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Chapter 1 Chemical Foundations

Chapter 1 Chemical Foundations. AP Chemistry. Objectives. Recall units of measure Describe uncertainty in measurement Use scientific notation for numbers Apply significant figure rules. Units of Measure. SI Base Units Mass Kilogram – kg 1 kilogram is about 2.20 pounds Length

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Chapter 1 Chemical Foundations

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  1. Chapter 1Chemical Foundations AP Chemistry

  2. Objectives • Recall units of measure • Describe uncertainty in measurement • Use scientific notation for numbers • Apply significant figure rules

  3. Units of Measure • SI Base Units • Mass • Kilogram – kg • 1 kilogram is about 2.20 pounds • Length • Meter – m • 1 meter is about 3 feet

  4. Units of Measure • Time • Second – s • The time needed for a cesium-133 atom to perform 9,192,631,770 complete oscillations. • Temperature • Kelvin – K • 273 K = 0 degrees C • Amount • Mole – mol • 1 mol = 6.022x1023 particles

  5. Units of Measure • Current • Ampre – A • Luminous Capacity • Candela – cd • First five are the most commonly used in chemistry

  6. Volume • Volume is not an SI Base Unit • Metric system • Powers of 10 • 1 Liter is 1/1000 of a cubic meter • 1 Liter (L) = 1000 cm3 = 1000 mL

  7. Cubic Meter Liter Milliliter

  8. Volume? • Uncertainty is in the last digit.

  9. Accuracy and Precision • Accuracy • The nearness of a measurement to its accepted value • Precision • The agreement between numerical values • You can be precise without being accurate

  10. Accurate & Precise Precise Neither • Loss of accuracy due to systematic errors • Error in same direction every time • Random Error give erratic results • Poor technique

  11. Significant Figures • All known digits plus one estimated digit in a measurement

  12. What is the length? • 2 Sig. Fig • 1 Known Digit • 1 Estimated Digit

  13. Significant Figures • Rule #1 • All Nonzero digits are significant • Ex. 76.44 mL • Ex. 285.85 s

  14. Significant Figures • Rule #2 • “Captive Zeros” • Zeros appearing between nonzero digits are significant • Ex. 308.2001 g =

  15. Significant Figures • Rule #3 • “Leading Zeros” • Zeros appearing in front of nonzero digits are not significant • Ex. 0.007036 g • Takes care of unit changes

  16. Significant Figures • Rule #4 • “Ending Zeros” • Ending zeros are significant if there is a decimal place • Ex. 53.00 m • Ex. 40000. m • Ex. 40000 m • 40000. is much more precise than 40000

  17. What is the length?

  18. Significant Figures • Rule #5 • “Exact Numbers” • Exact number have an unlimited number of significant figures • Exact numbers are counting numbers or definitions • 2 cars or 1000g/1kg

  19. Significant Figures • Rule #6 • “Scientific Notation” • All numbers that come before the x10n are significant • Must be in proper form • Ex. 3.33x105 • Ex. 2.04x10-4

  20. Rounding • 5 and larger round up • 4 and smaller round down • Round the following • 34.567 to 2 SF = • 756.44 to 4 SF = • 0.004325 to 3 SF = • 3436543 to 2 SF =

  21. Addition and Subtraction w/ SF • Addition and Subtraction • The answer must have the same number of digits to the right of the decimal as there are in the measurement having the fewest digits to the right of the decimal point • Ex. 12.11 m + 15 m = • Number of SF’s does not matter!

  22. Multiplication and Division w/ SF • Multiplication and Division • The answer can have no more SF’s than are in the measurement with the fewest total SF’s • Ex. 55 m / 11.34 s =

  23. Scientific Notation • A method of representing very large or very small numbers • M x 10n • M is a number 1 or larger and less than 10 • n is an integer (positive or negative) • All digits in M are significant (If in proper form)

  24. Converting to Sci. Notation • Move decimal so that M is between 1 and 10 • Determine n by counting the number of places the decimal point was moved • Moved to the left, n is positive • Moved to the right, n is negative

  25. Examples • 340,000,000 = • 5.04x105 = • 0.00000300 = • 2.212x10-4 =

  26. Sci. Notation on Calculators • Enter digits in you calculator using the EE key. • For TI 83’s it is the 2nd of the comma • For TI 30’s it is a key • Saves key strokes • Fewer OOR mistakes • 3.4x106 = 3.4E6 • 7.4x10-5 = 7.4E-5

  27. Sci. Notatation Math Operations • Multiply and Divide • Multiply or divide first number • Add exponents (Multiply) • Subtract exponents (Divide) • Addition and Subtraction • Exponents must be the same • Then add or subtract first number • Exponents stay the same

  28. Calculations • 3.0x105 + 4.0x105 = • 4.0x103 – 2.0x102 = • 7.0x105 * 2.0x104 = • 8.0x106 / 4.0x10-3 = • _____4.5x104____ = 6.2x106 * 3.1x10-8

  29. Objectives • Recall metric prefixes • Convert numbers from one unit to another • Describe different temperature scales • Explain density and perform calculations • Classify matter into groups

  30. The Metric System • A system based on powers of ten • Uses SI Units • Allows easy work with both large and small numbers • Prefixes tell us which power of 10 we are using

  31. SI Prefixes (10x larger)Page 9 in your book • Tera • Giga • Mega • Kilo • Hecto • Deca • Base T G M k h da 1012 109 106 103 102 101 100 1000000000000 1000000000 1000000 1000 100 10 1

  32. SI Prefixes (10x smaller) • Base • Deci • Centi • Milli • Micro • Nano • Pico 1 .1 .01 .001 .000001 .000000001 .000000000001 100 10-1 10-2 10-3 10-6 10-9 10-12 d c m μ n p

  33. Conversions • To convert between units set up conversion factors • Ratios of equality

  34. Convert 67 kg to g

  35. Convert 450 cL to dL

  36. Convert 3.4x108 ng to kg

  37. Converting From Metric To English • Find ratios that are true • Page 18 has some equivalents

  38. Convert 763 cm to yd

  39. Convert 1.2 mi/hr to ft/s

  40. Convert 3.8 m2/hr to cm2/s

  41. Temperature • Many different temp. scales • All 0 marks based on different ideas • 0 ºF Coldest saltwater stays a liquid • 0 ºC Normal Freezing Point of water • 0 K Molecular motion stops 1 K = 1 ºC = 1.8 ºF

  42. Temperature Conversion • Temp K = 273 + Temp C • Temp C = Temp K – 273 • 0 ºC = 273 K • If you need any others look up the equ. • TC= (TF – 32)(5/9) • TF = TC(9/5) + 32

  43. Density • Ratio of mass to volume • Density = __Mass__ Volume • Periodic Trend • Units • Solids – g/cm3 • Liquids – g/mL • Gases – g/L

  44. Density Determination • Mass is determined on a balance • Volume is measured in two ways • Regular objects can be measured • All objects can use water displacement

  45. Density • Physical Property • Can be used to identify a substance Lead 11.35 Iron 7.87 Magnesium 1.74 Zinc 7.13 Copper 8.96

  46. Example: A metal cube has sides measuring 3.00 cm. It has a mass of 242.13g. What is the density? What is the metal?

  47. Density • Physical Property • Can be used to identify a substance Lead 11.35 Iron 7.87 Magnesium 1.74 Zinc 7.13 Copper 8.96

  48. Homework • p.33 #'s 33a-f,36a-d,42 47,57,68

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