1 Physical and Chemical Properties of MatterUnit 1 Physical and Chemical Properties of Matter

2 2.1 Properties of Matter

3 Matter Anything that has mass and takes up space

4 Properties Extensive Property: Intensive Property: Substance:Depends on the amount of matter in a sample Examples: Mass and Volume Intensive Property: Depends on the TYPE of matter in a sample Examples: Density Substance: Matter with uniform and definite composition Every sample would have the same Intensive Properties Examples: Water, Hydrogen, Oxygen

5 Practice: Properties Brainstorm1.) On a piece of scratch paper, brainstorm as many properties as you can for each object. 2.) Label these properties as extensive or intensive.

6 What is a Physical Property?Quality or condition that can be observed without changing the composition of the substance Examples: color, solubility, odor, hardness, density, melting point, boiling point, magnetism, etc.

7 States of Matter Solid, liquid, and gas (plasma) AKA Phases of Matter

8 Important Properties of the Physical States of Matter:Property Solid (s) Liquid (l) Gas (g) Shape Definite Indefinite Volume Expansion Very slight Moderate Great Compressibility Poor

9 Vapor vs. Gas Vapor and Gas are interchangeable words.There is a difference though: Gas is used when the substance is in the gaseous state at room temperature (oxygen). Vapor describes the gaseous state of a substance that is generally a liquid or solid at room temperature (water).

10 What is a Physical Change?Altering a material without changing its chemical composition Examples: cut, grind, bend, melt, boil, dissolve, etc.

11 Contrast: What is a Chemical Change?Altering a material by changing its chemical composition Examples: burn, rust, decompose, corrode, explode, etc.

12 2.2 Mixtures

13 Mixtures A physical blend of two or more substances Two types:Heterogeneous: Not uniform in composition Examples: salad, sand, muddy water, oil and water, soup, mixture consisting of two or more phases, etc. Homogeneous: Uniform in composition Examples: air, salt water, soda water, solutions, mixture consisting of one phase AKA solution

14 Phase Describes any part of a sample with uniform composition and properties HOMOGENEOUS mixtures have only 1 phase. HETEROGENEOUS mixtures can be made of 2 or more phases. YOU CAN USUALLY SEE PHASES SEPARATELY!

15 Separating Mixtures Many mixtures can be separated by simple means.Must use Physical Properties Boiling point differences Melting point differences Particle size Density Magnetism Solubility Color Smell

16 Sample Problem 2.1 How could a mixture of aluminum nails and iron nails be separated?

17 2.3 Elements and Compounds

18 Elements Simplest form of matter that can exist under normal laboratory conditions Cannot be separated into simpler substances by chemical means Building blocks for all other substances Examples: nitrogen (N), hydrogen (H), and carbon (C)

19 Compounds Substances that can be separated into simpler substances, but only by chemical means Chemical combination of 2 or more elements Variety of chemical processes to separate substances into elements Compounds can be _________________, but Elements cannot.

20 Remember, what is a Chemical Change?Chemical Change : Altering a material by changing its chemical composition Examples: burn, rust, decompose, corrode, explode, etc.

21 Substances vs. MixturesPure Substance: Any chemical composition that is FIXED or uniform Can be changed to another substance only through “chemical” means Can either be “Element” or “Compound” Mixture: A composition that varies or is made of multiple substances Can be separated by “physical” means Can either be “Heterogeneous” or “Homogeneous”

22 Concept Map Construct a Concept Map using the following terms: MatterSubstance Mixture Element Compound Heterogeneous Homogeneous

23 2.4 Chemical Reactions

24 Chemical Property Quality of how a substance reacts with other substances to form new products The ability of a substance to undergo specific chemical changes Examples: Reactivity, Electronegativity, # of Valence electrons, Family

25 During a chemical change…one or more substances change into new substances Reactant: a starting substance in a chemical reaction Product: a substance formed in a chemical reaction Example: Nitrogen and hydrogen gas can react to form ammonia under certain conditions. Reactants Yield Products N2(g) H2(g) → NH3(g)

26 Chemical Change IndicatorsChange in heat Change in color Change in odor Light is given off Formation of a new gas (be careful) Formation of a solid (precipitate) Irreversible

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28 Conservation of Mass During any chemical reaction, the mass of the products is always equal to the mass of the original reactants

29 Chemistry I Chapter 3 9/16/2016

30 Qualitative vs Quantitative ObservationsQualitative Observations are Descriptive and do not include a numerical measurement Quantitative Observations include a numerical measurement

31 What is the Difference Between Accuracy and Precision?measure of how close a measurement comes to the actual or true value of whatever is measured Example: hitting the bull’s-eye on a dartboard Precision: Measure of how close a series of measurements are to one another Quality of measurements Example: hitting roughly the same spot on a dartboard several times

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33 Precision and MeasurementsHow precise a measurement is depends on the tool used to make the measurement. Precision is determined by the number of quantitative markers on the measuring tool. We always measure/estimate to one decimal place beyond the tools marks.

34 What can/should we record as the length of the object?

35 What can/should we record as the length of the object?

36 What can/should we record as the length of the object?

37 What can/should we record as the volume in this graduated cylinder?

38 Percent Error CalculationsPercent Error (or Experimental Error) The percentage an experimentally obtained value is from what it theoretically should be Percent Error = | theor - Exp | x 100% theor

39 Example: The boiling point of pure water is measured to be 99.1°C. Calculate the percent error.

40 Warm up: Block 2 (9/21 or 9/22)

41 Significant Figures The number of figures (digits or numbers) known in a measurement, including the last digit which should always be estimated Significant Figures are related to PRECISION

42 Every nonzero digit in a reported measurement is assumed to be significant24.7 meters = ________ sig figs meter = ________ sig figs 71 meters = _______ sig figs

43 Zeros appearing between nonzero digits are significant7003 meters = _______ sig figs 40.7 meters = _______ sig figs meters = _______ sig figs

44 Leftmost zeros appearing in front of nonzero digits are not significant. They act as placeholders. By writing the measurements in scientific notation, you can eliminate such placeholding zeros meter or 7.1 x 10-3 meter = ______ 0.422 meter or x 10-1 meter = ______ meter or 9 x 10-5 meter = ______

45 Zeros at the end of a number and to the right of a decimal point are always significant43.00 meters = _______ sig figs 1.10 meters = _______ sig figs 9.0 meters = _______ sig figs

46 Zeros at the rightmost end of a measurement that lie to the left of an understood decimal point are not significant if they serve as placeholders to show the magnitude of the number 300 meters = _______ sig figs 7200 meters = _______ sig figs 27,210 meters = _______ sig figs **If such zeros were known measured values, however, then they would be significant. Writing the value in scientific notation makes it clear that these zeros are significant** SPECIAL CASE 300 meters (if it really should have 3 sig figs) = 3.00 x 102 m

47 There are two situations in which numbers have an unlimited number of significant figures. The first involves counting. A number counted is exact. 23 people in your classroom has _______ sig figs The second situation involves exactly defined quantities such as those found within a system of measurement 60 min = 1 hr has _______ sig figs 100 cm = 1 m has _______ sig figs

48 Sample Problem 3.3 How many significant figures are in each measurement? 123 m d. 22 meter sticks 40,506 mm e m 9.8000x104 m f. 98,000 m

49 Warm-up 9/24/16 Measure the volume of the liquid in the 2 graduated cylinders at the front of the room. Be as precise as possible

50 Why have Significant Figures?All measurements have ERROR. This error needs to be communicated to other scientists These errors get passed on through calculations Our calculations can’t be more precise than the measurements used in them!

51 Rounding to Significant FiguresSample Problem 3.4 (page 69) Round off each measurement to the number of significant figures shown in parentheses. Write answers in scientific notation a meters (four) b meter (two) c meters (two)

52 ADDITION AND SUBTRACTIONThe answer should be rounded to the same number of decimal places as the measurement with the least number of decimal places 25.46 g g g g g = g g = g Why? Answer can only be as precise as the least precise measurement used!

53 MULTIPLICATION AND DIVISIONThe answer should be rounded with the same number of significant figures as the measurement with the least number of significant figures. (1.311m) (2.20m) = 6.884 cm / 2s =

54 MULTIPLICATION AND DIVISIONThe answer should be rounded with the same number of significant figures as the measurement with the least number of significant figures. (1.311m) (2.20m) = m2 = 2.88m2 6.884 cm / 2s =

55 MULTIPLICATION AND DIVISIONThe answer should be rounded with the same number of significant figures as the measurement with the least number of significant figures. (1.311m) (2.20m) = m2 = 2.88m2 6.884 cm / 2s = cm/s = 3 cm/s Why? Errors are compounded in multiplication and division!

56 Journal: Practice Calculate the volume of a rectangular block by using its 3 dimensions Measure the Length, Width, and Height of the block of wood in cm. Calculate the volume of the block of wood in cm3 Calculate the volume of 5 marbles by water displacement in a 50-mL graduated cylinder. Record the initial volume of the water Record the final volume of the water Calculate the volume of the 5 marbles

57 Sample Problem 3.5 Sample Problem 3.6Perform the following addition and subtraction operations. Give each answer to the correct number of significant figures. 12.52 meters meters meters meters – meters Sample Problem 3.6 Perform the following operations. Give the answers to the correct number of significant figures 7.55 meters X 0.34 meters 2.10 meters X 0.70 meters meters2 / 8.4 meters 0.365 meters2 / meter

58 Warm-up 9/26 How many significant figures do each of the following measurements have? g 15.0 cm 150 mL Add 2.5cm+0.45cm+10.01cm= Round the answer where you think you should. Divide 2.45𝑔 11𝑚𝐿 = Round the answer where you think you should. Show 4 different ways to represent the number 200m Each should be a different “precision” of the number.

59 Section 3.2: Student ResponsibleUse Textbook to fill in

60 Some common metric equivalents1 km = 1000 m 1 m = 100 cm 1 m = 1000 mm 1 kg = 1000 g 1 g = 100 cg 1 g = 1000 mg 1 kL = 1000 L 1 L = 100 cL 1 L = 1000 mL 1 cm3 = 1 mL 1 min = 60 s 1 h = 60 min Convert 25 km to m Convert kL to mL Convert L to cm3

61 What is Density and how is it Determined?Density – the ratio of the mass of an object to its volume Density = Mass Volume The density of substances change with temperature, but primarily depends on the substance. Does Density depend on Mass? Does Density depend on Volume?

62 What is Density and how is it Determined?Density – the ratio of the mass of an object to its volume Density = Mass Volume The density of substances change with temperature, but primarily depends on the substance. Does Density depend on Mass? No Does Density depend on Volume?

63 What is Density and how is it Determined?Density – the ratio of the mass of an object to its volume Density = Mass Volume The density of substances change with temperature, but primarily depends on the substance. Does Density depend on Mass? No Does Density depend on Volume? No

64 Warm-up 9/29 Convert 2.5 km/hr into cm/s How many cm3 fit in 1 L?Sample Problem 3.8 A copper penny has a mass of 3.1 g and a volume of cm3. What is the density of copper?

65 Water: Special DensityThe metric system has been designed to give a special value for the density of water Pure Water: 1 g of H2O = 1mL of H2O Density = 1 g/mL

66 Section 3.3: Textbook and in classPage 84: What happens when a measurement is multiplied by a conversion factor? What is Dimensional Analysis?

67 Factor Label Method: Dimensional AnalysisKeys to the system: Different units can be EQUIVALENT (equal to each other) Any quantity divided by its EQUIVALENT is equal to 1 Any quantity multiplied by 1, does not change in true value; it just changes units.

68 Sample Problem 3.9 (page 86) How many seconds are in a workday that lasts exactly eight hours?

69 Sample Problem 3.11 (page 88) Express 750 dg in grams

70 Sample Problem 3.12 (page 89) What is the volume of a pure silver coin that has a mass of 14 g? The density of silver (Ag) is 10.5 g/cm3.

71 Sample Problem 3. 14 (page 91) The density of manganese, a metal, is 7Sample Problem 3.14 (page 91) The density of manganese, a metal, is 7.21 g/cm3. What is the density of manganese expressed in the units of kg/m3?

72 Warm-up 9/30 If the density of Aluminum is 2.70 g/cm3, what will be the mass of a cube of Aluminum that is 6cm x 6cm x 6cm? How much will it cost to drive a car for 8.00 hours if the car averages a speed of 55.0 miles per hour, averages fuel consumption at 32.5 miles per gallon, and the cost of fuel is $2.98 per gallon?

73 Sample Problem: Complex Dimensional Analysis Tina’s car gets 35Sample Problem: Complex Dimensional Analysis Tina’s car gets 35.0 miles per gallon on the freeway. She is driving from San Francisco to Boston (3050 miles away). She will spend an average of $3.97 per gallon of high-octane fuel. Assuming all her driving is on the freeway, how much will her trip cost in gas?

74 Additional Conversions for ChemistryTemperature Celsius to Kelvin # K = # ⁰C + 273 Kelvin to Celsius # ⁰C = # K – 273 Celsius to F # ⁰F = (9/5)# ⁰C + 32

75 Warm-up 10/3 The density of water is 1.00 g/cm3. Express this density in lb/ft3. (1 inch = 2.54cm, 1kg = lbs)