Chapter 3 Scientific Measurement

1. Chapter 3Scientific Measurement

2. Section 3.1The Importance of Measurement • OBJECTIVES: • Distinguish between quantitative and qualitative measurements.

3. Section 3.1The 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 3.2Uncertainty in Measurements • OBJECTIVES: • Distinguish among the accuracy, precision, and error of a measurement.

9. Section 3.2Uncertainty 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 decimalplaces 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 3.3International System of Units • OBJECTIVES: • List SI units of measurement and common prefixes.

18. Section 3.3International 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 3.4Density • OBJECTIVES: • Calculate the density of an object from experimental data.

30. Section 3.4Density • 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 • 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 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 3.5Temperature • 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