Measurement
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Transcript of Measurement
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Measurement
• Quantitative Observation
• Comparison Based on an Accepted Scale– e.g. Meter Stick
• Has 2 Parts – the Number and the Unit– Number Tells Comparison– Unit Tells Scale
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Scientific Notation
• Technique Used to Express Very Large or Very Small Numbers
• Based on Powers of 10
• To Compare Numbers Written in Scientific Notation– First Compare Exponents of 10– Then Compare Numbers
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Writing Numbers in Scientific Notation1 Locate the Decimal Point2 Move the decimal point to the right of the non-
zero digit in the largest place– The new number is now between 1 and 10
3 Multiply the new number by 10n
– where n is the number of places you moved the decimal point
4 Determine the sign on the exponent n– If the decimal point was moved left, n is +
– If the decimal point was moved right, n is –
– If the decimal point was not moved, n is 0
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Writing Numbers in Standard Form
1 Determine the sign of n of 10n
– If n is + the decimal point will move to the right– If n is – the decimal point will move to the left
2 Determine the value of the exponent of 10– Tells the number of places to move the decimal
point
3 Move the decimal point and rewrite the number
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Related Units in the Metric System
• All units in the metric system are related to the fundamental unit by a power of 10
• The power of 10 is indicated by a prefix
• The prefixes are always the same, regardless of the fundamental unit
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Table 2.2: The Commonly used Prefixes in the Metric System
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Length• SI unit = meter (m)
– About 3½ inches longer than a yard• 1 meter = one ten-millionth the distance from the North Pole to the
Equator = distance between marks on standard metal rod in a Paris vault = distance covered by a certain number of wavelengths of a special color of light
• Commonly use centimeters (cm)
– 1 m = 100 cm
– 1 cm = 0.01 m = 10 mm
– 1 inch = 2.54 cm (exactly)
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Table 2.3: The Metric System for Measuring Length
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Figure 2.1: Comparison of English and metric units for length on a ruler
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Volume• Measure of the amount of three-dimensional space occupied by a
substance• SI unit = cubic meter (m3)• Commonly measure solid volume in cubic centimeters (cm3)
– 1 m3 = 106 cm3 – 1 cm3 = 10-6 m3 = 0.000001 m3
• Commonly measure liquid or gas volume in milliliters (mL)– 1 L is slightly larger than 1 quart– 1 L = 1 dm3 = 1000 mL = 103 mL – 1 mL = 0.001 L = 10-3 L– 1 mL = 1 cm3
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Figure 2.2: Cubes
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Figure 2.3: A 100-ml Graduated Cylinder
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Mass
• Measure of the amount of matter present in an object
• SI unit = kilogram (kg)• Commonly measure mass in grams (g) or
milligrams (mg)– 1 kg = 2.2046 pounds, 1 lbs.. = 453.59 g– 1 kg = 1000 g = 103 g, 1 g = 1000 mg = 103 mg– 1 g = 0.001 kg = 10-3 kg, 1 mg = 0.001 g = 10-3 g
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Figure 2.4: An electronic analytical
balance used in chemistry labs
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Uncertainty in Measured Numbers
• A measurement always has some amount of uncertainty
• Uncertainty comes from limitations of the techniques used for comparison
• To understand how reliable a measurement is, we need to understand the limitations of the measurement
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Reporting Measurements
• To indicate the uncertainty of a single measurement scientists use a system called significant figures
• The last digit written in a measurement is the number that is considered to be uncertain
• Unless stated otherwise, the uncertainty in the last digit is ±1
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Figure 2.5: Measuring a Pin
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Rules for Counting Significant Figures
• Nonzero integers are always significant• Zeros
– Leading zeros never count as significant figures– Captive zeros are always significant– Trailing zeros are significant if the number has
a decimal point
• Exact numbers have an unlimited number of significant figures
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Exact Numbers• Exact Numbers are numbers known with certainty • Unlimited number of significant figures• They are either
– counting numbers• number of sides on a square
– or defined• 100 cm = 1 m, 12 in = 1 ft, 1 in = 2.54 cm
• 1 kg = 1000 g, 1 LB = 16 oz
• 1000 mL = 1 L; 1 gal = 4 qts.
• 1 minute = 60 seconds
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Calculations with Significant Figures
• Calculators/computers do not know about significant figures!!!
• Exact numbers do not affect the number of significant figures in an answer
• Answers to calculations must be rounded to the proper number of significant figures– round at the end of the calculation
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Rules for Rounding Off
• If the digit to be removed• is less than 5, the preceding digit stays the
same• is equal to or greater than 5, the preceding
digit is increased by 1
• In a series of calculations, carry the extra digits to the final result and then round off
• Don’t forget to add place-holding zeros if necessary to keep value the same!!
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Multiplication/Division with Significant Figures
• Result has the same number of significant figures as the measurement with the smallest number of significant figures
• Count the number of significant figures in each measurement
• Round the result so it has the same number of significant figures as the measurement with the smallest number of significant figures
4.5 cm x 0.200 cm = 0.90 cm2
2 sig figs 3 sig figs 2 sig figs
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Adding/Subtracting Numbers with Significant Figures
• Result is limited by the number with the smallest number of significant decimal places
• Find last significant figure in each measurement
• Find which one is “left-most”• Round answer to the same decimal place
450 mL + 27.5 mL = 480 mLprecise to 10’s place precise to 0.1’s place precise to 10’s place
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Problem Solving and Dimensional Analysis
• Many problems in chemistry involve using equivalence statements to convert one unit of measurement to another
• Conversion factors are relationships between two units– May be exact or measured– Both parts of the conversion factor should have the same
number of significant figures
• Conversion factors generated from equivalence statements– e.g. 1 inch = 2.54 cm can give or
in1cm54.2
cm54.2in1
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• Arrange conversion factors so starting unit cancels– Arrange conversion factor so starting unit is on
the bottom of the conversion factor
• May string conversion factors
Problem Solving and Dimensional Analysis
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Converting One Unit to Another
• Find the relationship(s) between the starting and goal units. Write an equivalence statement for each relationship.
• Write a conversion factor for each equivalence statement.
• Arrange the conversion factor(s) to cancel starting unit and result in goal unit.
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Converting One Unit to Another
• Check that the units cancel properly
• Multiply and Divide the numbers to give the answer with the proper unit.
• Check your significant figures
• Check that your answer makes sense!
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Temperature Scales• Fahrenheit Scale, °F
– Water’s freezing point = 32°F, boiling point = 212°F
• Celsius Scale, °C– Temperature unit larger than the Fahrenheit
– Water’s freezing point = 0°C, boiling point = 100°C
• Kelvin Scale, K– Temperature unit same size as Celsius
– Water’s freezing point = 273 K, boiling point = 373 K
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Figure 2.6: Thermometers based on the three temperature scales in (a) ice water and (b) boiling water
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Figure 2.7: The three major temperature scales
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Figure 2.8: Converting 70°C to units measured on the Kelvin scale
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Figure 2.9: Comparison of the Celsius and Fahrenheit scales
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Density• Density is a property of matter representing the mass per unit
volume
• For equal volumes, denser object has larger mass
• For equal masses, denser object has small volume
• Solids = g/cm3
– 1 cm3 = 1 mL
• Liquids = g/mL
• Gases = g/L
• Volume of a solid can be determined by water displacement
• Density : solids > liquids >>> gases
• In a heterogeneous mixture, denser object sinks
VolumeMass
Density
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Using Density in Calculations
VolumeMass
Density
DensityMass
Volume
Volume Density Mass
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Figure 2.10: (a) Tank of water. (b) Person submerged in the tank, raising the level of the water.