2. Chemical Foundations Chapter 1

3. Chemistry Chemistry deals with situations in which the nature of a substance is changed by altering its composition so that entirely new substances are synthesized or particular properties of existing substances are enhanced.

4. Science Science is both a noun and a verb. Science is a body of knowledge and a method of adding to that body of knowledge.

5. Steps in the Scientific Method 1.Observations - quantitative - measurement involves a number and a unit. -  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. 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. Law vs. Theory A law summarizes what happens; a theory (model) is an attempt to explain why it happens.

8. The various parts of the scientific method.

9. Problems of the Scientific Method Scientists must be objective when using the scientific method. The scientific method is affected by: profit motivesreligious beliefs warsmisinterpretation of data budgetsemotions fadsprejudices politicspeer pressure

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

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

12. The Fundamental SI Units

14. One liter is defined as a cubic decimeter and 1 mL is one cubic centimeter.

15. Common types of laboratory equipment used to measure liquid volume.

16. Mass & Weight Mass is a measure of the resistance of an object to a change in its state of motion -- a constant. Weight is the measure of the pull of gravity on an object and varies with the object’s location.

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

18. Measurement of volume using a buret. The volume is read at the bottom of the meniscus.

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 elements of the same quantity.

20. a) is neither precise nor accurate, b) is precise but not accurate (small random, large systematic errors) c) both precise and accurate (small random, no systematic errors.

21. 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.

22. Accuracy Sample Exercise 1.2 on page 13. TrialGraduated CylinderBuret 125 mL26.54 mL 225 mL26.51 mL 325 mL26.60 mL 425 mL26.49 mL 525 mL26.57 mL Average25 mL26.54 mL Which is more accurate? Graduated cylinder produces systematic error --value is too low. Buret

23. Exponential Notation Also called scientific notation and powers of ten notation. Exponential notation has two advantages: the number of significant digits can easily be indicated fewer zeros are needed to write a very large or very small number.

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

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

26. Rules for Counting Significant Figures - Details Zeros - Leading zeros do not count as significant figures. 0.0486 has 3 sig figs.

27. Rules for Counting Significant Figures - Details Zeros - Captive zeros always count as significant figures. 16.07 has 4 sig figs.

28. Rules for Counting Significant Figures - Details Zeros -  Trailing zeros are significant only if the number contains a decimal point. 9.300 has 4 sig figs.

29. Rules for Counting Significant Figures - Details Exact numbers have an infinite number of significant figures. Can come from counting or definition. 15 atoms 1 inch = 2.54 cm, exactly

30. 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)

31. Rules for Significant Figures in Mathematical Operations Addition and Subtraction: # sig figs 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)

32. 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 five, the preceding digit stays the same. b. is equal to or greater than five, the preceding digit is increased by 1.

33. Dimensional Analysis Also called unit cancellation is a method of solving problems by using unit factors to change from one unit to another. Unit factor -- the unit that you have goes on bottom, and the unit that you want goes on top.

34. Dimensional Analysis Proper use of “unit factors” leads to proper units in your answer.

35. Dimensional Analysis What is the dimension of a 25.5 in bicycle frame in centimeters? (25.5 in)(2.54 cm/1 in) = 64.8 cm Units must be cancelled and the answer must have correct sig figs, be underlined, and include proper units!!

36. Temperature Celsius scale =   C Kelvin scale = K Fahrenheit scale =   F

37. Three major temperature scales.

38. Temperature

39. Temperature Calculations Convert - 40.0 o C to Kelvin. K = C + 273.15 K = -40.0 + 273.15 K = 233.2 K

40. Temperature Calculations Convert - 40.0 o C to Fahrenheit. 100 F - 3200 = -7200 100 F = -4000 F = - 40.0 o F

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

42. Density Calculations If an object has a density of 0.7850 g/cm 3 and a mass of 19.625 g, what is its volume? V = 25.00 cm 3

43. Matter: Anything occupying space and having mass.

44. 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

45. 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)

46. HOMOGENEOUS MATTER - a substance with the same properties throughout -- a pure substance. Elements and compounds are pure substances (homogeneous matter).

47. HETEROGENEOUS MATTER - has different properties throughout -- a mixture. »Salt and pepper »soil »granite »sea water »spaghetti & meat balls

48. SEPARATION OF MIXTURES - mixtures can be separated into pure substances by physical means. »distillation »filtration »centrifuging »magnet »evaporation »chromatography

49. Simple laboratory distillation apparatus.

50. CENTRIFUGE

52. Paper Chromatography Chromatography has two phases of matter: a stationary phase (the paper) and a mobile phase ( the liquid).

53. Compounds & Elements 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.

54. Universe Matter Energy Homogeneous Physical Change Heterogeneous Pure Substance SolutionMixture ElementCompound Chemical Change Electron Levels Nucleus Electrons ProtonsNeutrons Potential Energy Kinetic Energy Position Composition Gravitational Electrostatic

55. “TO BUILD FROM MATTER IS SUBLIMELY GREAT, BUT GODS AND POETS ONLY CAN CREATE.” Pitt