1. Chemical Foundations 1

2. Nature of Measurement Part 1 - number Part 2 - scale (unit) Examples: 20 grams 6.63 x 10 -34 Joule seconds Measurement - quantitative observation consisting of 2 parts consisting of 2 parts

3. Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.  Measurements are performed with instruments  No instrument can read to an infinite number of decimal places

4. Ex: Reading a Meterstick. l 2.... I.... I 3....I.... I 4.. cm First digit (known)= 2 2.?? cm Second digit (known)= 0.7 2.7? cm Third digit (estimated) between 0.05- 0.07 Length reported=2.75 cm or2.74 cm or2.74 cm or2.76 cm

5. Rules for Counting Significant Figures - Details Nonzero integers always count as significant figures. 1. Nonzero integers always count as significant figures. 3456 has 4 sig figs.

6. Rules for Counting Significant Figures - Details Note: “leading” means ANY zero that appears before the first nonzero digit, whether the zeros are before OR after a decimal. Zeros - 2. Leading zeros do not count as significant figures. 0.0486 has 3 sig figs.

7. Rules for Counting Significant Figures - Details Zeros - 3. Sandwiched zeros always count as significant figures. 16.07 has 4 sig figs. Note: “sandwiched” means zeros that appears between nonzero digits

8. Rules for Counting Significant Figures - Details Zeros 4. Trailing zeros are significant only if the number contains a decimal point. 9.300 has 4 sig figs. Note: “trailing” means ALL zeros that appear after the last nonzero digit

9. Rules for Counting Significant Figures - Details Exact numbers have an infinite number of significant figures. 5. Exact numbers have an infinite number of significant figures. 1 inch = 2.54 cm, exactly

10. Sig Fig Practice #1 How many significant figures in each of the following? 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs 3.29 x 10 3 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000  2 sig figs

11. Rules for Significant Figures in Mathematical Operations Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in the calculation. #1. Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in the calculation. 6.38 x 2.0 = 12.76  13 (2 sig figs)

12. Sig Fig Practice #2 3.24 m x 7.0 m CalculationCalculator says:Answer 22.68 m 2 23 m 2 100.0 g ÷ 23.7 cm 3 4.219409283 g/cm 3 4.22 g/cm 3 0.02 cm x 2.371 cm 0.04742 cm 2 0.05 cm 2 710 m ÷ 3.0 s 236.6666667 m/s240 m/s 1818.2 lb x 3.23 ft5872.786 lb·ft 5870 lb·ft 1.030 g x 2.87 mL 2.9561 g/mL2.96 g/mL

13. Rules for Significant Figures in Mathematical Operations Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement. #2: Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement. 6.8 + 11.934 =18.734  18.7 (1 decimal place, 3 sig figs)

14. Sig Fig Practice #3 3.24 m + 7.0 m CalculationCalculator says:Answer 10.24 m 10.2 m 100.0 g - 23.73 g 76.27 g 76.3 g 0.02 cm + 2.371 cm 2.391 cm 2.39 cm 713.1 L - 3.872 L 709.228 L709.2 L 1818.2 lb + 3.37 lb1821.57 lb 1821.6 lb 2.030 mL - 1.870 mL 0.16 mL 0.160 mL

15. Rules for Rounding Answers 1.Complete all calculations, then round ONLY the final answer. 2.Identify the correct digit to round (the last sig fig).ex: 18.734 3.Look ONLY at the number immediately to the right of this digit: 18.734 »If this number is 5 or greater, round the last sig fig up. »If this number is less than 5, the last sig fig remains the same. 18.7

16. The Fundamental SI Units (le Système International, SI)

17. SI Units

18. SI Prefixes Common to Chemistry PrefixUnit Abbr.Exponent MegaM10 6 Kilok10 3 Decid10 -1 Centic10 -2 Millim10 -3 Micro  10 -6 Nanon10 -9 Picop10 -12

19. Precision and Accuracy Accuracy refers to the agreement of a particular value with the true value. Precision refers to the degree of agreement among several measurements made in the same manner. Neither accurate nor precise Precise but not accurate Precise AND accurate

20. Types of Error Random Error (Indeterminate Error) - measurement has an equal probability of being high or low. Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration. This can result in measurements that are precise, but not accurate.

21. Error Analysis Practice Ex 1: The data collected when the same sample of silver was weighed five times is as follows: 2.31g, 2.51g, 2.30g, 2.44g, 2.40g The actual mass of the silver is 2.71. Are the student’s measurements accurate? Are they precise? Practice: Section 1.3 & 1.4 # 3, 4, 6.

22. Dimensional Analysis There are times when you need to change the units in which a measurement is expressed. Ex: You might want to convert from hours to minutes. 6.2 hours = ? minutes To do so, you must find the defined relationship between the 2 units. 1 hour = 60 minutes

23. Dimensional Analysis Then create a conversion factor that will cancel the units of your given value.

24. Conversion Factors Fractions in which the numerator and denominator are EQUAL quantities expressed in different units Example: 1 hr. = 60 min Factors: 1 hr. and 60 min 60 min 1 hr. Which one of these conversion factors will cancel the units of our given value, 6.2 hours?

25. Conversion Factors 6.2 hours x 1 hour = ? Minutes 60 min. OR 6.2 hours x 60 min = ? Minutes 1 hour The second conv. factor allows us to cancel the hour units (since “hr” appears in numerator & denominator) so this is the one we want.

26. Multi-step Conversions Sometimes you must use more than one conversion factor. When there isn’t a direct relationship between the 2 units of interest.

27. Multi-step Conversions, cont. How many seconds are in 1.4 days? Unit plan: days hr min seconds Defined Relationships: 1 day = 24 hr 1 hr = 60 min 1 min = 60 s 1.4 days x 24 hr x 60 min x 60 s = 1 day 1 hr 1 min ANSWER: 120,960 s.

28. Complex Conversions Sometimes it is necessary to convert with measurements that involve more than one unit! Ex: convert 60 mi/hr into ft/sec 1 mile=5280 ft 1 hr=60 min 1 min=60 sec 60mi x 5280 ft 1 hr x 1 min = 90 ft/sec hr 1 mi 60 min 60 sec 1

29. Summary: Dimensional Analysis By using dimensional analysis the UNITS ensure that you have the conversion right side up, and the UNITS are calculated as well as the numbers! ASSIGNMENT: Study Guide, Section 1.6, #15-18, 24-26 p 16-17

30. Steps in the Scientific Method 1. Observations -quantitative - qualitative 2.Formulating hypotheses - possible explanation for the observation 3.Performing experiments - gathering new information to decide whether the hypothesis is valid whether the hypothesis is valid

31. Outcomes Over the Long-Term Theory (Model) - A set of tested hypotheses that give an overall explanation of some natural phenomenon. overall explanation of some natural phenomenon. Natural Law - The same observation applies to many different systems different systems - Example - Law of Conservation of Mass

32. Law vs. Theory A law summarizes what happens A theory (model) is an attempt to explain why it happens. A theory (model) is an attempt to explain why it happens.

33. Converting Celsius to Kelvin 33 Kelvins =  C + 273°C = Kelvins - 273

34. Density Is a physical property of matter & can help you identify unknown element samples. Is the amount of mass per volume. Often expressed in g/mL 34

35. Properties of Matter Extensive properties Intensive properties Volume Mass Energy Content (think Calories!) depend on the amount of matter that is present. do not depend on the amount of matter present. Melting point Boiling point Density 35

36. Three Phases 36

37. Phase Differences Solid Solid – definite volume and shape; particles packed in fixed positions. Liquid Liquid – definite volume but indefinite shape; particles close together but not in fixed positions Gas Gas – neither definite volume nor definite shape; particles are at great distances from one another Plasma – high temperature, ionized phase of matter as found on the sun. 37

38. Classification of Matter 38

39. Separation of a Mixture The constituents of the mixture retain their identity and may be separated by physical means. 39

40. Separation of a Mixture The components of dyes such as ink may be separated by paper chromatography. 40

41. Separation of a Mixture By Distillation 41

42. Organization of Matter Matter Matter Mixtures: a) Homogeneous (Solutions) b) Heterogeneous Pure Substances Compounds Elements Elements Atoms NucleusElectrons ProtonsNeutrons Quarks Quarks 42

43. Separation of a Compound Separation of a Compound The Electrolysis of water Water  Hydrogen + Oxygen H 2 O  H 2 + O 2 Reactant  Products Compounds must be separated by chemical means. With the application of electricity, water can be separated into its elements 43