure that out? • 1 ft = 12 inches • 1 ft = 12 inches is a conversion factor. It can be written as a fraction where the numerator and the denominator are equivalent but have different units. For example, we can use the following conversion factors for changing between inches and feet: • 12 inches1 foot • 1 foot or 12 inches

  43. Some Handy Conversions • Let’s look at a meterstick and list all the conversions we can get from it. • We can say that one meter is equal to… • 10 dm • 100 cm • 1000 mm • Now let’s use those prefixes to figure out how we can modify units for liters. • A liter should contain _____ deciliters • A liter should contain _____ centiliters • A liter should contain _____ milliliters

  44. Practicing with Dimensional Analysis • 70 kg = ____ g • 63 cm = ____ mm • 2.5 L = _____ mL • 1 m2 = _____ cm2

  45. Practice: How many atoms are in 7.00g of gold? (1 atom of gold = 3.271 x 10-22 g) 7.00 g gold 1 atom gold = 2.14 x 1022 atoms gold 3.271 x 10-22 g gold

  46. Some dimensional analysis problems require several steps How many seconds are in 5.0 days? • 5.0 days 24 hours 60 min 60sec = 430,000 sec • 1 day 1 hour 1 min How many students are in a 10 room building if each classroom contains 25 students? 10 rooms 25 students = 250 Students 1 room

  47. An example of this would be the conversion of speed in miles per hour to meters per second. An object was traveling at 400. m/min. What was its speed in cm/s? • 400. m 100 cm 1 min = 667 cm/sec • 1 min 1 m 60sec Convert 423 m/sec to km/min. 423 m 1 km 60sec = 25.4 km/min 1 sec 1000 m 1 min

  48. The density of manganese is 7.21 g/cm3. What is the density of manganese expressed in units of kg/m3? • 7.21 g 1 kg 1003 cm3 = 7210 kg/m3 • 1 cm3 1000g 13 m3

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