Chapter 3 Scientific Measurement

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Chapter 3Scientific Measurement

• Charles Page High School
• Dr. Stephen L. Cotton

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Section 3.1The Importance of Measurement

• OBJECTIVES
• Distinguish between quantitative and qualitative
  measurements.

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Section 3.1The Importance of Measurement

• OBJECTIVES
• Convert measurements to scientific notation.

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Measurements

• Qualitative measurements - words
• Quantitative measurements involves numbers
  (quantities)
• Depends on reliability of instrument
• Depends on care with which it is read
• Scientific Notation
• Coefficient raised to power of 10

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Working with Scientific Notation

• Multiplication
• Multiply the coefficients, add the exponents
• Division
• Divide the coefficients, subtract the denominator
  exponent from numerator exponent

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Working with Scientific Notation

• Before adding or subtracting in scientific
  notation, the exponents must be the same
• Calculators will take care of this
• Addition
• Line up decimal add as usual the coefficients
  exponent stays the same

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Working with Scientific Notation

• Subtraction
• Line up decimal subtract coefficients as usual
  exponent remains the same

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Section 3.2Uncertainty in Measurements

• OBJECTIVES
• Distinguish among the accuracy, precision, and
  error of a measurement.

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Section 3.2Uncertainty in Measurements

• OBJECTIVES
• Identify the number of significant figures in a
  measurement, and in the result of a calculation.

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Uncertainty in Measurements

• Need to make reliable measurements in the lab
• Accuracy how close a measurement is to the true
  value
• Precision how close the measurements are to
  each other (reproducibility)
• Fig. 3.4, page 54

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Uncertainty in Measurements

• Accepted value correct value based on reliable
  references
• Experimental value the value measured in the
  lab
• Error the difference between the accepted and
  experimental values

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Uncertainty in Measurements

• Error accepted experimental
• Can be positive or negative
• Percent error the absolute value of the error
  divided by the accepted value, times 100
• error
• accepted value

x 100
error
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Significant Figures(sig. figs.)

• Significant figures in a measurement include all
  of the digits that are known, plus a last digit
  that is estimated.
• Note Fig. 3.6, page 56
• Rules for counting sig. figs.?
• Zeroes are the problem
• East Coast / West Coast method

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Counting Significant Fig.

• Sample 3-1, page 58
• Rounding
• Decide how many sig. figs. Needed
• Round, counting from the left
• Less than 5? Drop it.
• 5 or greater? Increase by 1
• Sample 3-2, page 59

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Sig. fig. calculations

• Addition and Subtraction
• The answer should be rounded to the same number
  of decimal places as the least number in the
  problem
• Sample 3-3, page 60

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Sig. Fig. calculations

• Multiplication and Division
• Round the answer to the same number of
  significant figures as the least number in the
  measurement
• Sample 3-4, page 61

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Section 3.3International System of Units

• OBJECTIVES
• List SI units of measurement and common prefixes.

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Section 3.3International System of Units

• OBJECTIVES
• Distinguish between the mass and weight of an
  object.

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International System of Units

• The number is only part of the answer it also
  need UNITS
• Depends upon units that serve as a reference
  standard
• The standards of measurement used in science are
  those of the Metric System

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International System of Units

• Metric system is now revised as the International
  System of Units (SI), as of 1960
• Simplicity and based on 10 or multiples of 10
• 7 base units
• Table 3.1, page 63

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International System of Units

• Sometimes, non-SI units are used
• Liter, Celsius, calorie
• Some are derived units
• Made by joining other units
• Speed (miles/hour)
• Density (grams/mL)

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Length

• In SI, the basic unit of length is the meter (m)
• Length is the distance between two objects
  measured with ruler
• We make use of prefixes for units larger or
  smaller
• Table 3.2, page 64

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Common prefixes

• Kilo (k) 1000 (one thousand)
• Deci (d) 1/10 (one tenth)
• Centi (c) 1/100 (one hundredth)
• Milli (m) 1/1000 (one thousandth)
• Micro (?) (one millionth)
• Nano (n) (one billionth)

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Volume

• The space occupied by any sample of matter
• Calculated for a solid by multiplying the length
  x width x height
• SI unit cubic meter (m3)
• Everyday unit Liter (L), which is non-SI

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Volume Measuring Instruments

• Graduated cylinders
• Pipet
• Buret
• Volumetric Flask
• Syringe
• Fig. 3.12, page 66

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Volume changes?

• Volume of any solid, liquid, or gas will change
  with temperature
• Much more prominent for GASES
• Therefore, measuring instruments are calibrated
  for a specific temperature, usually 20 oC, which
  is about normal room temperature

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Units of Mass

• Mass is a measure of the quantity of matter
• Weight is a force that measures the pull by
  gravity- it changes with location
• Mass is constant, regardless of location

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Working with Mass

• The SI unit of mass is the kilogram (kg), even
  though a more convenient unit is the gram
• Measuring instrument is the balance scale

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Section 3.4Density

• OBJECTIVES
• Calculate the density of an object from
  experimental data.

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Section 3.4Density

• OBJECTIVES
• List some useful application of the measurement
  of specific gravity.

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Density

• Which is heavier- lead or feathers?
• It depends upon the amount of the material
• A truckload of feathers is heavier than a small
  pellet of lead
• The relationship here is between mass and volume-
  called Density

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Density

• The formula for density is
• mass
• volume
• Common units are g/mL, or possibly g/cm3, (or g/L
  for gas)
• Density is a physical property, and does not
  depend upon sample size

Density
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Things related to density

• Note Table 3.7, page 69 for the density of corn
  oil and water
• What happens when corn oil and water are mixed?
• Why?
• Will lead float?

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Density and Temperature

• What happens to density as the temperature
  increases?
• Mass remains the same
• Most substances increase in volume as temperature
  increases
• Thus, density generally decreases as the
  temperature increases

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Density and water

• Water is an important exception
• Over certain temperatures, the volume of water
  increases as the temperature decreases
• Does ice float in liquid water?
• Why?
• Sample 3-5, page 71

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Specific Gravity

• A comparison of the density of an object to a
  reference standard (which is usually water) at
  the same temperature
• Water density at 4 oC 1 g/cm3

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Formula

• D of substance
  (g/cm3)
• D of water
  (g/cm3)
• Note there are no units left, since they cancel
  each other
• Measured with a hydrometer p.72
• Uses? Tests urine, antifreeze, battery

Specific gravity
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Section 3.5Temperature

• OBJECTIVES
• Convert between the Celsius and Kelvin
  temperature scales.

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Temperature

• Heat moves from warmer object to the cooler
  object
• Glass of iced tea gets colder?
• Remember that most substances expand with a temp.
  increase?
• Basis for thermometers

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Temperature scales

• Celsius scale- named after a Swedish astronomer
• Uses the freezing point(0 oC) and boiling point
  (100 oC) of water as references
• Divided into 100 equal intervals, or degrees
  Celsius

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Temperature scales

• Kelvin scale (or absolute scale)
• Named after Lord Kelvin
• K oC 273
• A change of one degree Kelvin is the same as a
  change of one degree Celsius
• No degree sign is used

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Temperature scales

• Water freezes at 273 K
• Water boils at 373 K
• 0 K is called absolute zero, and equals 273 oC
• Fig. 3.19, page 75
• Sample 3-6, page 75