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### Ch 11 notes complete

1. 1. CH 11: The MoleSection 1 Counting Matter
2. 2. Counting MatterWhat are some everyday wayswe count matter?1. DOZEN = 12 things2. 1 GROSS = 144 things What about molecules? Or atoms?3. BUSHEL of corn = 21.772 kg 1 MOLE = ???
3. 3. The Mole  It represents a counted number of things.  IN Chemistry the term MOLE represents the number of particles in a substance.In Chemistry is NOT this furry little animalor the spot on your face…
4. 4. Just how many is a mole? One mole represents 6.02 x 1023 of things (units, molecules, compounds, formula units). This is called Avogadro’s number. One mole of most elements contains 6.02 x 1023 atoms. 1 mole O2 = 6.02x1023 molecules of O2 602 000 000 000 000 000 000 000
5. 5. Just how big is a mole? Listen to The Mole Song! An Avogadros number of standard soft drink cans would cover the surface of the earth to a depth of over 200 miles If you had Avogadros number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.
6. 6. Solving the Problems SamplesRequired: dimensional analysis/factor label How many molecules are in 3.00 moles of N2? 6.02 x 1023 moleculesN 23 molesN 2 1.81 1024 molecules 1 mole N 2 How many moles of Na are in 1.10 x 1023 atoms? 23 1 mole Na1.10 10 atoms 0.182 moles Na 6.02 x 10 23 atoms Na
7. 7. Practice moles to particles Determine the number of atoms in 2.50 mol Zn. Given 3.25 mol AgNO3, determine the number of formula units. Calculate the number of molecules in 11.5 mol H2O.
8. 8. Practice Particles to Moles How many moles contain each of the following?  5.75 x 1024 atoms Al  3.75 x 1024 molecules CO2  3.58 x 1023 formula units ZnCl2  2.50 x 1020 atoms Fe
9. 9. Mixed Avogadro’s number Practice1. How many moles are in 2x1060 atoms of Cu?2. If there are 1.25x1045 molecules of dihydrogen monoxide in a glass, how many moles are there?3. 5 mol of Iron III sulfate is equal to how many compounds of Fe2(SO4)34. For every 3.5 mol of S there are how many atoms of sulfur?5. How many mol of O atoms are in 1.50 mol SO2?6. How many mol of H are in 1 mol (NH4) 2CO3?7. How many atoms of H are in 1 mol (NH4) 2CO3?
10. 10. 11-2 Mass & the Mole11-3 Moles of Compounds
11. 11. Molar Mass Defn: is the mass (think grams) of one mole of a substance  Atomic masses (from periodic table) represent molar mass.  Units g/mol  1 mole of Carbon has 6.02 x 1023 atoms of C and they have a mass of 12.01 grams.  To calculate the molar mass of a compound, you add up the molar masses of all the elements in that compound
12. 12. Molar Mass Practice What is the mass of 1.00 mole of Oxygen? Of Nitrogen? 1 mole O = 16.0 grams 1 mole N = 14.0 grams Find the molar mass for:  SO3 SO3= 80 g/mole  Na2SO4 1 Mole = 142.043g Tutorial Site Molar Mass Calculator for homework help
13. 13. Molar Mass Practice  When you see 1.00 mole = _?_ g, think “g means GO to the PERIODIC TABLE” to find the molar mass. http://www.webelements.com/
14. 14. Practice Problems Determine the molar mass of each of the following ionic compounds:  NaOH  CaCl2  KC2H3O2  HCN  CCl4  H2O
15. 15. Grams-Mole Conversions How many moles are in 56.8 g of HCl? 1 moleHCl 56.8 g HCl 1.58 molesHCl 36 g HCl How many grams are in .05 moles Na2SO4? 90 g Na 2SO 4.05 mole Na 2SO 4 4.5 g Na 2SO 4 1 mole Na 2SO 4
16. 16. Practice Mole to Grams Determine the mass in grams of each of the following.  3.57 mol Al  42.6 mol Si  3.45 mol Co  2.45 mol Zn
17. 17. Practice Gram to Mole How many atoms are in each of the following samples?  55.2 g Li  0.230 g Pb  11.5 g Hg  45.6 g Si  0.120 kg Ti
18. 18. Ex: Mass to Particle Conversion Gold is one of a group of metals called the coinage metals (copper, silver and gold). How many atoms of gold (Au) are in a pure gold nugget having a mass of 25.0 g?Known: Unknown:Mass = 25.0 g Au Number of atoms = ? Atoms AuMolar mass Au = 196.97 g/mol Au25.0 g Au x 1 mole Au x 6.02 x 1023 atoms Au = 7.65 x 1022 atoms Au 196.97 g Au 1 mol Au
19. 19. Practice Problems How many atoms are in each of the following samples?  55.2 g Li  0.230 g Pb  11.5 g Hg  45.6 g Si  0.120 kg Ti
20. 20. Ex: Particle to Mass Conversion A party balloon contains 5.50 x 1022 atoms of helium (He) gas. What is the mass in grams of the helium?Known: Unknown:Number of atoms = 5.50 x 1022 atoms He Mass = ? G HeMolar mass He = 4.00 g/mol He5.50 x 1022 atoms He x 1 mol He x 4.00 g He = 0.366 g He 6.02 x 1023 atoms He 1 mol He
21. 21. Practice Problems What is the mass in grams of each of the following?  6.02 x 1024 atoms Bi  1.00 x 1024 atoms Mn  3.40 x 1022 atoms He  1.50 x 1015 atoms N  1.50 x 1015 atoms U
22. 22. Molar Volume The volume of a gas is usually measured at standard temperature and pressure (_STP_) Standard temp = ___0°_ C Standard pressure = ___1___ atmosphere (atm) 1 mole of any gas occupies __22.4__ L of space at STP
23. 23. Molar Volume Practice How many moles would 45.0 L of He gas be? How many liters of O2 would 3.8 moles occupy?
24. 24. Putting it all together 1.0 mole = _6.02 x 1023___atoms or molecules 1.0 mole = _?__ g(PT) 1.0 mole = _22.4 L (at STP)
25. 25. Helpful Chart! MOLES Volume Grams in Liters Atoms or Molecules
26. 26. Chemical Formulas and the Mole Chemical formula for a compound indicates the types of atoms and the number of each contained in one unit.  Ex. CCl2F2 - Freon  Ratio of carbon to chlorine to fluorine is 1:2:2  Ratios can be written: or  In one mole of freon you would have 1 mole of carbon, 2 moles of chlorine and 2 moles of fluorine.
27. 27. Ex: Mole Relationship from Chemical Formulas Determine the moles of aluminum ions (Al3+) in 1.25 moles of aluminum oxide1.25 molAl2O3 x 2 mol Al3+ ions = 2.50 mol Al3+ ions 1 mole Al2O3
28. 28. Practice Problems  Determine the number of moles of chloride ions in 2.53 mol ZnCl2.  Calculate the number of moles of each element in 1.25 mol glucose (C6H12O6).  Determine the number of moles of sulfate ions present in 3.00 mol iron (III) sulfate (Fe2(SO4)3).  How many moles of oxygen atoms are present in 5.00 mol diphosphorus pentoxide?  Calculate the number of moles of hydrogen atoms in 11.5 mol water.
29. 29. Practice Problems Mixed How many chloride ions are in 2.50 mol of ZnCl2 Calculate the number of atoms of each element in 1.25 mol glucose (C6H12O6). Determine the number of moles of sulfate ions present in 6.02x1023 compounds of iron (III) sulfate (Fe2(SO4)3). How many atoms of oxygen are present in 3.01x1023 molecules of diphosphorus pentoxide? Calculate the number of molecules in 11.5 mol water.
30. 30. Practice Problems A sample of silver chromate has a mass of 25.8 g.  How many Ag+ ions are present?  How many CrO42- ions are present?  What is the mass in grams of one unit of silver chromate? What mass of sodium chloride contains 4.59 x 1024 units? A sample of ethanol (C2H5OH) has a mass of 45.6 g.  How many carbon atoms does the sample contain?  How many hydrogen atoms are present?  How many oxygen atoms are present?
31. 31. 11-4 Empirical &Molecular Formulas
32. 32. Percent Composition the percentage by mass of each element in a compound The percent comp. is found by using the following formula: mass of 1 element% Mass 100 molarmass of compound
33. 33. Percent Composition ExampleEx. Compound XY is 55g element X and 45g element Y 55 g of element X x 100 = 55 % element X 100 g of compound 45 g of element Y x 100 = 45 % element Y 100 g of compound
34. 34. Percent Composition fromChemical Formula First find the molar mass of each element and the molar mass of the compound Ex: what is the % composition of H in 1 mole of H2O?  Multiplythe molar mass of the element by its subscript in the formula.  1.01 g/mol H x 2 mol = 2.02 g H % by mass H = 2.02 g x 100 = 11.2% H 18.02 g H2O
35. 35. Percent Composition fromChemical FormulaExample continued – Molar mass of O for each mole of H2O? 16.00 g/mol O x 1mol O = 16.00 g O 16.00 g x 100 = 88.8 % O 18.02 g
36. 36. % Composition Practice What is the percent of C & H in C2H6? What is the percent of each element in Na2SO3?
37. 37. Empirical Formulas This is the LOWEST whole number ratio of the elements in a compound. For example, the empirical formula for  Molecular Formula C6H6  Empirical Formula CH What is the empirical formula for each?  C2H6  C6H12O6
38. 38. 11-3: Calculating EmpiricalFormula Steps for caluculating Empirical Formula give mass or percent composition: 1. If given a percent sign (%), remove the sign & change to GRAMS.  You are assuming you have 100 g of the compound. 2. Convert grams ---> moles. 3. Select lowest number of moles 4. Divide each number of moles by this number. 5. If the number divides out evenly, these are the subscripts of the elements in the compound. 6. If any of the numbers have a .5, MULTIPLY them ALL by TWO & then place these numbers as the subscripts.
39. 39. Example Using PercentComposition The percent composition of an oxide of sulfur is 40.05% S and 59.95% O. Assuming you have a 100g sample, it contains 40.05g S and 59.95g O. Convert to moles using molar mass: 40.05g S x 1 mol = 1.249 mol S 32.07g 59.95g O x 1 mol = 3.747 mol O 16.00g
40. 40. Example Using PercentCompositionThe mole ratio of S atoms to O atoms in the oxide is 1.249 : 3.747.Recognize that S has the smallest possible number of moles at ~1. Make the mole value of S equal to 1 by dividing both mole values by 1.249. 1.249 mol S = 1 mol S 1.249 3.747 mol O = 3 mol O 1.249The simplest whole number mole ratio of S atoms to O atoms is 1 : 3. The empirical formula for the oxide of sulfur is SO3.
41. 41. Sample Problems What is the empirical formula for a compound which is 75 % C and 25 % H? What is the empirical formula for a compound which has  48.64 % C,  8.16 % H  43.20 % O
42. 42. Sample Problems #2 What is the empirical formula of  40.68 %C  5.08 % H  54.24 % O
43. 43. Practice Problems A blue solid is found to contain 36.894% N and 63.16% O. What is the empirical formula for this solid? Determine the empirical formula for a compound that contains 35.98% Al and 64.02% S. Propane is a hydrocarbon, a compound composed only of carbon and hydrogen. It is 81.82% C and 18.18% H. What is the empirical formula? The chemical analysis of aspirin indicates that the molecule is 60.00% C, 4.44% H and 35.56% O. Determine the empirical formula. What is the empirical formula for a compound that contains 10.89% Mg, 31.77% Cl, and 57.34% O?
44. 44. Molecular Formula Specifies the actual number of atoms of each element in one molecule or formula unit of the substance  n=ratio between experimentally determined mass of compound and the molar mass of the empirical formula.
45. 45. Calculating Molecular Formula Molar mass of acetylene – 26.04 g/mol Mass of empirical formula (CH) – 13.02 g/mol n – Obtained by dividing the molar mass by the mass of the empirical formula indicates that the molar mass of acetylene is two times the mass of the empirical formula.Experimentally determined molar mass of acetylene = 26.04 g/mol = 2.000 mass of empirical formula CH 13.02 g/mol Molecular Formula = (CH)2 Acetylene = C2H2
46. 46. Determining a MolecularFormula Succinic acid is a substance produced by lichens. Chemical analysis indicates it is composed of 40.68% C, 5.08% H, and 54.24% O and has a molar mass of 118.1 g/mol. Determine the empirical and molecular formulas for succinic acid.Known: Unknown:Percent by mass = 40.68% C empirical formula = ?Percent by mass = 5.08% H molecular formula = ?Percent by mass = 54.24% O
47. 47. Practice Problems Analysis of a chemical used in photographic developing fluid indicates a chemical composition of 65.45% C, 5.45% H, and 29.09% O. The molar mass is found to be 110.0 g/mol. Determine the molecular formula. A compound was found to contain 49.98 g C, and 10.47 g H. The molar mass of the compound is 58.12 g/mol. Determine the molecular formula. A colorless liquid composed of 46.68% N and 53.32% O has a molar mass of 60.01 g/mol. What is the molecular formula?
48. 48. 11-4 The Formulafor a Hydrate
49. 49. Naming HydratesHydrate: a compound that has a specific number of water molecules bound to its atoms.In the formula for a hydrate, the number of water molecules associated with each formula unit of the compound is written following a dot. ex. Na2CO3·10H2O called sodium carbonate decahydrate deca- means 10 and hydrate means water Therefore there are 10 water molecules are associated with one formula unit of the compound.
50. 50. Formulas for Hydrates and ExamplesPrefix Molecules H2O Formula NameMono- 1 (NH4)2C2O4·H2O Ammonium oxalate monohydrate Di- 2 CaCl2·2H2O Calcium chloride dihydrate Tri- 3 NaC2H3O2·3H2O Sodium acetate trihydrateTetra- 4 FePO4·4H2O Iron(III) phosphate tetrahydratePenta- 5 CuSO4·5H2O Copper(II) sulfate pentahydrateHexa- 6 CoCl2·6H2O Cobalt(II) chloride hexahydrateHepta- 7 MgSO4·7H2O Magnesium sulfate heptahydrateOcta- 8 Ba(OH)2·8H2O Barium hydroxide octahydrateDeca- 10 Na2CO3·10H2O Sodium carbonate decahydrate
51. 51. Analyzing a Hydrate Must drive off the water by heating the compound Substance remaining after heating is anhydrous (without water)  Example: hydrated cobalt(II) chloride is a pink solid that turns a deep blue when the water of hydration is driven off and anhydrous cobalt(II) chloride is produced
52. 52. Formula for a Hydrate To determine the formula for a hydrate, you must determine the number of moles of water associated with one mole of the hydrate.
53. 53. Example Problem – Determining theFormula for a Hydrate A mass of 2.5 g of blue, hydrated copper sulfate (CuSO4·xH2O) is placed in a crucible and heated. After heating, 1.59 g white anhydrous copper sulfate (CuSO4) remains. What is the formula for the hydrate? Name the hydrate.
54. 54. Example Problem – (cont.)Known:Mass of hydrated compound = 2.50 g CuSO4·xH2OMass of anhydrous compound = 1.59 g CuSO4Molar mass = 18.02 g/mol H2OMolar mass = 159.6 g/mol CuSO4Unknown:Formula for hydrate = ?Name of hydrate = ?
55. 55. Example Problem – (cont.)Subtract the mass of the anhydrous copper sulfate from the mass of the hydrated copper sulfate to determine the mass of water lost: mass of hydrates copper sulfate 2.50 g mass of anhydrous copper sulfate - 1.59 g mass of water lost 0.91 gCalculate the number of moles of H2O and anhydrous CuSO4 1.59 g CuSO4 x 1 mol CuSO4 = 0.00996 mol CuSO4 159.6 g CuSO4 0.91 g H2O x 1 mol H2O = 0.050 mol H2O 18.02 g H2O
56. 56. Example Problem – (cont.)Determine the value of x. x = moles H2O = 0.050 mol H2O = 5.0 mol H2O = 5 moles CuSO4 0.00996 mol CuSO4 1.0 mol CuSO4 1The ratio of H2O to CuSO4 is 5 : 1, so the formula for the hydrate is CuSO4·5H2O, copper(II) sulfate pentahydrate.
57. 57. Practice Problems A hydrate is found to have the following percent composition: 48.18% MgSO4 and 51.2% H2O. What is the formula and name for this hydrate? If 11.75 g of the common hydrate cobalt(II) chloride is heated, 9.25 g of anhydrous cobalt chloride remains. What is the formula and name for this hydrate?
58. 58. Uses of Hydrates Drying agents: CaCl2, CaSO4 Storage of solar energy: Na2SO4·10H2O