1. Introduction: Matter & Measurement AP Chemistry Chapter 1

2. Chemistry What is chemistry? composition of matter It is the study of the composition of matter and the changes that matter undergoes. What is matter? It is anything that takes up space and has mass.

3. Elements, Compounds & Mixtures A substance is matter that has a definite composition and constant properties. It can be an element or a compound.

4. Elements, Compounds & Mixtures An element is the simplest form of matter. It cannot be broken down further by chemical reactions.

5. Elements, Compounds & Mixtures A compound can be separated into simpler forms. It is a combination of two or more elements.

6. Mixtures A mixture is a physical blend of two or more substances. 1. Heterogeneous Mixtures –Not uniform in composition –Properties indefinite & vary –Can be separated by physical methods

7. Mixtures 2. Homogeneous Mixtures –Completely uniform in composition –Properties constant for a given sample –Cannot be separated by physical methods (need distillation, chromatography, etc) –Also called solutions.

8. Separating mixtures Only a physical change- no new matter Filtration- separate solids from liquids with a barrier. Distillation- separate different liquids or solutions of a solid and a liquid using boiling points. –Heat the mixture. –Catch vapor and cool it to retrieve the liquid. Chromatography- different substances are attracted to paper or gel, so move at different speeds.

9. Filtration

10. Distillation

11. Chromatography

13. Physical & Chemical Properties Physical property – characteristics of a pure substance that we can observe without changing the substance; the chemical composition of the substance does not change.

14. Physical & Chemical Properties Chemical property – describes the chemical reaction of a pure substance with another substance; chemical reaction is involved.

15. Physical & Chemical Properties Physical properties appearance odor melting point boiling point hardness density solubility conductivity Chemical properties reaction with oxygen (flammability) rxn with water rxn with acid Etc….

16. Intensive & Extensive Properties Intensive properties Do not depend on the amount of sample being examined –temperature –odor –melting point –boiling point –hardness –density Extensive properties Depend on the quantity of the sample –mass –volume Etc….

17. Physical & Chemical Changes Physical changes The composition of the substance doesn’t change Phase changes (like liquid to gas) –Evaporation, freezing, condensing, subliming, etc. Tearing or cutting the substance Chemical changes The substance is transformed into a chemically different substance All chemical reactions

18. Signs of a Chemical Changes 1.permanent color change 2.gas produced (odor or bubbles) 3.precipitate (solid) produced 4.light given off 5.heat released (exothermic) or absorbed (endothermic)

19. Making Measurements A measurement is a number with a unit.A measurement is a number with a unit. All measurements, MUST have units.All measurements, MUST have units.

20. Types of Units Energy Joule J Pressure Pascal Pa

21. Prefixes giga- G 1,000,000,00010 9 mega - M 1,000,00010 6 kilo - k 1,00010 3 deci-d0.110 -1 centi-c0.01 10 -2 milli-m0.001 10 -3 micro-  0.000001 10 -6 nano-n0.000000001 10 -9 pico-p0.00000000000110 -12

22. Measurements There are two types of measurements:  Qualitative data are words, such as color, heavy or hot.  Quantitative measurements involve numbers (quantities), and depend on: 1. The reliability of the measuring instrument. 2. The care with which it is read – this is determined by YOU!

23. Accuracy & Precision  Accuracy – how close a measurement is to the true value.  Precision – how close the measurements are to each other (reproducibility).

24. Precision and Accuracy Neither accurate nor precise Precise, but not accurate Precise AND accurate Our goal!

25. Which are Precise? Accurate?

26. Uncertainty in Measurements Measurements are performed with instruments, and no instrument can read to an infinite number of decimal places Which of the balances below has the greatest uncertainty in measurement? 12 3

27. Uncertainty Basis for significant figures All measurements are uncertain to some degree Precision- how repeatable Accuracy- how correct - closeness to true value. Random error - equal chance of being high or low- addressed by averaging measurements - expected

28. Uncertainty Systematic error- same direction each time Want to avoid this Bad equipment or bad technique. Better precision implies better accuracy. You can have precision without accuracy. You can’t have accuracy without precision (unless you’re really lucky).

29. Significant Figures in Measurements Significant figures in a measurement include all of the digits that are known, plus one more digit that is estimated. Sig figs help to account for the uncertainty in a measurement.

30. To how many significant figures can you measure this pencil? What is wrong with this ruler? What is it missing?

31. Rules for Counting Significant Figures Non-zeros always count as significant figures: 3456 has 4 significant figures

32. Rules for Counting Significant Figures Zeros Leading zeroes do not count as significant figures: 0.0486 has 3 significant figures

33. Rules for Counting Significant Figures Zeros Captive zeroes always count as significant figures: 16.07 has 4 significant figures

34. Rules for Counting Significant Figures Zeros Trailing zeros are significant only if the number contains a written decimal point: 9.300 has 4 significant figures

35. Rules for Counting Significant Figures Two special situations have an unlimited (infinite) number of significant figures: 1.Counted items a)23 people, or 36 desks 2.Exactly defined quantities b)60 minutes = 1 hour

36. Sig Fig Practice #1 How many significant figures in the following? 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs 3.29 x 10 3 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000 mL  2 sig figs 3 cats  infinite These all come from some measurements This is a counted value

37. Significant Figures in Calculations  In general a calculated answer cannot be more accurate than the least accurate measurement from which it was calculated.  Sometimes, calculated values need to be rounded off.

38. Rounding Calculated Answers  Rounding  Decide how many significant figures are needed  Round to that many digits, counting from the left  Is the next digit less than 5? Drop it.  Next digit 5 or greater? Increase by 1

39. Rules for Significant Figures in Mathematical Operations  Addition and Subtraction  The answer should be rounded to the same number of decimal places as the least number of decimal places in the problem.

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

41. Sig Fig Practice #2 3.24 m + 7.0 m CalculationCalculator says:Answer 10.24 m 10.2 m 100.0 g - 23.73 g 76.27 g 76.3 g 0.02 cm + 2.371 cm 2.391 cm 2.39 cm 713.1 L - 3.872 L 709.228 L709.2 L 1818 lb + 3.37 lb1821.37 lb 1821 lb 2.030 mL - 1.870 mL 0.16 mL 0.160 mL *Note the zero that has been added.

42. Rounding Calculated Answers  Multiplication and Division  Round the answer to the same number of significant figures as the least number of significant figures in the problem.

43. Rules for Significant Figures in Mathematical Operations Multiplication and Division: # sig figs in the result equals the number in the least accurate measurement used in the calculation. 6.38 x 2.0 = 12.76  13 (2 sig figs)

44. Other Special Cases What if your answer has less significant figures than you are supposed to have?What if your answer has less significant figures than you are supposed to have? –Calculator Example: 100.00 / 5.00 = 20 Add zeros!Add zeros! –20 is 1 sf –20. is 2 sf –20.0 is 3 sf

45. Sig Fig Practice #3 3.24 m x 7.0 m CalculationCalculator says:Answer 22.68 m 2 23 m 2 100.0 g ÷ 23.7 cm 3 4.219409283 g/cm 3 4.22 g/cm 3 0.02 cm x 2.371 cm 0.04742 cm 2 0.05 cm 2 710 m ÷ 3.0 s 236.6666667 m/s240 m/s 1818.2 lb x 3.23 ft5872.786 lb·ft 5870 lb·ft 1.030 g x 2.87 mL 2.9561 g/mL2.96 g/mL

46. Dimensional Analysis Using the units to solve problems

47. Dimensional Analysis Use conversion factors to change the units Conversion factors = 1 1 foot = 12 inches (equivalence statement) 12 in = 1 = 1 ft. 1 ft. 12 in 2 conversion factors multiply by the one that will give you the correct units in your answer.

48. Examples 11 yards = 2 rod 40 rods = 1 furlong 8 furlongs = 1 mile The Kentucky Derby race is 1.25 miles. How long is the race in rods, furlongs, meters, and kilometers? A marathon race is 26 miles, 385 yards. What is this distance in rods and kilometers?

49. Examples Science fiction often uses nautical analogies to describe space travel. If the starship U.S.S. Enterprise is traveling at warp factor 1.71, what is its speed in knots? Warp 1.71 = 5.00 times the speed of light speed of light = 3.00 x 10 8 m/s 1 knot = 2000 yd/h exactly

50. Because you never learned dimensional analysis, you have been working at a fast food restaurant for the past 35 years wrapping hamburgers. Each hour you wrap 184 hamburgers. You work 8 hours per day. You work 5 days a week. you get paid every 2 weeks with a salary of $840.34. How many hamburgers will you have to wrap to make your first one million dollars? Examples

51. A senior was applying to college and wondered how many applications she needed to send. Her counselor explained that with the excellent grade she received in chemistry she would probably be accepted to one school out of every three to which she applied. She immediately realized that for each application she would have to write 3 essays, and each essay would require 2 hours work. Of course writing essays is no simple matter. For each hour of serious essay writing, she would need to expend 500 calories which she could derive from her mother's apple pies. Every three times she cleaned her bedroom, her mother would made her an apple pie. How many times would she have to clean her room in order to gain acceptance to 10 colleges?

52. Temperature and Density

53. Temperature A measure of the average kinetic energy Different temperature scales, all are talking about the same height of mercury. We make measurements in lab using the Celsius scale, but most chemistry problems require you to change the temperature to Kelvin before using in an equation.

54. Converting ºF to ºC and vice versa Fahrenheit to Celsius(°F - 32) x 5 / 9 = °C Celsius to Fahrenheit(°C × 9 / 5 ) + 32 = °F

55. 0ºC 32ºF 0ºC = 32ºF

56. 100ºC212ºF 100ºC = 212ºF 0ºC = 32ºF 0ºC 32ºF

57. Converting o C to K and vice versa Celsius to Kelvin K = o C + 273.15 Kelvin to Celsius o C = K - 273.15

58. Density Ratio of mass to volume D = m/V Useful for identifying a compound Useful for predicting weight An intrinsic property- does depend on what the material is.

59. Density Problem An empty container weighs 121.3 g. Filled with carbon tetrachloride (density 1.53 g/cm 3 ) the container weighs 283.2 g. What is the volume of the container?

60. Density Problem A 55.0 gal drum weighs 75.0 lbs. when empty. What will the total mass be when filled with ethanol? density 0.789 g/cm 3 1 gal = 3.78 L 1 lb = 454 g