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SECTION 2-3

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Objectives 1. Distinguish between accuracy and precision 2. Determine the number of significant figures in measurements 3. Perform mathematical operations involving significant figures 4. Convert measurements into scientific notation 5. Distinguish between inversely and directly proportional relationships

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Accuracy and Precision Accuracy – the closeness of measurements to the correct or accepted value of the quantity measured Precision – the closeness of a set of measurements of the same quantity made in the same way. Visual Concept – Click Here

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Percent Error Calculated by subtracting the experimental value from the accepted value, dividing the difference by the accepted value, and then multiplying by 100 Percent error = value accepted - value experimental x 100 value accepted Sample Problem A student measures the mass and volume of a substance and calculates its density as 1.40 g/mL. The correct, or accepted, value of the density is 1.30 g/mL. What is the percentage error of the student’s measurement? Percent error = value accepted - value experimental value accepted = 1.30 g/mL g/mL 1.30 g/mL = % x 100 Video Concept – Click Here

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Error in Measurement 1. Skill of the measurer 2. Measuring instruments 3. Conditions of measurement

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Significant Figures Consists of all digits known with certainty, plus one final digit 6.75 cm certain estimated Total of 3 significant figures, to two places after the decimal mL **Read from bottom of meniscus.

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Rules for determining significant zeros Visual Concept – Click Here

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Sample Problem How many significant figures are in each of the following measurements? a) 28.6 g 3 s.f. b) cm 4 s.f. c) 910 m 2 s.f. d) L 4 s.f. e) kg 5 s.f.

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Rounding Greater than 5inc. by 42.7 Less than 5stay17.32 17.3 5, followed by nonzero inc. by 2.79 5, not followed by nonzero inc. by 4.64 Preceded by odd digit 5, not followed by nonzero stays78.65 78.6 Preceded by Even digit Visual Concept – Click Here

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Addition/subtraction with significant figures The answer must have the same number of digits to the right of the decimal point as there are in the measurement having the fewest digits to the right of the decimal point So :

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Multiplication/Division with significant figures The answer can have no more significant figures than are in the measurement with the fewest number of significant figures g ÷ mL = g/mL 3 s.f.4 s.f. Should be 3 s.f. So : g/mL

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Sample Problems Carry out the following calculations. Express each answer to the correct number of significant figures. a) 5.44 m m = Answer: 2.84 m b) 2.4 g/mL x mL = Answer: 38 g

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Conversion Factors and Significant Figures There is no uncertainty exact conversion factors. Significant figures are only for measurements, not known values!

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Scientific Notation Numbers are written in the form M x 10 n, where the factor M is a number greater than or equal to 1 but less than 10 and n is a whole number. 65 ooo km M is 6.5 Decimal moved 4 places to left X 10 4 So: 6.5 x 10 4 km Why? Makes very small or large numbers more workable molecules 6.02 x molecules Visual Concept – Click Here

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Extremely small numbers – negative exponent Ex: g 5.67 x g M should be in significant figures Mathematical Operations Using Scientific Notation 1. Addition and subtraction —These operations can be performed only if the values have the same exponent (n factor). a) example: 4.2 × 10 4 kg × 10 3 kg OR

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2. Multiplication —The M factors are multiplied, and the exponents are added algebraically. a) example: (5.23 × 10 6 µm)(7.1 × 10 −2 µm) = (5.23 × 7.1)(10 6 × 10 −2 ) = × 10 4 µm 2 = 3.7 × 10 5 µm 2 3. Division — The M factors are divided, and the exponent of the denominator is subtracted from that of the numerator. a) example: = × 10 3 = 6.7 10 2 g/mol

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Using Sample Problems Analyze Read the problem carefully at least twice and to analyze the information in it. Plan Develop a plan for solving the problem. Compute Substitute the data and necessary conversion factors. Evaluate 1) Check to see that the units are correct. 2) Make an estimate of the expected answer. 3) Expressed in significant figures.

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Sample Problem 1. Analyze mass = kg density = 2.70 g/cm 3 volume = ? 2. Plan conversion factor : 1000 g = 1 kg Calculate the volume of a sample of aluminum that has a mass of kg. The density of aluminum is 2.70 g/cm 3. D V = m D V = m

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3. Compute mass = kg x Round using sig figs × 10 3 cm 3 4. Evaluate cm 3, correct units. correct number of sig figs is three 1 kg 1000 g =3057 g D V = m = 2.70 g/cm 3 = cm 3 Video Concept – Click Here

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Direct Proportions Two quantities if divided by the other gives a constant value If one doubles so does the other y = kx

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Inverse Proportions Two quantities who’s product is a constant If one doubles, the other is cut in half xy = k Visual Concept – Click Here

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