Chapter 2 (read pp. 29-37)

1. Chapter 2(read pp. 29-37) The Scientific Method and Units of Measurement Test is Friday Aug 31st

2. Scientific Method • A logical approach to solving problems by observing and collecting data, formulating hypotheses, testing hypotheses, and formulating theories supported by data. • Observations-using senses to obtain information • Descriptive data – qualitative • Numerical data –quantitative • Examples of qualitative and quantitative:

3. Scientific Method • Inferences – interpretations or explanations • Experiment- carrying out a procedure under controlled conditions to make observations and collect data • Chemists study systems – a specific portion of matter in a given region of space that has been selected for study during an experiment • Examples of a system:

4. Scientific Method • Scientists use generalizations about data collected to formulate hypotheses. • Hypothesis- a testable statement, serves as a basis for further experimenting • only two possible answers • hypothesis is right • hypothesis is wrong • Modify hypothesis - repeat the cycle

5. Cycle repeats many times. • The hypothesis gets more and more certain. • Becomes a theory - A broad generalization that explains a body of facts or phenomenon • A model may be developed to support the theory Observations Hypothesis Experiment

6. Theories are useful because they predict results of new experiments • Help us form mental pictures of processes (models) Observations Hypothesis Experiment

7. Another outcome is that certain behavior is repeated many times • Scientific Law is developed • Description of how things behave • Law - how • Theory- why Observations Hypothesis Experiment

8. Theory (Model) Modify Prediction Experiment Law Observations Hypothesis Experiment

9. The Metric System/SI System An easy way to measure

10. Measurements are quantitative information Units Matter

11. Measuring – number + unit • The numbers are only half of a measurement • Recipe: 1 salt, 3 sugar, 2 flour ??? • Numbers without units are meaningless. • How many feet in a yard • A mile • A rod

12. The Metric System • Easier to use because it is a decimal system • Every conversion is by some power of 10. • A metric unit has two parts • A prefix and a base unit. • prefix tells you how many times to divide or multiply by 10.

13. Prefixes • Tera- T 1,000,000,000,000 1012 • giga- G 1,000,000,000 109 • mega - M 1,000,000 106 • kilo - k 1,000 103 • deci- d 0.1 10-1 • centi- c 0.01 10-2 • milli- m 0.001 10-3 • micro- m 0.000001 10-6 • nano- n 0.000000001 10-9 • pico- p 0.000000000001 10-12

14. Base Units • Length - meter - m • Mass - gram – g • Time - second - s • Temperature – Kelvin - K • Celsius º C • Energy - Joules- J • Volume - Liter - L • Amount of substance - mole – mol

15. Mass • is the amount of matter in an object. • Tool - balance scale • Standard SI unit – kilogram • Base unit - gram • Common units = mg,mg, g, kg • Weight – pull of gravity on matter

16. Length • The distance between two points • Tool – metric ruler • Standard unit - meter • Common units – mm, cm, m, km

17. Derived Units • Many SI units are combinations of base units called derived units • Examples we will use at this time are volume and density

18. Volume • The amount of space an object occupies • V = L x W x H • Tools – metric ruler, graduated cylinder, buret, volumetric flask • SI unit - m3 • 1 Liter = 1 dm3 • 1 mL = 1 cm3 = 1 cc

19. Using Scientific Measurements (pp. 44-52) • All measurements have a certain degree of uncertainty • Uncertainty can result in limitations that depend on the instrument or the experimenter • Scientists use two word to describe how good the measurements are

20. How good are the measurements? • Accuracy- how close the measurement is to the actual value • Precision- how closely the numerical values of a set of measurements agree with each other

21. Differences • Accuracy can be true of an individual measurement or the average of several • Precision requires several measurements before anything can be said about it • There can be precision without accuracy • There can be no accuracy without precision

22. Let’s use a golf anaolgy

23. Accurate? No Precise? Yes

24. Accurate? Yes Precise? Yes

25. Precise? No Accurate? No

26. In terms of measurement • Three students measure the room to be 10.2 m, 10.3 m and 10.4 m across. • Were they precise? • Were they accurate?

27. Percent Error Accuracy is judged using percent error. The formula is: Actual Value – Experimental Value x 100 Actual Value

28. Significant figures (sig figs) • Scientists record measurements in significant figures. • Sig figs consist of all the digits known with certainty plus a final digit that is estimated.

29. 1 2 3 4 5 Significant figures (sig figs) • When using measuring devices, the location of the estimated digit depends on the smallest division on the scale

30. Significant figures (sig figs) • The more marks the better we can estimate. • Scientist always understand that the last number recorded is actually an estimate 1 2 3 4 5

31. Rules for Determining Sig Figs • All nonzero digits are significant • Exact numbers (from counting or definitions) do not limit sig figs • All zeros between nonzero digits are significant

32. Rules for Determining Sig Figs • All zeros to the right of a decimal point and after a nonzero digit are significant • Zeros used for placing the decimal point are not significant

33. Atlantic/Pacific Rule for Determining Sig Figs • If a decimal point is Present, count from the Pacific side • If a decimal point is Absent, count from the Atlantic Side • Begin counting with the first nonzero digit you come to and then keep counting

34. Sig figs. • How many sig figs in the following measurements? • 458 g 3500 g • 4085 g 0.057010 m • 4850 g • 0.0485 g • 0.004085 g • 40.004085 g

35. Sig Figs. • 405.0 g • 4050 g • 0.450 g • 4050.05 g • 0.0500060 g • Next we learn the rules for calculations

36. Adding and subtracting with sig figs • Round the answer so that the estimated digit is in the same place value as the least precise measurement

37. 27.93 + 6.4 27.93 27.93 + 6.4 6.4 For example • First line up the decimal places Then do the adding Find the estimated numbers in the problem 34.33 This answer must be rounded to the tenths place

38. Rounding rules • look at the number behind the one you’re rounding. • If it is 0 to 4 don’t change it • If it is 5 to 9 make it one bigger • round 45.462 to four sig figs • to three sig figs • to two sig figs • to one sig fig

39. Multiplication and Division • The answer should have the same number of significant figures as the measurement with the least number of sig figs • 3.6 x 653 • 2350.8 • 3.6 has 2 s.f. 653 has 3 s.f. • answer can only have 2 s.f. • 2400

40. Practice • 4.8 + 6.8765 • 520 + 94.98 • 0.0045 + 2.113 • 6.0 - 3.82 • 5.4 - 3.28 • 6.7 - .542 • 500 -126 • 6.01 - 3.8

41. Multiplication and Division • 4.5 / 6.245 • 4.5 x 6.245 • 9.8764 x .043 • 3.876 / 1983 • 16547 / 714

42. Homework • Workbook – p. 25 – 26 • # 1,2,3,4,8,10,16

43. Scientific Notation • Shorthand technique used by scientists to write extremely small or large numbers The form is: M x 10n M is a number greater than or equal to 1 but less than 10. The exponent, n, is a positive or negative integer

44. Examples and Practice • 7400 m • 328 500 g • 0.00900 kg • .00705 cm • 0.002 m • 6.3 x 104 cm • 5.42 x 105 g • 12.25 x 102 cm • 6.2 x 10-2 g

45. Dimensional Analysis • A problem solving method that treats units in calculations as algebraic factors • Units common to both numerators and denominators are cancelled and removed from the expressions • A conversion factors is used to convert from one unit to the other • Exact conversions do not limit significant figures

46. Density • D = M / V • An intensive property (it is unaffected by the size of the sample) • Density is often used to identify substances. • Common units - g/ cm3, g/mL, g/L • Tools? -

47. Density • As the mass of the substance increases the volume increases proportionately and the ratio of mass to volume (density) is constant • This is a direct proportion therefore the graph is a straight line that passes through the origin. (See p. 55)

48. Density • Because most substances expand with an increase in temperature (increasing the volume), density usually decreases with increasing volume. • Density varies with temperature

49. Density of water • 1 g of water is 1 mL of water. • density of water is 1 g/mL (at 4ºC) • Specific gravity - the density of an object compared to the density of water • Specific gravity of water is 1.0

50. 0ºC Measuring Temperature • The average kinetic energy of the particles in a sample of matter • Celsius scale • water freezes at 0ºC • water boils at 100ºC • body temperature 37ºC • room temperature 20 - 25ºC