Water, Water Everywhere

1. Water, Water Everywhere

2. Water as Liquid - Rainwater

3. Water Vapor (Steam)

4. Snow and Snow Flakes

5. Water as Solid - Iceberg

6. The Three States of WaterMacroscopic and Microscopic Views

7. Water Cycle

8. Colors in Nature

9. Chemical Reaction

10. Where does Chemistry fit in? • Chemistry provides the links between the macroscopic world that we experience and the microscopic world of atoms and molecules. • It is relevant to all form of scientific studies.

11. Roles of Chemistry

12. The Central Science • Chemistry is the study of matter and the changes/reactions they undergo. • Chemistry is a central science. • It is essential in understanding both biological and non-biological worlds;

13. What is Matter? • The materials of the universe  anything that has mass and occupies space

14. What Type of Change? • Physical and Chemical; • Physical Changes: Processes that alter the states of substances, but not their fundamental compositions. • Chemical Changees: Processes that alter the fundamental compositions of substances and their identity.

15. Study of Matter & Changes In chemistry you will study: • the existence of matter at macroscopic and microscopic levels; • the different states they can exist, • factors that determine their stability, and • their physical and chemical properties.

16. Atoms vs. Molecules • Matter is composed of tiny particles called atoms. • Atom: smallest part of an element that is still that element. • Molecule: Two or more atoms joined and acting as a unit.

17. Chemical Reaction • One substance becomes another substance(s), such that the fundamental compositions of products are different from those of reactants.

18. Roles of Science • Science is not a just list of facts or knowledge; • Science is a framework for gaining and organizing knowledge/fact about matter, including changes they undergo;

19. Roles of Scientists • Scientists continuously challenge our current beliefs about nature, and always: • asking questions about what we have already known; • Testing the fact/knowledge to confirm it or to gain new insight.

20. The Process: Scientific Method

21. The Scientific Method

22. Fundamental Steps in Scientific Method • Identify the problems and collect information/data; • Develop a hypothesis based on available data; • Test the hypothesis (Design & perform experiments) • Collect and analyze more data to support hypothesis • Make a Conclusion: • Observations may become Law; • Hypotheses may become Theory.

23. Terms in the Scientific Method • Hypothesis: a possible explanation for an observation. • Theory: a set of (tested) hypotheses that gives an overall explanation of certain natural phenomenon. • Scientific Law: a concise statement that summarizes repeatable observed (measurable) behavior.

24. Units and Measurements Measurement • Quantitative observations consist of: • Number & Units (without unit, the value becomes meaningless. • Examples: • 65 kg (kg = kilogram; unit for mass) • 4800 km (km = kilometer; unit for distance) • 3.00 x 108 m/s (m/s = meter per second; unit for speed)

25. Units and Measurements The Number System • Decimal Numbers: 384,400 0.08206 • Scientific Notations: 3.844 x 105 (but 384.4 x 103 is not) 8.206 x 10-2

26. Meaning of 10n and 10-n • The exponent 10n : • if n = 0, 100 = 1; • if n > 0, 10n > 1; • Examples: 101 = 10; 102 = 100; 103 = 1,000; • The exponent 10-n : • if n > 1, 10-n < 1; • Examples: 10-1 = 0.1; 10-2 = 0.01; 10-3 = 0.001

27. Units of Measurements • Units give meaning to numbers. Without UnitWith Units 384,400 ? 384,400 km (implies very far) 384,400 cm (not very far) 144 ? 144 eggs (implies quantity) 0.08206 ? 0.08206 L.atm/(K.mol) (No meaning)

28. English Units Mass: ounce (oz.), pound (lb.), ton; Length: inches (in), feet (ft), yd, mi., etc; Volume: pt, qt, gall., in3, ft3,etc.; Area: acre, hectare, in2, ft2, yd2, mi2.

29. Metric Units Mass: gram (g); kg, mg, mg, ng; Length: meter (m), cm, mm, km, mm, nm, pm; Area: cm2, m2, km2 Volume: L, mL, mL, dL, cm3, m3; (cm3 = mL)

30. Fundamental SI Units Physical QuantityName of UnitAbbreviation Mass kilogram kg Length meter m Time second s Temperature Kelvin K Amount of substance mole mol Energy Joule J Electrical charge Coulomb C Electric current ampere A

31. Prefixes in the Metric System • Prefix Symbol 10n Decimal Forms Giga G 109 1,000,000,000 Mega M 106 1,000,000 kilo k 103 1,000 deci d 10-1 0.1 centi c 10-2 0.01 milli m 10-3 0.001 micro m 10-6 0.000,001 nano n 10-9 0.000,000,001 pico p 10-12 0.000,000,000,001 —————————————————————

32. Mass and Weight • Mass is a measure of quantity of substance; • Mass does not vary with condition or location. • Weight is a measure of the gravitational force exerted on an object; • Weight varies with location if the gravitational force changes.

33. Errors in Measurements • Random errors • values have equal chances of being high or low; • magnitude of error varies from one measurement to another; • error may be minimize by taking the average of several measurements of the same kind;

34. Errors in Measurements • Systematic errors • Errors due to faulty instruments; • reading is either higher or lower than the correct value by a fixed amount; • the magnitude of systematic error is the same, regardless of quantity measured; • For balances with systematic errors, weighing by difference can eliminate systematic errors.

35. Accuracy and Precisionin Measurements • Accuracy The agreement of an experimental value with the “true” or accepted value; • Precision Degree of agreement among values of same measurements; (degree of repeatability)

36. Accuracy and Precision

37. Accuracy and Precision • Accuracy and degree of precision in a measurement is defined by the type of instrument used.

38. Balances with Different Precisions Centigram Balance (precision: ± 0.01 g) Milligram Balance (precision: ± 0.001 g)

39. Analytical Balance(precision: ± 0.0001 g)

40. Significant Figures • Way of expressing measured values with degree of certainty; • For examples: • Mass of an object on a centigram balance = 2.51 g • Mass of same object on analytical balance = 2.5089 g Absolute error for centigram balance = 0.4%; Absolute error for analytical balance = 0.004%; Analytical balance gives mass with more significant figures (5) and more precise (a greater degree of certainty), compared with a centigram balance that gives 3 significant digits for the same mass.

41. How many significant figures are in the following measurements?

42. What is the Buret Reading shown in the Diagram? • Reading liquid volume in a buret; • Read at the bottom of meniscus; • Suppose meniscus is read as 20.15 mL: • Certain digits: 20.15 • Uncertain digit: 20.15

43. What is the volume of liquid in the graduated cylinder?

44. Rules for Counting Significant Figures • All nonzero integers are counted as significant figures Examples: 453.6 has 4 significant figures; 4.48 x 105 has 3 significant figures;

45. Rules for Counting Significant Figures 2. Leading zeroes – zeroes that precede all nonzero digits are NOT counted as significant figures. Examples: 0.0821 has 3 significant figures 0.00055 has 2 significant figures

46. Rules for Counting Significant Figures 3. Captive zeros – these are zeros between nonzero digits; they are always counted as significant figures. Examples: 1.079 has 4 significant figures 0.08206 has 4 significant figures

47. Rules for Counting Significant Figures 4. Trailing zeroes – these are zeroes at the right end of the number. They are counted as significant figures if the number contains a decimal point, otherwise it is not counted. Examples: 208.0 has 4 significant figures; 2080. also has 4 significant figures, but 2080 has 3 significant figures;

48. Rules for Counting Significant Figures 5. Exact numbers – these are numbers given by definition or obtained by counting. They have infinite number of significant figures; the value has no error. Examples: 1 yard = 36 inches; 1 inch = 2.54 cm (exactly); there are 24 eggs in the basket; this class has 60 students enrolled; (There are 35,600 spectators watching the A’s game at the Coliseum is not an exact number, because it is an estimate.)

49. How many significant figures? • 0.00239 • 0.01950 • 1.00 x 10-3 • 100.40 • 168,000 • 0.082060 • One dime equals to 10 pennies • Express 1000 as a value with two significant figures.

50. Rounding off Values in Calculations • In Multiplications and/or Divisions Round off the final answer so that it has the same number of significant figures as the value with the least significant figures. Examples: (a) 9.546 x 3.12 = 29.8 (round off from 29.78352) (b) 9.546/2.5 = 3.8 (round off from 3.8184) (c) (9.546 x 3.12)/2.5 = 12 (round off from 11.913408)