Chemistry Matters and Measurement Explained with Everyday Examples

1. Matters and Measurement Edward Wen, PhD

2. Chemistry is about Everyday experience • Why Cookies tastes different from Cookie Dough? • Why Baking Powder or Baking Soda? • Why using Aluminum Foil, not Paper Towel? • What if the Temperature is set too high? Photo credit: itsnicethat.com A great entertaining documentary where many of us can see how we were just like those chemists: http://www.pbs.org/video/2365543486/

3. Chapter Outline • Classification of matters • Measurement, Metric system (SI) • Scientific Notation • Significant figures • Conversion factor • Density

4. In Your Room • Everything you can see, touch, smell or taste in your room is made of matter. • Chemists study the differences in matter and how that relates to the structure of matter.

5. What is Matter? • Matter: anything that occupies space and has mass • Matter is actually composed of a lot of tiny little pieces: Atoms and Molecules

6. Atoms and Molecules • Atoms: the tiny particles that make up all matter. Helium gas (for blimp) is made up of Helium atoms. • Molecules: In most substances, the atoms are joined together in units. Liquid water is made up of water molecules (2 Hydrogen atoms + 1 Oxygen atoms)

7. Physical States of Matters • Matter can be classified as solid, liquid or gas based on what properties it exhibits (0-l:30) https://www.youtube.com/watch?v=ndw9XYA4iF0

8. Why different States of a Matter? Structure Determines Properties • the atoms or molecules have different structures in solids, liquid and gases

9. Solids • Particles in a solid: packed close together and are fixed in position • though they may vibrate • Incompressible • retaining their shape and volume • Unable to flow

10. Liquids • Particles are closely packed, but they have some ability to move around Incompressible Able to flow, yet not to escape and expand to fill the container (not “antigravity”)

11. Gases • The particles have complete freedom from each other (not sticky to each other) • The particles are constantly flying around, bumping into each other and the container • There is a lot of empty space between the particles (low density)  Compressible  Able to flow and Fill space (“antigravity”)

12. Classifying Matter:Sugar, Copper, Coke, Gasoline/Water

13. Matter Pure Substance Mixture • Constant Composition • Variable Composition • Homogeneous Classification of Matter

14. Pure substance Matter that is composed of only one kind of piece. • Solid: Salt, Sugar, Dry ice, Copper, Diamond • Liquid: Propane, distilled water (or Deionized water, DI water) • Gas: Helium gas (GOODYEAR blimp)

15. Classifying Pure Substances:Elements and Compounds Elements: Substances which can not be broken down into simpler substances by chemical reactions. (A,B) Compounds: Most substances are chemical combinations of elements. (C) • Examples: Pure sugar, pure water • can be broken down into elements • Properties of the compound not related to the properties of the elements that compose it

16. Elements • Example: Diamond (pure carbon), helium gas. • 116 known, 91 are found in nature • others are man-made • Abundance = percentage found in nature • Hydrogen: most abundant in the universe • Oxygen: most abundant element (by mass) on earth and in the human body • Silicon: abundant on earth surface • every sample of an element is made up of lots of identical atoms

18. Compounds • Composed of elements in fixed percentages • water is 89% O & 11% H • billions of known compounds • Organic (sugar, glycerol) or inorganic (table salt) • same elements can form more than one different compound • water and hydrogen peroxide contain just hydrogen and oxygen • carbohydrates all contain just C, H & O (sugar, starch, glucose)

19. Mixture Matter that is composed of different kinds of pieces. Different samples may have the same pieces in different percentages. (D) Examples: • Solid: Flour, Brass (Copper and Zinc), Rock • Liquid: Salt water, soda, Gasoline • Gas: air

20. Classification of Mixtures • Homogeneous = composition is uniform throughout • appears to be one thing • every piece of a sample has identical properties, though another sample with the same components may have different properties • solutions (homogeneous mixtures): Air; Tap water • Heterogeneous = matter that is non-uniform throughout • contains regions with different properties than other regions: gasoline mixed with water; Italian salad dressing

21. What is a Measurement? • Quantitative observation • comparison to an agreed upon standard Every measurement has a numberand a unit: • 77Fahrenheit: Room temperature • 7.5 pounds: Average newborn body weight in the US: • 55 ± 0.5grams: amount of sugar in one can of Coca Cola UNIT: what standard you are comparing your object to the number tells you • what multiple of the standard the object measures • the uncertainty in the measurement (±)

22. Some Standard Units in the Metric System

23. Related Units in the SI System All units in the SI system are related to the standard unit by a power of 10 (exactly!) • 1 kg = 103 g • 1 km = 103 m • 1 m = 102 cm • The power of 10 is indicated by a prefix • The prefixes are always the same, regardless of the standard unit

24. Prefixes: What is the exponent form? • kilo = 1000 times base unit = _____ • 1 kg = 1000 g = 103 g • deci = 0.1 times the base unit = _____ • 1 dL = 0.1 L = 10-1 L; 1 L = 10 dL • centi = 0.01 times the base unit = _____ 1 cm = 0.01 m = 10-2 m; 1 m = 100 cm • milli = 0.001 times the base unit = _____ • 1 mg = 0.001 g = 10-3 g; 1 g = 1000 mg • micro = ______ times the base unit • 1 m = 10-6 m; 106m = 1 m • nano = _______ times the base unit • 1 nL = 10-9L; 109nL = 1 L

25. Common Prefixes in the SI System

26. Standard Unit vs. Prefixes Using meter (m) as example: 1 km = 1000 m = 103 m 1 m = 10 dm = 100 cm = 102cm = 1000 mm = 103mm = 1,000,000 m = 106m = 1,000,000,000 nm = 109 nm

27. Length • Two-dimensional distance an object covers • SI unit: METER (abbreviation as m) • About 3½ inches longer than a yard 1 m = 10-7 the distance from the North Pole to the Equator • Commonly use centimeters (cm) • 1 m = 100 cm = 1.094 yard • 1 cm = 0.01 m = 10 mm • 1 inch = 2.54 cm (exactly)

28. Mass • Amount of matter present in an object • SI unit: kilogram (kg) • about 2 lbs. 3 oz. • Commonly measure mass in grams (g) or milligrams (mg) • 1 kg = 2.2046 pounds (1 lbs. = 0.45359) • 1 g = 1000 mg = 103 mg • 1 g = 0.001 kg = 10-3 kg

29. Volume • Amount of three-dimensional space occupied • SI unit = cubic meter (m3) • Commonly measure solid volume in cubic centimeters (cm3) • 1 m3 = 106 cm3 • 1 cm3 = 10-6 m3 = 0.000001 m3 • Commonly measure liquid or gas volume in milliliters (mL) • 1 gallon (gal) = 3.78 L = 3.78  103 mL • 1 L = 1 dm3 = 1000 mL = 103 mL • 1 mL = 1 cm3 = 1 cc (cubic centimeter)

30. Common Everyday Units and Their EXACT Conversions

31. Common Units and Their Equivalents

32. Units • Always write every number with its associated unit • Always include units in your calculations • you can do the same kind of operations on units as you can with numbers • cm × cm = cm2 • cm + cm = cm • cm ÷ cm = 1 • using units as a guide to problem solving

33. Conversion Factor Relationships to Convert one unit of measurement to another: US dollar  Canadian dollar, dollar  cent Conversion Factors: Relationships between two units Both parts of the conversion factor have the same number of significant figures Conversion factors generated from equivalence statements e.g. 1 inch = 2.54 cm can give or 33

34. How to Use Conversion Factor Arrange conversion factors so starting unit cancels Arrange conversion factor so starting unit is on the bottom of the conversion factor unit 2 unit 1 unit 1 x = unit 2 Conversion Factor 34

35. We have been using the Conversion Factor ALL THE TIME!  How are we converting #cents into #dollars? Why? From 1 dollar = 100 cents Conversion Factor dollar cents 1 dollar 100 cents 45,000 cents x = 450 dollars 35

36. Convert 0.0501 g to ng (nanogram) Convert 325 mL to L (liters) 0.325 L 5.01 x 107 g 36

37. the sun’s diameter is 1,392,000,000 m Scientific Notation Very Large vs. Very Small numbers: The sun’s diameter is 1,392,000,000 m Scientific Notation: 1.392 x 109 m An atom’s diameter is 0.000 000 000 3 m Scientific Notation: 3 x 10-10 m

38. Scientific Notation (SN) Power of 10 (Math language): • 10 x 10 = 100  100 = 102 (2nd power of 10) • 10 x 10 x 10 = 1,000  1,000 = 103 (3rd power of 10) each Decimal Place in our number system represents a different power of 10 • 24 = 2.4 x 101 = 2.4 x 10 • 1,000,000,000 (1 billion) = 109 • 0.0000000001 (1/10 billionth ) = 10-10 Easily comparable by looking at the power of 10

39. exponent 1.23 x 10-8 decimal part exponent part Exponents 10Y • when the exponent on 10 (Y) is positive, the number is that many powers of 10 larger • sun’s diameter = 1.392 x 109 m = 1,392,000,000 m • when Y is negative, the number is that many powers of 10 smaller • avg. atom’s diameter = 3 x 10-10 m = 0.0000000003 m 1.23 x 105 > 4.56 x 102 4.56 x 10-2 > 7.89 x 10-5 7.89 x 1010 > 1.23 x 1010

40. Writing Numbers in SN Big numbers: 12,340,000 Small numbers: 0.0000234 1.234 x 107 2.34 x 10-5

41. Writing a Number in Standard Form 1.234 x 10-6 • since exponent is -6, move the decimal point to the left 6 places • if you run out of digits, add zeros 000 001.234 If the exponent > 1, add trailing zeros: 1.234 x 1010 1.2340000000 0.000 001 234 12,340,000,000

42. Scientific calculators

43. 1.23 Input 1.23 +/- -1.23 Press EXP -1.23 00 Press -1.23 03 Input 3 +/- -1.23 -03 Press Inputting Scientific Notation into a Calculator -1.23 x 10-3 • input decimal part of the number • if negative press +/- key • (–) on some • press EXP key • EE on some (maybe 2nd function) • input exponent on 10 • press +/- key to change exponent to negative

44. Significant Figures (Sig. Fig.) 12.3 cm: 3 sig. figs. range 12.30.1 cm Definition: The non-place-holding digits in a reported measurement • some zero’s are place holders (0.005010), NOT counted as significant figures. What is Sig. Fig. for? the range of values to expect for repeated measurements • the more significant figures there are in a measurement, the smaller the range of values is, more precise. 12.30 cm:4 sig. figs. range 12.30 0.01 cm

45. Significant Figure is SIGNIFICANT FIGURE For the same measurement: More sig figs = more precision Higher precision measurement is like higher resolution of image: High resolution, sharper image

46. Significant Figures (SF) vs. Decimal Places (DP) • All non-zero digits are significant • 1.5 : 2 SF, 1 DP • Interior zeros are significant • 1.05 : 3 SF, 2 DP • Zero: Only Trailing zeros are significant • 1.050 : 4 SF, 3 DP • 0.001050 : 4 SF, 6 DP (Place-holding zero) = SN : 1.050 x 10-3

47. Counting Significant Figures (Contd) 4. Exact numbers has infinite () number of significant figures: example: • 1 pound = 16 ounces • 1 kilogram = 1,000 grams = 1,000,000 milligrams • 1 water molecule contains 2 hydrogen atoms 5. Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation, or add decimal point to specify SF. • 150: ambiguous number • 150. : 3 SF • 1.50 x 102 : 3 SF

48. Example–Counting Sig. Fig. in a Number How many significant figures are in each of the following numbers? 0.0035 1.080 2.97 × 105 1 m = 1000 mm

49. Practice: How many SF and DP in measurement?

50. Sig. Fig. in Multiplication/Division: SF • When multiplying or dividing measurements with Sig. Fig., the result has the same number of SF as the measurement with the fewest number of sig. fig. Round off: 5.02 × 89,665 × 0.10 = 45.0118 = Add trailing zero: 5.00 ÷250. = 0.02 = Use scientific notation for large numbers 5.89 × 6,103 = 35946.67 =