2. Unit 1 Notes

3. 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. Example: centimeter prefix Base unit A centimeter= one hundredth of a meter

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

5. Short cut to Converting how far you have to move on this chart, tells you how far, and which direction to move the decimal place. The box is the base unit, meters, Liters, grams, etc. khDdcm

6. Short cut to Conversion Change 5.6 m to millimeters khDdcm l starts at the base unit and move three to the right. l move the decimal point three to the right 5600

7. MEASUREMENT Accuracy and Precision

8. Types of Measurement Quantitative- use numbers to describe Qualitative- use description without numbers 4 feet extra large Hot 100ºF

9. Scientists prefer Quantitative- easy check. Easy to agree upon, no personal bias. The measuring instrument limits how good the measurement is.

10. Accuracy vs. Precision Accuracy - how close a measurement is to the accepted value Precision - how close a series of measurements are to each other ACCURATE = CORRECT PRECISE = CONSISTENT

11. Accuracy… To determine if data is accurate, ask yourself, “how close is this data point to the actual, accepted, or true value? In short, the closer your data is to the accepted or true value, the more accurate your data is.

12. Precision… To determine if your data is precise, ask yourself, “How close are my data points to one another?” In short, the closer your data points are to one another, the more precise your data points are.

13. Sketch these in your journal… Accurate Precise Not Accurate Precise Not Accurate Not Precise Accurate Not Precise

14. 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.

15. Percent Error Indicates accuracy of a measurement your value accepted value

16. Percent Error A student determines the density of a substance to be 1.40 g/mL. Find the % error if the accepted value of the density is 1.36 g/mL. % error = 2.9 %

17. Measurements When we measure something, we can (and do) always estimate between the smallest marks. 21345

18. Measurements The better the marks, the better we can estimate. Scientist always understand that the last number measured is actually an estimate. 21345

19. Significant Figures Indicate precision of a measurement. Recording significant figures Significant figures in a measurement include the known digits plus a final estimated digit 2.35 cm

20. Significant Figures Counting Sig Figs Count all numbers EXCEPT: Leading zeros -- 0.0025 Trailing zeros without a decimal point -- 2,500

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

22. 4. 0.080 3. 5,280 2. 402 1. 23.50 Significant Figures Counting Sig Fig Examples 4 sig figs 3 sig figs 2 sig figs

23. 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.

24. In science, we deal with some very LARGE numbers: 1 mole = 602,000,000,000,000,000,000,000 particles In science, we deal with some very SMALL numbers: Mass of an electron = 0.000000000000000000000000000000091 kg Scientific Notation

25. Imagine the difficulty of calculating the mass of 1 mole of electrons! 0.000000000000000000000000000000091 kg x 602000000000000000000000 x 602000000000000000000000 ???????????????????????????????????

26. Scientific Notation: A method of representing very large or very small numbers in the form: M x 10 n M x 10 n  M is a number between 1 and 10  n is an integer

27. 2 500 000 000 Step #1: Insert an understood decimal point. Step #2: Decide where the decimal must end up so that one number is to its left up so that one number is to its left Step #3: Count how many places you bounce the decimal point the decimal point 1234567 8 9 Step #4: Re-write in the form M x 10 n

28. 2.5 x 10 9 The exponent is the number of places we moved the decimal.

29. 0.0000579 Step #2: Decide where the decimal must end up so that one number is to its left up so that one number is to its left Step #3: Count how many places you bounce the decimal point the decimal point Step #4: Re-write in the form M x 10 n 12345

30. 5.79 x 10 -5 The exponent is negative because the number we started with was less than 1.

31. Scientific Notation Converting into Sci. Notation: Move decimal until there’s 1 digit to its left. Places moved = exponent. Large # (>1)  positive exponent Small # (<1)  negative exponent Only include sig figs. 65,000 kg  6.5 × 10 4 kg

32. Physical Properties A characteristic that can be observed or measured without changing the identity of the substance

33. States of Matter

34. Solids  very low kinetic energy - particles vibrate but can’t move around  Retains size and shape  Definite shape and volume

35. States of Matter Liquids  low kinetic energy - particles can move around but are still close together  Takes the shape of its container  definite volume

36. States of Matter Gases  high KE - particles can separate and move throughout container  Easily compressed  No definite shape  No definite volume

37. Some Properties of Solids, Liquids, and Gases Property Solid Liquid Gas Shape Has definite shapeTakes the shape of Takes the shape the container of its container Volume Has a definite volumeHas a definite volume Fills the volume of the container Arrangement of Fixed, very closeRandom, close Random, far apart Particles Interactions between Very strong Strong Essentially none particles

38. Physical Change A change in a substance that does not involve a change in the identity of the substance. Examples: Phase Changes – boiling point, melting point, freezing point A substance dissolving in another substance - solubility

39. Chemical Properties Relates to a substance’s ability to undergo change that transforms it into a different substance  Ability to : combust, oxidize, neutralize, reactivity, etc.

40. Chemical Change A change in which one or more substances are converted into different substances. Evidence of Chemical Change:  Heat and light  Change in color  Production of gas  Precipitation of a solid

41. Temperature is Average Kinetic Energy Fast Slow “HOT” “COLD”

42. Celcius & Kelvin Temperature Scales

43. Measuring Temperature Kelvin starts at absolute zero (-273 º C) degrees are the same size C = K -273 K = C + 273 Kelvin is always bigger. Kelvin can never be negative. 273 K