chapter 3 scientific measurement pioneer high school mr. david norton
TRANSCRIPT
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Chapter 3Scientific Measurement
Pioneer High School
Mr. David Norton
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Section 3.1The Importance of
MeasurementOBJECTIVES:Distinguish between quantitative and
qualitative measurements.
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Section 3.1The Importance of
MeasurementOBJECTIVES: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 valuex 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
UnitsOBJECTIVES:List SI units of measurement and
common prefixes.
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Section 3.3International System of
UnitsOBJECTIVES: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
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