1. DENSITY Video 1.1

2. Density depends on:  Mass: the amount of matter an object contains. (This is different than weight, which is mass plus gravity)  Volume: The amount of space a substance occupies

3. How do we measure mass in the lab?  Electronic Balance

4. How can we measure volume?  l x w x h (regular solid)  ex. V = 1cm 3  Graduated cylinder (liquids)  Read bottom of MENISCUS  ex. V = 27.5 mL

5. Measuring Volume: Irregular Solid Water displacement method: 1. Measure initial volume 2. Measure final volume with object 3. The Difference is the volume of the object

6. Example What is the volume of the solid a. 46 mL b. 54 mL c. 8 mL d. 26 mL

7. Density  Ratio of mass of an object to its volume  Use density formula  Located on Table T

8. Example 1 What is the density of an object with a mass of 60 g and a volume of 2 cm 3 ? 60/2 = 30 g/cm 3

9. Example 2 An object has a volume of 825 cm 3 and a density of 13.6 g/cm 3. Find its mass. 13.6 = x 825 11200g

10. Example 3: How to solve for mass or volume if density is not given: USE TABLE S Example: The volume of an aluminum sample is 251 cm 3. What is the mass of the sample? The density of aluminum on table S is 2.70g/cm 3 2.70 = x 251 678 g

11. Example 4 Determine the volume of an aluminum object with a mass of 40g. 2.70= 40 x 2.70 x = 251 93.0 cm 3

12. METRIC AND SCIENTIFIC NOTATION REVIEW Video 1.2

13. Our country still uses an old system with non uniform measurements such as: fractions of an inch... 12 inches to a foot…. 3 feet to a yard…. 5.5 yards to a rod... 320 rods to a mile... 43,560 sq ft to an acre... But almost all other countries use the metric system, which is disadvantageous for us.

14. But we do use the metric for a few things: We buy cola in liters... We buy memory cards in bites… We run 10 km races... We swim in 25 meter pools... Why haven’t we switched entirely to metric?

15. Measuring length in meters  When measuring a person we would use meters.  If we are measuring an ant, would meters still be feasible? What should we use?  If we are measuring the distance from your house to the school, what should we use?  Always pick a prefix with a value close to what you are measuring.

18. The Metric System  If a unit is getting larger (m  km) the number must get smaller. [If the unit gets smaller (m  cm) the number gets larger.]  Examples:  1. 23.5cm = km  2. 3567mL = L  3. 0.0984g = ug 0.000235 3.567 98400

19. Temperature Conversions  Notice that each scales is marked with BP and FP of water as well as absolute zero.  The degree size of Celsius is equal to Kelvin. Therefore we adjust only for zero points: C = K – 273 K = C + 273

20. Thinking in Celsius  -10° Celsius = frigid (14° F)  0° Celsius = cold (32° F)  10° Celsius = cool (50° F)  20° Celsius = comfortably warm (68° F)  30° Celsius = hot (86° F)  40° Celsius = very hot (104° F)  50° Celsius = Phoenix Hot (120°F)

21. Scientific Notation What is the purpose for using scientific notation in science?

22. Scientific Notation M x 10 n M is between 1 and 10 n is the number of decimal spaces moved to make M

23. Rules 1. Find the decimal point. If it is not written, it is at the end of the number. 2. Move the decimal point to make the number between 1 and 10 3. Place the number of space you moved the decimal in the n spot. 4. If you original number was above 1, the exponent is positive. If the number was smaller that 1, the exponent is negative.

24. Scientific Notation  1020000 is equal to  1.02x10 6  0.00789 is equal to  7.89x10 -3  3.45x10 5 is equal to  345000  1.23x10 -4 is equal to  0.000123

25. Scientific Notation Mathematics Multiplication and Division: 1. Multiply or divide the base numbers. 2. When multiplying, add exponents. When dividing, subtract exponents. (8x10 5 )(2x10 3 ) = 16x10 8 or 1.6x10 9 (8x10 5 )/(2x10 3 ) = 4x10 2

26. Scientific Notation Mathematics Addition and Subtraction: 1. The exponents must be the same. Change your numbers to make this possible. 2. Add or subtract base numbers and do not change the exponent. * Remember: if the decimal move makes the base number smaller, the exponent increases. 5x10 5 + 3x10 4 = 5x10 5 + 0.3x10 5 = 5.3x10 5

27. SIGNIFICANT FIGURES Video 1.3

28. Precision Versus Accuracy  Precision: reproducibility, repeatability  Accuracy: closeness to the correct answer 1. A student obtains the following data: 2.57mL 2.59mL 2.58mL 2.98mL Compare these pieces in terms of precision and accuracy.

29. Precision Versus Accuracy Describe these diagrams in terms of precision and accuracy: The first shows precision, not accuracy. The second shows accuracy, not precision.

30. In this classroom, what is more important: Precision or Accuracy? Think

31. Precision! Due to lack of precise equipment and variable climates we will most likely not end up with accurate results. Therefore, we will focus on refining our lab skills and strive for precise results.

32. Significant Figures  When scientists take measurements their equipment can measure with varying degrees of precision.  A scientists final calculation can only be as precise as their least precise measurement.  Count digits in your measured number to determine their level of precision.

33. Counting Significant Figures  All natural numbers 1-9 count once. 569 has 3SF3.456 has 4SF  The number zero is tricky…  Zeros always count between natural numbers: 109 have 3SF50089 has 5SF  Zeros before a decimal and natural number never count. 0.00789 has 3SF001234 has 4SF  Zeros after natural numbers only count IF there is a decimal present. 100 has 1SF100. has 3SF100.0 has 4SF

34. Adding/Subtracting Significant Figures  Remember: you can only be as precise as your least precise measurement. Therefore, when adding and subtracting, round your answer to the least number of DECIMAL PLACES.  2.0 + 5.61 = 7.61 = 7.6  5.67 + 102.111 = 107.781 = 107.78  23 + 11.10 = 34.10 = 34

35. Multiplying/Dividing Significant Figures  Remember: you can only be as precise as your least precise measurement. Therefore, when multiplying or dividing, round your answer to the least number of significant figures.  2.0 x 35.1 = 70.2 = 70.  5.11 x 98.654 = 504.12194 = 504  72.1 / 3.123 = 23.0867755 = 23.1

36. CLASSIFYING MATTER Video 1.4

37. Matter Anything that takes up space (has a volume) and has mass!

38. Substances A substance has a uniform and definite composition.  Elements are the simplest form of matter which cannot be broken down chemically. They are listed by name on Table S.  Compounds are made up of two or more elements that are chemically combined in a fixed ratio.

39. Mixtures A mixture is comprised of two or more substances.  Homogeneous: a uniform mixture also known as a solution. You can not see its parts. Sometimes gets the symbol (aq) for aqueous which means dissolved in water.  Heterogeneous: non uniform mixture. You can see the parts.

40. Mixtures  Mixtures have two parts:  Solute: The substance(s) dissolving.  Solvent: The substance that does the dissolving. What is the “universal solvent”?

41. Element, Compound or Mixture? Iron Air Water (tap) Water (purified) Gold Methane Sand Sugar Iced tea Steel Chlorine Salt Soap Brass Sodium Calcium carbonate

42. Homogeneous or Heterogeneous?

43. Rules of thumb:  If a single symbol, with one capitol letter is given or can be used to name the species, it is an element.  If two or more combined symbols, with two more more capitol letters are given or can be used to name a species, it is a compound.  All elements and compounds are homogeneous.  If no chemical formula is given or can be found, it doesn’t have a fixed ratio and must be a mixture. These species can be homogeneous or heterogeneous.

44. Separating Mixtures There are many ways to separate a mixture. If it is heterogeneous, you many be able to just sort them with your hands. If not, you can use: 1. Decanting: The least accurate way to separate; you can pour off the top layer of liquid from the bottom layer. (You have probably done this at the beach with water and sand mixtures.)

45. Separating Mixtures 2. Filtering: To separate a solid from a liquid, filter it using a funnel and filter paper. The solid stays on top, liquids pour through. (What happens if the solid is completely dissolved in the liquid?)

46. Separating Mixtures 3. Distillation: When two or more liquids are mixed homogeneously, you can boil off each liquid separately using the following apparatus: This is how many water bottling companies clean their water.

47. Separating Mixtures 4. Chromatography: Mixtures can be dissolved and small samples can be placed on filter paper. The filter paper will be placed in a wet container, and the sample will rise to the top, separating each component out. This method is usually used to determine of the sample is pure or not. It is not used to completely separate the sample.

48. List observations for the solutions:

49. For Fun

50. Solids Liquids and Gases

51. Solids  Definite shape  Definite volume  Constant vibration  Molecules are packed tightly in a geometric (crystalline) pattern

52. Liquids  No definite shape  Definite volume  Constant motion  No arrangement  Molecules are closer together than a gas but further than a solid

53. Gases  No definite shape  No definite volume  No arrangement  Spread out  Compressable

54. Pressure  Gases exert a pressure on surrounding substances because they are constantly moving and colliding with other surfaces.  Only in a vacuum, where there are no molecules, there is no pressure.  Gas pressure can be measured in atmospheres or kilopascals, according to reference table A.

55. 1. Which represents a liquid? Solid? Gas? 2. Which has the highest melting point? 3. Which has the weakest IMF? 4. Which has the lowest boiling point? 5. Which has the strongest IMF? 6. Which has a definite shape? 7. Which have a definite volume?

56. What is a vapor?  Vapors are the gaseous form of a substance that is normally a liquid or solid at room temperature.

57. Heat transfers from areas of high to low temperature  Endothermic reaction require you to put heat in.  Exothermic reactions require you to take heat out; heat exits.

58. Phase Changes Identify the phase change and if it’s endothermic or exothermic:  Evaporation  Condensation  Melting  Freezing  Sublimation  Deposition Liquid to gas endothermic Gas to liquid exothermic Solid to liquid endothermic Liquid to solid exothermic gas to solid exothermic Solid to gas endothermic

59. Describe the following:  Melting point  Boiling Point  Freezing Point MP and FP are the same Temperature for a pure substance since it uses the same phases!

60. Thermochemistry  The study of energy changes that occur in chemical reactions.  Kinetic Energy refers to energy of motion. (Temperature)  Potential Energy refers to stored energy.

61. Phase Change Diagrams Where is the KE increasing? Where is the PE increasing? Where is KE stable? Where is the melting and boiling point?

62. 1. Which line segment represents a liquid? A solid? 2. What is the boiling point? 3. What is the freezing point? 4. On what line segments is the PE increasing? 5. What is PE doing when it is not increasing? 6. Describe KE and PE. 7. Where is the solid/liquid equilibrium? 8. What would happen if it was heated further? B C E D A

63. Cooling Curve Identify the lines for solid, liquid and gas. Where is the KE decreasing? Where is the PE decreasing? Where is PE stable? Where is the melting and boiling point? B E D C A F

64. CHANGES IN THE LAB Video 1.5

65. Changes in the Lab Physical Changes require that the chemical not change composition. (Breaking, grinding, etc.) Chemical Changes require that the chemical changes into a new substance. (Burning, reacting, etc.)

66. Physical or Chemical Changes?

67. Observations Observations are made many ways. They can be either: Qualitative: appearance or behaviors: not measured Quantative: a mathematical description. and either Extensive: dependant on the amount of matter Intensive: dependant on type of matter

68. Qual or Quant and Intensive or Extensive? Rough or smooth Shiny or dull Large or small Kinetic energy