1. Chapter 3 Scientific Measurement Pioneer High School Mr. David Norton

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 x 100% error =

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

24. 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)

25. 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 (m 3 )  Everyday unit = Liter (L), which is non-SI

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

28. 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 o C, which is about normal room temperature

29. 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

31. 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

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

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

35. 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

36. Density  The formula for density is: mass volume Common units are g/mL, or possibly g/cm 3, (or g/L for gas) Density is a physical property, and does not depend upon sample size Density =

38. 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?

39. 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

40. 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

41. 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 o C = 1 g/cm 3

42. Formula D of substance (g/cm 3 ) D of water (g/cm 3 ) Note there are no units left, since they cancel each other Measured with a hydrometer – p.72 Uses? Tests urine, antifreeze, battery Specific gravity =

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

44. 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

45. Temperature scales  Celsius scale- named after a Swedish astronomer  Uses the freezing point(0 o C) and boiling point (100 o C) of water as references  Divided into 100 equal intervals, or degrees Celsius

46. Temperature scales  Kelvin scale (or absolute scale)  Named after Lord Kelvin  K = o C + 273  A change of one degree Kelvin is the same as a change of one degree Celsius  No degree sign is used

47. Temperature scales  Water freezes at 273 K  Water boils at 373 K  0 K is called absolute zero, and equals –273 o C  Fig. 3.19, page 75  Sample 3-6, page 75