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Matter

Matter. Matter : Anything that occupies space and has mass. Physical Properties. Physical Properties : They can be measured and observed without changing the composition or identity of a substance. Examples Odor, Color, Volume, Matter, Density, Melting Point, Boiling Point.

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Matter

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  1. Matter Matter: Anything that occupies space and has mass.

  2. Physical Properties Physical Properties: They can be measured and observed without changing the composition or identity of a substance. Examples Odor, Color, Volume, Matter, Density, Melting Point, Boiling Point

  3. A Further Breakdown: Extensive vs. Intensive Physical Properties • Extensive Properties: depend on amt of substance (mass, volume) • Intensive Properties: do NOT depend on amt of substance (melting point, boiling point)

  4. Chemical Properties • Properties in which there is a change in composition • Reactivity, flammability, etc. • Subdivided into physical and chemical changes

  5. Physical Changes Physical Change: change in physical properties Examples Ice melting, water boiling

  6. Chemical Changes Chemical Changes: Forming new substance(s) Examples Rusting of nails, digestion of food in our stomach, the growth of grass

  7. Practice Classify the following as a physical or chemical change or physical or chemical property: (a) Gallium metal melts in your hand (and in your mouth). (b) A Page is White. (c) Copper sheet acquires a green color over the years. (d) Milk turns sour. (e) Wax is melted over a flame. (f) Propane gas is flammable. (g) Bromine liquid is reddish-brown in color.

  8. Pure Substances: Elements and Compounds • Element: A substance that cannot be separated into simpler substances by chemical means. Example Gold and…? • Compound: A substance composed of atoms of 2 or more elements chemically united in fixed proportions. Example Sodium Chloride and…?

  9. Mixtures Mixture: A combination of 2 or more substances in which the substances retain their identity though no longer seen. Examples Air, Soft Drinks, Wine, Coffee, Water pumped from the Earth. Can you think of anymore…? They can be separated into pure substances: Elements and/or Compounds. They can converted into two or more pure substances.

  10. Mixtures • Homogeneous Mixture: The composition of the mixture, after sufficient stirring, is the same throughout the solution. A homogeneous mixture is called a solution. It has one layer. • Ex: Salt dissolved in water. • Heterogeneous Mixture: The individual components of a mixture remain physically separated and can be seen as separate components. It has more than one layer. • Ex: A glass full of oil and water or sand in a bucket of water.

  11. Practice Classify the following as a pure substance, a homogeneous mixture (solution) or a heterogeneous mixture: (a) Soda (b) Kool-Aid (c) Oil and Vinegar (d) Common Table Salt (Sodium Chloride) (e) A vein of gold embedded in quartz

  12. Separation of Mixtures Distillation: is the process of vaporizing a liquid in a boiling pot and then condensing (gas  liquid) it again where it will collect in another vessel. • Used to separate water from dissolved materials (solid or liquid) • Used to make moon-shine; i.e., separate ethanol from impurities

  13. Simple Distillation

  14. Separation of Mixtures Filtration: the process of causing a liquid-solid heterogeneous mixture to encounter a porous barrier so that the liquid passes through. The solid is left behind. • The liquid that passes through is called the filtrate. • The remaining solid is the residue, or filter cake. There are two purposes for filtrations: (1) to remove solid impurities from a liquid. (2) to separate solid products from a liquid.

  15. Scientific Notation Handling Numbers Associated with Measurements Scientific Notation: Expresses a number as a product of a number between 1 and 10 and the appropriate power of 10. These numbers are very large and very small. They are cumbersome Example: 702,400,000,000,000,000,000 0.00000000000000000000768

  16. Using Scientific Notation • Any number can be represented as the product of a number between 1 and 10 and a power of 10 (either positive or negative). • The decimal point should be placed with a one non-zero number to its left. • The power of 10 depends on the number of places the decimal point is moved and in which direction. • If the decimal point is moved to the left, the power of 10 is positive. If the decimal point is moved to the right, the power of 10 is negative.

  17. Examples • Express 685,000 in scientific notation: • The decimal point must be moved five places to the left • Thus, the decimal point has one non-zero number to its left • 6.85 x 105 • Express 0.00000663 in scientific notation: • The decimal point must be moved six places to the right • Thus, the decimal point has one non-zero number to its left • 6.63 x 10-6 • Try these: • 809,000,000,000 • 0.0000000006

  18. Fundamental SI Units Units: The units part of a measurement tells us what quantity is being used to represent the results of the measurement. • SI = Systeme Internationale (French) Physical QuantityName of UnitAbbreviation mass kilogram kg length meter m time second s temperature kelvin K amount of substance mole mol

  19. Measurements of Length, Volume, and Mass • Length: Measurement of how long a thing is from end to end. • The SI base unit of length is the meter (m). • Volume: Amount of 3-D space occupied by a substance. • Its SI derived unit is m3. • Another common unit of volume is the liter (l). • Mass: Quantity of matter present in an object. • The SI base unit of mass is the kilogram (kg). • Prefixes can be used for all units: • i.e., milligram, milliliter, millimeter

  20. Prefixes used with SI Units PrefixSymbolMeaning Tera T 1 x 1012 Giga G 109 Mega M 106 Kilo k 103 Deca D 101 deci d 10-1 centi c 10-2 milli m 10-3 micro m 10-6 nano n 10-9 pico p 10-12

  21. The Use of Prefixes • 1 dL = 1 x 10-1 L = 0.1 L • 1 mg = 1 x 10-3 g = 0.001 g • 1 km = 1 x 103 m = 1000 m

  22. Uncertainty in Measurement Measurements • 3.00 cm 3.01 cm 3.02 cm • Notice that the first two digits are the same. • These are called the certain numbers. • The third digit is estimated and can vary. • It is called an uncertain number. • Give the certain and uncertain numbers in the following measurements: • 2.509 kg 1.0596 L

  23. Precision & Accuracy • Precision: How well measurements agree with one another • Accuracy: agreement of measurement with accepted (book) value

  24. Practice • A 5-page package of high quality printing paper had its length measured in inches. The measurements obtained were: 11.003, 11.003, 11.004, 11.003, 11.003 • The cover says its length is 11.003 inches. • Do you have “good” or “bad” precision? • What about your accuracy: “good” or “bad”?

  25. More Practice • Five blank writable CD’s had the same piece of music burned on to them. The original CD said that the track was two minutes and thirty-three seconds (2’33”) long. • However, the length of the track on the burned CD’s was the following: 2’15”, 2’15”, 2’15”, 2’15”, 2’15” • Do you have “good” or “bad” precision? • What about your accuracy: “good” or “bad”?

  26. Significant Figures Significant Figures: Numbers recorded in a measurement. (All the certain numbers+the first uncertain number) • The more significant figures (sig figs) in a measurement the greater the precision. • 32.0 is less precise than 32.000000

  27. Guidelines for Using Significant Figures • Nonzero Integers: • Any digit that is not zero is significant. Example 894 has _________ significant figures. 2.341 has _________ significant figures.

  28. Guidelines for Using Significant Figures • Leading Zeros: • Zeros to the left of the first nonzero digit are not significant. • They are used to indicate the placement of the decimal point. Example 0.07 has __________ significant figures. 0.0000048 has __________ significant figures.

  29. Guidelines for Using Significant Figures • Captive Zeros: • Zeros between nonzero digits are significant. Example 707 has ___________ significant figures. 50,001 has __________ significant figures.

  30. Guidelines for Using Significant Figures • Trailing Zeros: • If a number is greater than 1, then all the zeros written to the right of the decimal point count as significant figures. Example 3.0 has __________ significant figures. 30.071 has __________ significant figures. 4.042 has __________ significant figures. 7.0000 has __________ significant figures. 8,500 has __________ significant figures.

  31. Guidelines for Using Significant Figures • Leading, Captive, and Trailing Zeros: • If a number is less than 1, then only the zeros that are at the end of the number, and zeros that are between nonzero digits are significant. Example 0.070 has ___________ significant figures. 0.4006 has ___________ significant figures. 0.00520 has __________ significant figures. 0.0006700 has __________ significant figures.

  32. Guidelines for Using Significant Figures • Exact Numbers: • They are assumed to have an unlimited number of significant figures. • 

  33. Guidelines for Using Significant Figures • Numbers With Trailing Zeroes And No Decimal Point: • For numbers that do not contain decimal points, the measurement is said to be ambiguous. Example 700: 1, 2, or 3 sig figs? Use Scientific Notation: 7x102 has one sig fig. 7.0x102 has two sig figs. 7.00 x 102 has three sig figs. (How many significant figures are in 701? Do you need a decimal pt?)

  34. Rounding Off Numbers:Rules for Rounding Off *We like to reduce our number to fewer digits.* 1. If the digit to be removed is less than 5, then the preceding digit stays the same. When rounding off, use only the first number to the right of the last significant figure. Do not round off sequentially. Example 8.934 rounds off to _________ if we only want 2 sig. figs.

  35. Rounding Off NumbersRules for Rounding Off 2. If the digit to be removed is equal to or greater than 5, then the preceding digit is increased by 1. When rounding off, use only the first number to the right of the last significant figure. Do not round off sequentially. Example 8.627 rounds off to ________ if we only want 3 sig. figs. 0.425 rounds off to ________ if we only want 2 sig. figs.

  36. Rules for Using Significant Figures in Calculations • Addition and Subtraction: • In the answer, the number of sig figs to the right of the decimal point are determined by the lowest number of sig figs to the right of the decimal point given by the measurements. • The measurement is said to be limiting. It limits the number of significant figures in the result. Example 90.442 + 1.1 = 91.542 Rounded Off to 91.5 3.000 - 0.10 = _________ Rounded Off to __________ 1081 - 7.25 = _________ *For Addition and Subtraction, the decimal points are counted as sig figs.*

  37. Rules for Using Significant Figures in Calculations • Multiplication and Division: • The number of sig figs is determined by the original number that has the smallest number of sig figs. • The measurement is said to be limiting. It limits the number of sig figs in the result. Example (2.7)x(3.5029) = 9.45783 Rounded Off to 9.5 (7.85)/(124.6) = _____ Rounded Off to ____________ *For Multiplication and Division, the whole measurements’ sig figs are counted.*

  38. Rules for Using Significant Figures in Calculations • What about: • Order of operations! • Follow the add/sub sig figs for each operation • Then divide, following division sig fig rules • Thus, 7.85 + 11.1 = 19.0 • And 124.6 – 4 = 121 • Therefore, 19.0/121 = 0.157

  39. Problem Solving and Dimensional Analysis • How do we convert from one unit of measurement to another? • We do this via conversion factors. For instance: 1 dollar = 100 pennies Both represent the Same Amount of Money • Conversion factors allow us to carry out conversions between different units that mean the same quantity. • They are not taken into sig fig consideration. • Found on A-11 thru A-13.

  40. Problem Solving and Dimensional Analysis Convert 57.4 m into mm Convert 6.1 dm into km Convert 8.1 m2 to cm2

  41. Problem Solving and Dimensional Analysis Convert 1.06 in. into cm Convert 23.80 L into gal Convert 7.62 g/mL into oz./gal

  42. Comparing Temperature Scales

  43. Temperature Conversions • Converting Between the Kelvin and Celsius Scales ToC + 273.15 = TK • Converting between the Fahrenheit and Celsius Scales ToF = 1.80(ToC) + 32

  44. Temperature Conversions Convert 172 K to oC. Convert 41.2oC to oF. Convert 239.05 oF to K.

  45. Density • Density: Amount of matter present in a given volume of substance • Density = mass/volume = g/mL • Not to be confused with weight!

  46. Example • The volume of a liquid in a graduated cylinder is 24.00 ml, and weighs 36.0 grams. What is the density of this liquid?

  47. Practice • Mercury has a density of 13.6 g/ml. What volume of mercury must be taken to obtain 100 grams of the metal?

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