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Chapter 2: Measurement and Calculations… Section 2-1: Scientific Method (pg29-31) will not be explicitly covered but used throughout this entire class… 1

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Section 2-2: Units of Measurement & Metric review and summary… Pgs 33-39 2

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The Fundamental Metric System (SI) Units…some of them

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SI Units 4

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SI Prefixes Common to Chemistry PrefixUnit Abbr.Exponent Kilok10 3 Decid10 -1 Centic10 -2 Millim10 -3 Micro 10 -6 We will use these units primarily in this class, but these are only a few of the units within the SI (Metric System) 5

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Density… Density is defined as the ratio of mass (matter) to volume. Density (D) is usually expressed D = m/v Units are most commonly g/ml or g/cm 3, but can be in ANY units of mass to volume! Density is a characteristic physical property that does not change with size of sample (BUT some conditions can affect density value…like what?) 6

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Conversion Factors (pg 40-42) We’ll cover this soon…not yet. 7

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(p. 44 - 57) Section 2-3: Using Scientific Measurements 8

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A. Accuracy vs. Precision Accuracy - how close a measurement is to the accepted value Precision - how close a series of measurements are to each other ACCURATE = CORRECT PRECISE = CONSISTENT 9

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B. Percent Error Indicates accuracy of a measurement your value accepted value 10

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B. Percent Error A student determines the density of a substance to be 1.40 g/mL. Find the % error if the accepted value of the density is 1.36 g/mL. % error = 2.9 % 11

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C. Significant Figures Indicate precision of a measurement. Recording measurements in Sig Figs… – Sig figs in a measurement include the known digits plus a final estimated digit 2.35 cm 12

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C. Significant Figures Counting Sig Figs (textbook Table 2-5, p.47) – Zeros that are in between any digits 1-9 are ALWAYS counted as significant Ex: 105 = 3 sig figs 1208.5 = 5 sig figs – Count all numbers (1 though 9) EXCEPT… Leading zeros -- 0.0025 Trailing zeros without a decimal point -- 2,500 13

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4. 0.080 3. 5,280 2. 402 1. 23.50 C. Significant Figures Counting Sig Fig Examples 1. 23.50 2. 402 3. 5,280 4. 0.080 4 sig figs 3 sig figs 2 sig figs 14

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C. Significant Figures Calculating with Sig Figs – Multiply/Divide - The # with the fewest sig figs determines the # of sig figs in the answer. (13.91g/cm 3 )(23.3cm 3 ) = 324.103g 324 g 4 SF3 SF 15

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C. Significant Figures Calculating with Sig Figs (con’t) – Add/Subtract - The # with the lowest decimal value determines the place of the last sig fig in the answer. 3.75 mL + 4.1 mL 7.85 mL 224 g + 130 g 354 g 7.9 mL 350 g 3.75 mL + 4.1 mL 7.85 mL 224 g + 130 g 354 g 16

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C. Significant Figures Calculating with Sig Figs (con’t) – Exact Numbers do not limit the # of sig figs in the answer. Counting numbers: 12 students Exact conversions: 1 m = 100 cm “1” in any conversion: 1 in = 2.54 cm 17

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C. Significant Figures 5. (15.30 g) ÷ (6.4 mL) Practice Problems = 2.390625 g/mL 18.1 g 6. 18.9g - 0.84 g 18.06 g 4 SF2 SF 2.4 g/mL 2 SF 18

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D. Scientific Notation (and standard notation) PROPER Sci. notation uses format: M x 10 n – M is larger than 1 and smaller than 10 (1.000 to 9.999) To converting into Sci. Notation: – Move decimal until there’s 1 digit to left of decimal (COUNTING places moved tells you about the exponent) – Large numbers (>1) positive exponents Small numbers (<1) negative exponent – Only include sig figs. (significant figures) 65,000 kg 6.5 × 10 4 kg 19

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D. Scientific Notation 2,400,000 g 0.00256 kg 7 10 -5 km 6.2 10 4 mm Practice Problems 2.4 10 6 g 2.56 10 -3 kg 0.00007 km 62,000 mm 20

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D. Scientific Notation Calculating with Sci. Notation (5.44 × 10 7 g) ÷ (8.1 × 10 4 mol) = 5.44 EXP EE ÷ ÷ EXP EE ENTER EXE 78.1 4 = 671.6049383= 670 g/mol= 6.7 × 10 2 g/mol Type on your calculator: 21

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E. Proportions Direct Proportion Inverse Proportion y x y x 22

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