1. CHAPTER 2 Metric System

2. THE METRIC SYSTEM

4. Measuring The numbers are only half of a measurement. It is 10 long. 10 what? Numbers without units are meaningless. How many feet in a yard? A mile? A rod?

5. The Metric System Easier to use because it is a decimal system. Every conversion is by some power of 10. A metric unit has two parts. A prefix and a base unit. prefix tells you how many times to divide or multiply by 10.

6. Lengthmeterm Massgramg TimeSeconds TemperatureKelvinK Amount of a substance MoleMol VolumeLiterL Length - straight distance between two points -Meters (m) Mass - how much matter in an object -grams (g) Volume - amount of space taken up by an object Cubic meters (m3) or -Liters (L)

7. Base Units Length - meter - more than a yard - m Mass - grams - about a raisin - g Time - second - s Temperature - Kelvin or ºCelsius K or ºC Energy - Joules- J Volume - Liter - half of a two liter bottle- L Amount of substance - mole - mol

8. Metric System Prefixes convert the base units into units that are appropriate for the item being measured. © 2009, Prentice-Hall, Inc.

9. Kilo-Hecta-Deka-UNITSDeci-Centi-Mili- khda dcm Prefix

10. Prefixes giga- G 1,000,000,00010 9 mega - M 1,000,00010 6 kilo - k 1,00010 3 deci-d0.1 10 -1 centi-c0.01 10 -2 milli-m0.001 10 -3 micro-  0.000001 10 -6 nano-n0.000000001 10 -9

11. Prefixes kilo k 1000 times deci d 1/10 centi c 1/100 milli m 1/1000 kilometer - about 0.6 miles centimeter - less than half an inch millimeter - the width of a paper clip wire

12. DIMENSIONAL ANALYSIS Using the units to solve problems

13. Chapter 1: Chemistry: Matter and Measurement 13 A Problem-Solving Method Chemistry problems usually require calculations, and yield quantitative (numerical) answers For example, 1 inch = 2.54 cm EOS The unit-conversion method is useful for solving most chemistry problems – the focus here is on “unit equivalents”

14. 4m = ________ cm 5km = ____________m 40mm = __________m 67m = __________ hm 135cm = _____________km 0.1km = _________dm 2km = ________________mm.34km = ________ cm 5km = ____________m 19cm = __________mm 98m = __________km 135m = _____________km 33 km = _________dm 87m = ________________mm

15. Converting how far you have to move on this chart, tells you how far, and which direction to move the decimal place. The box is the base unit, meters, Liters, grams, etc. khDdcm

16. Conversions convert 25 mg to grams convert 0.45 km to mm convert 35 mL to liters It works because the math works, we are dividing or multiplying by 10 the correct number of times. khDdcm

17. Chapter 1: Chemistry: Matter and Measurement 17 Other Equivalents and Conversion Factors A conversion factor is the fractional expression of the equivalents EOS

18. Dimensional Analysis We use dimensional analysis to convert one quantity to another. Most commonly dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm) © 2009, Prentice-Hall, Inc. 1 in. 2.54 cm 1 in. or

19. Dimensional Analysis Use the form of the conversion factor that puts the sought-for unit in the numerator. © 2009, Prentice-Hall, Inc. Given unit  desired unit desired unit given unit Conversion factor

20. Dimensional Analysis For example, to convert 8.00 m to inches, convert m to cm convert cm to in. © 2009, Prentice-Hall, Inc. 8.00 m 100 cm  1 m  1 in. 2.54 cm  315 in.

21. Dimensional Analysis Use conversion factors to change the units Conversion factors = 1 1 foot = 12 inches (equivalence statement) 12 in = 1 = 1 ft. 1 ft. 12 in 2 conversion factors multiply by the one that will give you the correct units in your answer.

22. Chapter 1: Chemistry: Matter and Measurement 22 Two Examples EOS How many cm are in 26 inches? 26 in× cm in 2.54 1 = 66 cm

23. Examples 11 yards = 2 rod 40 rods = 1 furlong 8 furlongs = 1 mile The Kentucky Derby race is 1.25 miles. How long is the race in rods, furlongs, meters, and kilometers? A marathon race is 26 miles, 385 yards. What is this distance in rods and kilometers?

24. Units to a Power How many m 3 is 1500 cm 3 ? 3 1500 cm 3 1 m 100 cm 1 m 100 cm 1 m 100 cm 3 1500 cm 3 1 m 100 cm 3

25. Units to a Power How many cm 2 is 15 m 2 ? 36 cm 3 is how many mm 3 ?

26. Multiple units The speed limit is 65 mi/hr. What is this in m/s? 1 mile = 1760 yds 1 meter = 1.094 yds 65 mi hr 1760 yd 1 mi1.094 yd 1 m1 hr 60 min 1 min 60 s

27. Multiple units Lead has a density of 11.4 g/cm 3. What is this in pounds per quart? 454 g = 1 lb 1 L = 1.094 qt

28. Scientific Notation To write in Scientific Notation you need a number between 1 & 9 in front of the decimal. When going from right to left you add the exponent (positive exponent) When going from left to right you subtract the exponent (negative exponent)

29. Converting Cont. Examples I. 345 II..000345 III. 56890 IV..000000000134

30. WHICH IS HEAVIER? it depends

31. Density How heavy something is for its size The ratio of mass to volume for a substance D = M/V Independent of how much of it you have Gold- high density Air- low density

32. Floating Lower density floats on higher density Ice is less dense than water Most wood is less dense than water Helium is less dense than air. A ship is less dense than water

33. Density of Water 1g of water is 1mL of water Density of water is 1 g/mL At 4 0 C Otherwise it is less

34. Calculating The formula tells you how Units will be g/ml or g/cm 3 A piece of wood has a mass of 11.2g and a volume of 23mL. What is the density? A piece of wood has a density of 0.93 g/mL and a volume of 23 mL what is the mass?

35. UNCERTAINY IN MEASUREMENT © 2009, Prentice-Hall, Inc.

36. Significant Figures The term significant figures refers to digits that were measured. When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers. © 2009, Prentice-Hall, Inc.

37. Chapter 1: Chemistry: Matter and Measurement 37 Significant Figures EOS All digits in a number that are known with certainty plus the first uncertain digit The more significant digits obtained, the better the precision of a measurement The concept of significant figures applies only to measurements Exact values have an unlimited number of significant figures

38. Significant Figures 1. All nonzero digits are significant. 2. Zeroes between two significant figures are themselves significant. 3. Zeroes at the beginning of a number are never significant. 4. Zeroes at the end of a number are significant if a decimal point is written in the number. © 2009, Prentice-Hall, Inc.

39. Chapter 1: Chemistry: Matter and Measurement 39 Rules for Zeros in Significant Figures Zeros between two other significant digits ARE significant e.g., 10023 A zero preceding a decimal point is not significant e.g., 0.10023 EOS Zeros between the decimal point and the first nonzero digit are not significant e.g., 0.0010023

40. Chapter 1: Chemistry: Matter and Measurement 40 Rules for Zeros in Significant Figures Zeros at the end of a number are significant if they are to the right of the decimal point e.g., 0.10023001023.00 EOS Zeros at the end of a number may or may not be significant if the number is written without a decimal point e.g., 1000. compared to 1000

41. Chapter 1: Chemistry: Matter and Measurement 41 Rules for Significant Figures in Calculations KEY POINT: A calculated quantity can be no more precise than the least precise data used in the calculation EOS Analogy: a chain is only as strong as its weakest link … and the reported result should reflect this fact

42. Significant Figures When addition or subtraction is performed, answers are rounded to the least significant decimal place. When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation. © 2009, Prentice-Hall, Inc.

43. Chapter 1: Chemistry: Matter and Measurement 43 Significant Figures in Calculations Multiplication and Division: the reported results should have no more significant figures than the factor with the fewest significant figures 1.827 m × 0.762 m = ? EOS 0.762 has 3 sigfigs so the reported answer is 1.39 m 2

44. Chapter 1: Chemistry: Matter and Measurement 44 Significant Figures in Calculations Addition and Subtraction: the reported results should have the same number of decimal places as the number with the fewest decimal places EOS NOTE - Be cautious of round-off errors in multi- step problems. Wait until calculating the final answer before rounding.

45. Conservation of Mass Law of Conservation of Mass- in a physical or chemical reaction, mass is neither created nor destroyed; it is conserved. All mass can be accounted for. Mass of the Reactants = Mass of Products

46. Weight vs. Mass Weight Mass Measures the force of gravity on an object Weight can change if the force of gravity acting on the object changes How much matter is in an object Remains constant (the same) no matter where it is

47. Mass and Weight Mass is measure of resistance to change in motion Weight is force of gravity. Sometimes used interchangeably Mass can’t change, weight can

48. Mass Weight is a force. Mass is the amount of matter. 1 gram is defined as the mass of 1 cm 3 of water at 4 ºC. 1000 g = 1000 cm 3 of water 1 kg = 1 L of water

49. Mass 1 kg = 2.5 lbs 1 g = 1 paper clip 1 mg = 10 grains of salt

50. Volume calculated by multiplying L x W x H Liter the volume of a cube 1 dm (10 cm) on a side 1L = 1 dm 3 so 1 L = 10 cm x 10 cm x 10 cm 1 L = 1000 cm 3 1/1000 L = 1 cm 3 1 mL = 1 cm 3

51. Volume 1 L about 1/4 of a gallon - a quart 1 mL is about 20 drops of water or 1 sugar cube

52. Measurement How do we measure in science? What do measurements mean? How do we tell? Let’s start with Accuracy vs. Precision

53. How good are the measurements? Scientists use two words to describe how good the measurements are- Accuracy- how close the measurement is to the actual value Precision- how well can the measurement be repeated

54. Uncertainty Basis for significant figures All measurements are uncertain to some degree Precision- how repeatable Accuracy- how correct - closeness to true value. Random error - equal chance of being high or low- addressed by averaging measurements - expected

55. Accuracy versus Precision Accuracy refers to the proximity of a measurement to the true value of a quantity. Precision refers to the proximity of several measurements to each other. © 2009, Prentice-Hall, Inc.

56. Uncertainty Systematic error- same direction each time Want to avoid this Bad equipment or bad technique. Better precision implies better accuracy You can have precision without accuracy You can’t have accuracy without precision (unless you’re really lucky).

57. Temperature A measure of the average kinetic energy Different temperature scales, all are talking about the same height of mercury. Derive a equation for converting ºF toºC

58. Temperature is different from heat. Temperature is which way heat will flow. (from hot to cold) Heat is energy, ability to do work. A drop of boiling water hurts, kilogram of boiling water kills.

59. Calculating Temp. From Celsius to Fahrenheit F= (C x 9/5) + 32 From Fahrenheit to Celsius C= 5/9 (F-32) Examples: I. 49 0 F to 0 C II. 97 0 C to 0 F

60. Measuring Temperature Celsius scale. water freezes at 0ºC water boils at 100ºC body temperature 37ºC room temperature 20 - 25ºC 0ºC

61. Measuring Temperature Kelvin starts at absolute zero (-273 º C) degrees are the same size C = K -273 K = C + 273 Kelvin is always bigger. Kelvin can never be negative. 273 K

62. Problems How many???? 1. 349K to 0 C 2. 12 0 C to K 3. 34 0 F to 0 C 4. 101 0 C to 0 F

63. Units of heat are calories or Joules 1 calorie is the amount of heat needed to raise the temperature of 1 gram of water by 1ºC. A food Calorie is really a kilocalorie. How much energy is absorbed to heat 15 grams of water by 25ºC. 1 calorie = 4.18 J

64. Conservation of Mass Law of Conservation of Mass- in a physical or chemical reaction, mass is neither created nor destroyed; it is conserved. All mass can be accounted for. Mass of the Reactants = Mass of Products

65. Energy The ability to do work. Work - cause a change or move an object. Many types- all can be changed into the other.

66. Types of energy Potential- stored energy Kinetic Energy- energy something has because its moving Heat- the energy that moves because of a temperature difference. Chemical energy- energy released or absorbed in a chemical change. Electrical energy - energy of moving charges

67. Types of Energy Radiant Energy- energy that can travel through empty space (light, UV, infrared, radio) All types of energy can be converted into others. If you trace the source far enough back, you will end up at nuclear energy.

68. Conservation of Energy Energy can be neither created or destroyed in ordinary changes (not nuclear), it can only change form. Its not just a good idea, its the law.