1 / 55

SCIENTIFIC MEASUREMENT

SCIENTIFIC MEASUREMENT. CHEM IH: CHAPTER 3. Stating a Measurement. In every measurement there is a Number followed by a Unit from a measuring device The number should also be as precise as the measuring device. Ex: Reading a Meterstick.

terrel
Download Presentation

SCIENTIFIC MEASUREMENT

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. SCIENTIFIC MEASUREMENT • CHEM IH: CHAPTER 3

  2. Stating a Measurement In every measurement there is a • Number followed by a • Unit from a measuring device The number should also be as precise as the measuring device.

  3. Ex: Reading a Meterstick . l2. . . . I . . . . I3 . . . .I . . . . I4. . cm First digit (known) = 2 2.?? cm Second digit (known) = 0.7 2.7? cm Third digit (estimated) between 0.05- 0.07 Length reported =2.75 cm or 2.74 cm or 2.76 cm

  4. UNITS OF MEASUREMENT Use SI units — based on the metric system Length Mass Volume Time Temperature Meter, m Kilogram, kg Liter, L Seconds, s Celsius degrees, ˚C kelvins, K

  5. Metric Prefixes

  6. Conversion Factors Fractions in which the numerator and denominator are EQUAL quantities expressed in different units Example: 1 hr. = 60 min Factors: 1 hr. and 60 min 60 min 1 hr.

  7. How many minutes are in 2.5 hours? Conversion factor 2.5 hr x 60 min = 150 min 1 hr cancel By using dimensional analysis / factor-label method, the UNITS ensure that you have the conversion right side up, and the UNITS are calculated as well as the numbers!

  8. Sample Problem • You have $7.25 in your pocket in quarters. How many quarters do you have? 7.25 dollars 4 quarters 1 dollar = 29 quarters X

  9. Learning Check How many seconds are in 1.4 days? Unit plan: days hr min seconds 1.4 days x 24 hr x ?? 1 day

  10. Wait a minute! What is wrong with the following setup? 1.4 day x 1 day x 60 min x 60 sec 24 hr 1 hr 1 min

  11. Significant Figures • The numbers reported in a measurement are limited by the measuring tool • Significant figures in a measurement include the known digits plus one estimated digit

  12. Counting Significant Figures RULE 1. All non-zero digits in a measured number are significant. Only a zero could indicate that rounding occurred. Number of Significant Figures 38.15 cm 4 5.6 ft 2 65.6 lb ___ 122.55 m___

  13. Leading Zeros RULE 2. Leading zeros in decimal numbers are NOT significant. Number of Significant Figures 0.008 mm 1 0.0156 oz 3 0.0042 lb ____ 0.000262 mL ____

  14. Sandwiched Zeros RULE 3. Zeros between nonzero numbers are significant. (They can not be rounded unless they are on an end of a number.) Number of Significant Figures 50.8 mm 3 2001 min 4 0.702 lb ____ 0.00405 m ____

  15. Trailing Zeros RULE 4. Trailing zeros in numbers without decimals are NOT significant. They are only serving as place holders. Number of Significant Figures 25,000 in. 2 200. yr 3 48,600 gal ____ 25,005,000 g ____

  16. Learning Check A. Which answers contain 3 significant figures? 1) 0.4760 2) 0.00476 3) 4760 B. All the zeros are significant in 1) 0.00307 2) 25.300 3) 2.050 x 103 C. 534,675 rounded to 3 significant figures is 1) 535 2) 535,000 3) 5.35 x 105

  17. Learning Check In which set(s) do both numbers contain the samenumber of significant figures? 1) 22.0 and 22.00 2) 400.0 and 40 3) 0.000015 and 150,000

  18. Significant Numbers in Calculations • A calculated answer cannot be more precise than the measuring tool. • A calculated answer must match the least precise measurement. • Significant figures are needed for final answers from 1) adding or subtracting 2) multiplying or dividing

  19. Adding and Subtracting The answer has the same number of decimal places as the measurement with the fewest decimal places. 25.2one decimal place + 1.34two decimal places 26.54 answer 26.5one decimal place

  20. Learning Check In each calculation, round the answer to the correct number of significant figures. A. 235.05 + 19.6 + 2.1 = 1) 256.75 2) 256.8 3) 257 B. 58.925 - 18.2 = 1) 40.725 2) 40.73 3) 40.7

  21. Multiplying and Dividing Round (or add zeros) to the calculated answer until you have the same number of significant figures as the measurement with the fewest significant figures.

  22. Learning Check A. 2.19 X 4.2 = 1) 9 2) 9.2 3) 9.198 B. 4.311 ÷ 0.07 = 1)61.582) 62 3) 60 C. 2.54 X 0.0028 = 0.0105 X 0.060 1) 11.3 2) 11 3) 0.041

  23. CHEMICAL QUANTITIES: THE MOLEchemIh: Chapter 10chemi: chapter 12

  24. MEASURING MASS • A mole is a quantity of things, just as… 1 dozen = 12 things 1 gross = 144 things 1 mole = 6.02 x 1023 things • “Things” usually measured in moles are atoms, molecules, ions, and formula units

  25. You can measure mass, or volume, or you can count pieces • We measure mass in grams • We measure volume in liters • We count pieces in MOLES

  26. A MOLE… • is an amount, defined as the number of carbon atoms in exactly 12 grams of carbon-12 • 1 mole = 6.02 x 1023of the representative particles • Treat it like a very large dozen 6.02 x 1023is called: Avogadro’s number

  27. Similar Words for an amount: • Pair: 1 pair of shoelaces = 2 shoelaces • Dozen: 1 dozen oranges = 12 oranges • Gross: 1 gross of pencils= 144 pencils • Ream: 1 ream of paper= 500 sheets of paper

  28. What are Representative Particles? • The smallest pieces of a substance: • For a molecular compound: it is the molecule. • For an ionic compound: it is the formula unit (made of ions) • For an element: it is the atom • Remember the 7 diatomic elements? (made of molecules)

  29. How many oxygen atoms in the following? • CaCO3 3 atoms of oxygen • Al2(SO4)3 12 (4 x 3) atoms of oxygen • How many ions in the following? • CaCl2 • 3 total ions (1 Ca2+ ion and 2 Cl1- ions) • NaOH • 2 total ions (1 Na1+ ion and 1 OH1- ion) • Al2(SO4)3 • 5 total ions (2 Al3+ + 3 SO4 ions)

  30. CONVERSION FACTOR • MOLES = representative particles x ____________1 mole_____________ 6.02 x 1023 representative particles

  31. EXAMPLES: ATOMS  MOLES • How many atoms of Al are in 1.5 mol of Al? • Conversion: 1 mole = 6.02 x 1023 atoms 1.5 mol of Al 6.02 x 1023 atoms 1 mole 9.03 x 1023 atoms of Al =

  32. EXAMPLES: MOLECULES  MOLES How many atoms of H are there in 3 moles of H2O? Conversions: 1 mole = 6.02 x 1023 molecules H2O molecule = 2 atoms of Hydrogen 6.02 x 1023 molec 1 mole 2 atoms H 1 H2O molecule 3 moles of H2O = 3.612 x 1024 atoms H

  33. 20 Ca 40.08 MOLAR MASS Determined simply by looking at the periodic chart Molar mass = Atomic Mass * Thus, 1 mol Ca = 40 g Atomic Mass same as Molar Mass

  34. MOLAR MASS • To calculate the molar mass of a compound, find the number of grams in each element in one mole of the compound • Then add the masses within the compound Example: H2O H= 1.012 (1.01) + 1 (15.999)= 18.02 O= 15.999

  35. SOME PRACTICE PROBLEMS • How many atoms of O are in 3.7 mol of O? • 2.2 X 1024 atoms of oxygen • How many atoms of P are in 2.3 mol of P? • 1.4 x 1024 atoms of phosphorus • How many atoms of Ca are there in 2.5 moles of CaCl2? • 1.5 x 1024 atoms Ca • How many atoms of O are there in 1.7 moles of SO4? • 4.1 x 1024 atoms of oxygen

  36. Remember!!!! • The molar mass of any substance (in grams) equals 1 mole • This applies to ALL substance: elements, molecular compounds, ionic compounds • Use molar mass to convert between mass and moles • Ex: Mass, in grams, of 6 mol of MgCl2 ? mass of MgCl2 = 6 mol MgCl2 92.21 g MgCl2 1 mol MgCl2 = 571.26 g MgCl2

  37. VOLUME AND THE MOLE • Volume of 1 mol of solid and liquids are not the same • But gases are more predictable, under the same physical conditions • Avogadro’s hypothesis helps explain: equal volume of gases, at the same temp and pressure contains equal number of particles • Volume varies with changes in temperature • Ex: helium balloon

  38. Gases vary at different temperatures, makes it hard to measure • Because of variation use STP • Standard Temperature and Pressure • Temperature = 0° C • Pressure = 1 atm (atmosphere) or 101.3 kPal

  39. Standard Temperature Pressure • At STP: 1 mole, 6.02 x 1023 atoms, of any gas has a volume of 22.4 L • Called Molar Volume • Used to convert between # of moles and vol of a gas @ STP • Ex: what is the vol of 1.25 mol of sulfur Vol of S= 1.25 mol S 22.4 L = 28 L 1 mol

  40. MOLAR MASS FROM DENSITY • Different gases have different densities • Density of a gas measured in g/L @ a specific temperature • Can use the following formula to solve : grams = grams X 22.4 L mole L 1 mole • Ex: Density of gaseous compound containing oxygen and carbon is 1.964 g/ L, what is the molar mass? • grams = 1.964 g X 22.4 L then you solve mole 1 L 1 mole = 44.o g/mol

  41. Calculating Percent Composition of a Compound • Like all percent problems: a part ÷ the whole • Find the mass of each of the components (the elements) • Next, divide by the total mass of the compound • Then X 100 % = percent Formula: % Composition = Mass of element X 100% Mass of compound

  42. Example: A compound is formed when 9.03 g of Mg combines completely with 3.48 g of N. What is the percent composition of the compound? • First add the 2 mass of the 2 compounds to reach the total mass 9.03 g Mg + 3.48 g N = 12.51 g Mg3N2 • Find the % of each compound % Mg= 9.03 g Mg X 100% = 72.2 % 12.51 g Mg3N2 % N= 3.48 g N X 100% = 27.8 % 12.51 g Mg3N2

  43. % Composition from Chemical Formula • Can find the percent composition of a compound using just the molar mass of the compound and the element • % mass=mass of the element 1 mol cmpd X100% molar mass of the compound • Example: Find the percent of C in CO2 12.01 g C X 100% = 27.3% C 44.01 g CO2 Can find O % by subtracting 27.3% from 100%

  44. Using % Composition • Can use % composition as a conversion factor just like the mole • After finding the % comp. of each element in a cmpd. can assume the total compound = 100g • Example: C= 27.3% 27.3 g C O= 72.7 % 72.7 g O • In 100 g sample of compound there is 27.3 g of C & 72.7 g of O How much C would be contained in 73 g of CO2? 73 g CO2 27.3 g C = 19.93 g C 100 g CO2

  45. EMPIRICAL FORMULAS • Empirical formulas are the lowest WHOLE number ratios of elements contained in a compound

  46. REMEMBER… • Molecular formulas tells the actual number of of each kind of atom present in a molecule of the compound • Ex: H2O2 HO Molecular Empirical Formula Formula CO2 CO2 Molecular Empirical For CO2 they are the same Formula Formula

  47. Formulas for ionic compounds are ALWAYS empirical (the lowest whole number ratio = can not be reduced) • Examples: NaCl MgCl2 Al2(SO4)3 K2CO3 Simplest whole number ratio for NaCl

  48. A formula is not just the ratio of atoms, it is also the ratio of moles • In 1 mole of CO2 there is 1 mole of carbon and 2 moles of oxygen • In one molecule of CO2 there is 1 atom of C and 2 atoms of O • Formulas for molecular compounds MIGHT be empirical (lowest whole number ratio) • Molecular: H2O C6H12O6 C12H22O11 (Correct formula) • Empirical: (Lowest whole H2O CH2O C12H22O11 number ratio)

  49. CALCULATING EMPIRICAL • We can get a ratio from the percent composition • Assume you have a 100 g sample the percentage become grams (75.1% = 75.1 grams) • Convert grams to moles • Find lowest whole number ratio by dividing each number of moles by the smallest value

More Related