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Chapter 3 Scientific Measurement

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1 Chapter 3 Scientific Measurement
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2 Section 3.1 The Importance of Measurement
OBJECTIVES: Distinguish between quantitative and qualitative measurements.

3 Section 3.1 The Importance of Measurement
OBJECTIVES: Convert measurements to scientific notation.

4 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

5 Working with Scientific Notation
Multiplication Multiply the coefficients, add the exponents Division Divide the coefficients, subtract the denominator exponent from numerator exponent

6 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

7 Working with Scientific Notation
Subtraction Line up decimal; subtract coefficients as usual; exponent remains the same

8 Section 3.2 Uncertainty in Measurements
OBJECTIVES: Distinguish among the accuracy, precision, and error of a measurement.

9 Section 3.2 Uncertainty in Measurements
OBJECTIVES: Identify the number of significant figures in a measurement, and in the result of a calculation.

10 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

11 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

12 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 % error = x 100%

13 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

14 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

15 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

16 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

17 Section 3.3 International System of Units
OBJECTIVES: List SI units of measurement and common prefixes.

18 Section 3.3 International System of Units
OBJECTIVES: Distinguish between the mass and weight of an object.

19 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

20 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

21 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)

22 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

23 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)

24 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

25 Volume Measuring Instruments
Graduated cylinders Pipet Buret Volumetric Flask Syringe Fig. 3.12, page 66

26 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

27 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

28 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

29 Section 3.4 Density OBJECTIVES:
Calculate the density of an object from experimental data.

30 Section 3.4 Density OBJECTIVES:
List some useful application of the measurement of specific gravity.

31 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

32 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 =

33 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?

34 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

35 Density and water Sample 3-5, page 71 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

36 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

37 Formula Note there are no units left, since they cancel each other
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 =

38 Section 3.5 Temperature OBJECTIVES:
Convert between the Celsius and Kelvin temperature scales.

39 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

40 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

41 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

42 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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