1. 1 Chapter 1 Chemical Foundations

2. 2 Theory or Law Evolution? Atomic system? The germ theory of illness? The heliocentric solar system? Gravity? Heat transfer?

3. 3 Theory or Law 1.Studying more hours will ensure you will get an A+ in Chem101. 2.Solving large number of old exams would increase your grades in Chem101. 3.Drinking a lot of water is better for your health. 4.Doing sport will keep you healthy. 5.Basketball will always move downward. 6.A hot body will cool with time. 7.Conservation of mass

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5. 5 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

6. 6 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

7. 7 Law and Theory A law summarizes what happens; A theory (model) is an attempt to explain why it happens.

8. 8 Figure 1.4: The fundamental steps of the scientific method.

9. 9 Figure 1.5: The various parts of the scientific method.

10. 10 Hypothesis A hypothesis is an educated guess, based on observation. Usually, a hypothesis can be supported or refuted through experimentation or more observation. A hypothesis can be disproven, but not proven to be true. Example: If you see no difference in the cleaning ability of various laundry detergents, you might hypothesize that cleaning effectiveness is not affected by which detergent you use. You can see this hypothesis can be disproven if a stain is removed by one detergent and not another. On the other hand, you cannot prove the hypothesis. Even if you never see a difference in the cleanliness of your clothes after trying a thousand detergents, there might be one you haven't tried that could be different.

11. 11 Theory A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing. A theory is valid as long as there is no evidence to dispute it. Therefore, theories can be disproven. Basically, if evidence accumulates to support a hypothesis, then the hypothesis can become accepted as a good explanation of a phenomenon. One definition of a theory is to say it's an accepted hypothesis.

12. 12 Example of a Theory Example: It is known that on June 30, 1908 in Tunguska, Siberia, there was an explosion equivalent to the detonation of about 15 million tons of TNT. Many hypotheses have been proposed for what caused the explosion. It is theorized that the explosion was caused by a natural extraterrestrial phenomenon, and was not caused by man. Is this theory a fact? No. The event is a recorded fact. Is this theory generally accepted to be true, based on evidence to-date? Yes. Can this theory be shown to be false and be discarded? Yes.

13. 13 Example of a Law A law generalizes a body of observations. At the time it is made, no exceptions have been found to a law. Scientific laws explain things, but they do not describe them. One way to tell a law and a theory apart is to ask if the description gives you a means to explain 'why'. Example: Consider Newton's Law of Gravity. Newton could use this law to predict the behavior of a dropped object, but he couldn't explain why it happened. As you can see, there is no 'proof' or absolute 'truth' in science. The closest we get are facts, which are indisputable observations. Note, however, if you define proof as arriving at a logical conclusion, based on the evidence, then there is 'proof' in science. I work under the definition that to prove something implies it can never be wrong, which is different. If you're asked to define hypothesis, theory, and law, keep in mind the definitions of proof and of these words can vary slightly depending on the scientific discipline. What is important is to realize they don't all mean the same thing and cannot be used interchangeably.

14. 14 Nature of Measurement Measurement - quantitative observation consisting of 2 parts Part 1 - number Part 2 - scale (unit) Part 2 - scale (unit)Examples: 20 grams 6.63    Joule seconds

15. 15 International System (le Système International) Based on metric system and units derived from metric system.

16. 16 The Fundamental SI Units

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18. 18 Figure 1.6: Measurement of volume

19. 19 Figure 1.7: Common types of laboratory equipment used to measure liquid volume.

20. 20 Figure 1.8: An electronic analytical balance.

21. 21 Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

22. 22 Figure 1.9: Measurement of volume using a buret. The volume is read at the bottom of the liquid curve (called the meniscus). 20.16ml 20.17ml 20.15ml 20.18ml ±0.01ml

23. 23 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 elements of the same quantity.

24. 24 Figure 1.10: The results of several dart throws show the difference between precise and accurate.

25. 25 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.

26. 26 Rules for Counting Significant Figures - Overview 1.Nonzero integers 2.Zeros  leading zeros  captive zeros  trailing zeros 3.Exact numbers

27. 27 Rules for Counting Significant Figures - Details Nonzero integers always count as significant figures. 3456 has 4 sig figs.

28. 28 Rules for Counting Significant Figures - DetailsZeros  Leading zeros do not count as significant figures. 0.0486 has 3 sig figs.

29. 29 Rules for Counting Significant Figures - DetailsZeros  Captive zeros always count as  Captive zeros always count as significant figures. 16.07 has 4 sig figs.

30. 30 Rules for Counting Significant Figures - DetailsZeros  Trailing zeros are significant only  Trailing zeros are significant only if the number contains a decimal point. 9.300 has 4 sig figs.

31. 31 Rules for Counting Significant Figures - Details Exact numbers have an infinite number of significant figures. Independent of measuring device: 1 apple, 10 students, 5 cars…. 2πr The 2 is exact 4/3 π r 2 the 4 and 3 are exact From Definition: 1 inch = 2.54 cm exactly The 1 and 2.54 do not limit the significant figures

32. 32 100. has 3 sig. fig. = 1.00 x 10 2 100 has 1 sig. fig. = 1 x 10 2

33. 33 Rules For Rounding 1.In a series of calculations, carry the extra digits through to the final result, then round. 2.If the digit to be removed: A.Is less than 5, then no change e.g. 1.33 rounded to 2 sig. fig = 1.3 B.Is equal or greater than 5, the preceding digit increase by 1 e.g. 1.36 rounded to 2 sig. fig = 1.4

34. 34 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. 6.38  2.0 = 12.76  13 (2 sig figs)

35. 35 Rules for Significant Figures in Mathematical Operations Addition and Subtraction: # decimal places in the result equals the number of decimal places in the least precise measurement. 6.8 + 11.934 = 18.734  18.7 (3 sig figs)

36. 36 Rules for Counting Significant Figures.

37. 37 Dimensional Analysis Proper use of “unit factors” leads to proper units in your answer: 2.54 cm = 1 inch 1 inch/2.54 cm = 1 Unit factor What is the length in inch of 2.85 cm pencil 2.85 (cm) x 1 (inch)/2.54(cm) = 2.85/2.54 = 1.12 in

38. 38 Dimensional Analysis 1.Determine which unit conversion factor(s) are needed 2.Carry units through calculation 3.If all units cancel except for the desired unit(s), then the problem was solved correctly. 1 L = 1000 mL How many mL are in 1.63 L? 1L 1000 mL 1.63 L x = 1630 mL 1L 1000 mL 1.63 L x = 0.001630 L2L2 mL

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40. 40 Scientific Notation The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 6.022 x 10 23 The mass of a single carbon atom in grams: 0.0000000000000000000000199 1.99 x 10 -23 N x 10 n N is a number between 1 and 10 n is a positive or negative integer

41. 41 Scientific Notation 568.762 n > 0 568.762 = 5.68762 x 10 2 move decimal left 0.00000772 n < 0 0.00000772 = 7.72 x 10 -6 move decimal right Addition or Subtraction 1.Write each quantity with the same exponent n 2.Combine N 1 and N 2 3.The exponent, n, remains the same 4.31 x 10 4 + 3.9 x 10 3 = 4.31 x 10 4 + 0.39 x 10 4 = 4.70 x 10 4

42. 42 Temperature Celsius scale =  C Kelvin scale = K Fahrenheit scale =  F

43. 43 Temperature

44. 44 Figure 1.11: The three major temperature scales. 180/100= 9/5

45. 45 TFTF Tc

46. 46 Figure 1.12: Normal body temperature on the Fahrenheit, Celsius, and Kelvin scales.

47. 47 Density Density is the mass of substance per unit volume of the substance:

48. 48 Matter: Anything occupying space and having mass.

49. 49 Classification of Matter Three States of Matter: Solid: rigid - fixed volume and shape Liquid: definite volume but assumes the shape of its container Gas: no fixed volume or shape - assumes the shape of its container

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51. 51 Types of Mixtures Mixtures have variable composition. A homogeneous mixture is a solution (for example, vinegar) A heterogeneous mixture is, to the naked eye, clearly not uniform (for example, a bottle of ranch dressing)

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53. 53 Pure Substances Can be isolated by separation methods:  Chromatography  Filtration  Distillation

54. 54 Figure 1.15a: Paper chromatography of ink. (a) A line of the mixture to be separated is placed at one end of a sheet of porous paper.

55. 55 Figure 1.15b: Paper chromatograph y of ink. (b) The paper acts as a wick to draw up the liquid.

56. 56 Figure 1.15c: Paper chromatography of ink. (c) The component with the weakest attraction for the paper travels faster than the components that cling to the paper.

57. 57 Figure 1.14: Simple laboratory distillation apparatus.

58. 58 Element: A substance that cannot be decomposed into simpler substances by chemical means. Compound: A substance with a constant composition that can be broken down into elements by chemical processes.

59. 59 Figure 1.16: The organization of matter.

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62. 62 Q1. Express the result of this calculation to the correct number of significant figures. A)188.1 B)188 C)188.12 D)190. E)200.

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