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Ch 11 notes complete

  1. 1. CH 11: The Mole<br />Section 1 Counting Matter<br />
  2. 2. Counting Matter<br />What are some everyday ways we count matter?<br />1.<br />2.<br />3.<br />DOZEN = 12 things<br />1 GROSS = 144 things<br />What about molecules? Or atoms?<br />BUSHEL of corn = 21.772 kg <br />1 MOLE = ???<br />
  3. 3. The Mole<br />It represents a counted number of things. <br />IN Chemistry the term MOLE represents the number of particles in a substance.<br />In Chemistry is NOT this furry little animal or the spot on your face…<br />
  4. 4. Just how many is a mole?<br />One mole represents 6.02 x 1023 of things (units, molecules, compounds, formula units). This is called Avogadro’s number.<br />One mole of most elements contains 6.02 x 1023atoms.<br />1 mole O2 = 6.02x1023molecules of O2<br />602 000 000 000 000 000 000 000<br />
  5. 5. Just how big is a mole?<br />Listen to The Mole Song! <br />An Avogadro's number of standard soft drink cans would cover the surface of the earth to a depth of over 200 miles<br />If you had Avogadro's 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.<br />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. <br />
  6. 6. Solving the Problems Samples Required: dimensional analysis/factor label<br />How many molecules are in 3.00 moles of N2?<br />How many moles of Na are in 1.10 x 1023 atoms?<br />
  7. 7. Practice moles to particles<br />Determine the number of atoms in 2.50 mol Zn.<br />Given 3.25 mol AgNO3, determine the number of formula units.<br />Calculate the number of molecules in 11.5 mol H2O.<br />
  8. 8. Practice Particles to Moles<br />How many moles contain each of the following?<br />5.75 x 1024 atoms Al<br />3.75 x 1024 molecules CO2<br />3.58 x 1023 formula units ZnCl2<br />2.50 x 1020 atoms Fe<br />
  9. 9. 11-2 Mass & the Mole 11-3 Moles of Compounds <br />
  10. 10. Molar Mass<br />Defn: is the mass (think grams) of one mole of a substance<br />Atomic masses (from periodic table) represent molar mass.<br />Units g/mol<br />1 moleof Carbon has 6.02 x 1023 atoms of C and they have a mass of12.01grams.<br />To calculate the molar mass of a compound, you add up the molar masses of all the elements in that compound<br />
  11. 11. Molar Mass Practice<br />What is the mass of 1.00 mole of Oxygen? Of Nitrogen?<br />Find the molar mass for:<br />SO3<br />Na2SO4<br />1 mole O = 16.0 grams<br />1 mole N = 14.0 grams<br />SO3= 80 g/mole<br />1 Mole = 142.043g<br />Tutorial Site<br />Molar Mass Calculator for homework help<br />
  12. 12. Molar Mass Practice<br />When you see 1.00 mole = _?_ g, think “g means GO to the PERIODIC TABLE” to find the molar mass.<br /><br />
  13. 13. Practice Problems<br />Determine the molar mass of each of the following ionic compounds: <br />NaOH<br />CaCl2<br />KC2H3O2<br />HCN<br />CCl4<br />H2O<br />
  14. 14. Grams-Mole Conversions<br />How many moles are in 56.8 g of HCl?<br />How many grams are in .05 moles Na2SO4?<br />
  15. 15. Practice Mole to Grams<br />Determine the mass in grams of each of the following.<br />3.57 mol Al<br />42.6 mol Si<br />3.45 mol Co<br />2.45 mol Zn<br />
  16. 16. Practice Gram to Mole<br />How many atoms are in each of the following samples?<br />55.2 g Li<br />0.230 g Pb<br />11.5 g Hg<br />45.6 g Si<br />0.120 kg Ti<br />
  17. 17. Ex: Mass to Particle Conversion<br />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?<br />Known:Unknown:<br />Mass = 25.0 g Au <br />Molar mass Au = 196.97 g/mol Au<br />25.0 g Au x 1 mole Au x 6.02 x 1023 atoms Au = 7.65 x 1022 atoms Au<br /> 196.97 g Au 1 mol Au<br />Number of atoms = ? Atoms Au<br />
  18. 18. Practice Problems<br />How many atoms are in each of the following samples?<br />55.2 g Li<br />0.230 g Pb<br />11.5 g Hg<br />45.6 g Si<br />0.120 kg Ti<br />
  19. 19. Ex: Particle to Mass Conversion<br />A party balloon contains 5.50 x 1022 atoms of helium (He) gas. What is the mass in grams of the helium?<br />Known:Unknown:<br />Number of atoms = 5.50 x 1022 atoms He <br />Molar mass He = 4.00 g/mol He<br />5.50 x 1022 atoms He x 1 mol He x 4.00 g He = 0.366 g He<br /> 6.02 x 1023 atoms He 1 mol He<br />Mass = ? G He<br />
  20. 20. Practice Problems<br />What is the mass in grams of each of the following?<br />6.02 x 1024 atoms Bi<br />1.00 x 1024 atoms Mn<br />3.40 x 1022 atoms He<br />1.50 x 1015 atoms N<br />1.50 x 1015 atoms U<br />
  21. 21. Molar Volume<br />The volume of a gas is usually measured at standard temperature and pressure (_STP_)<br />Standard temp = ___0°_ C<br />Standard pressure = ___1___ atmosphere (atm)<br />1 mole of any gas occupies __22.4__ L of space at STP<br />
  22. 22. Molar Volume Practice<br />How many moles would 45.0 L of He gas be?<br />How many liters of O2 would 3.8 moles occupy?<br />
  23. 23. Putting it all together <br />1.0 mole = _6.02 x 1023___atoms or molecules 1.0 mole = _?__ g(PT)<br /> 1.0 mole= _22.4 L (at STP) <br />
  24. 24. Helpful Chart!<br />MOLES<br />Volume <br />in Liters<br />Grams<br />Atoms <br />or Molecules<br />
  25. 25. Chemical Formulas and the Mole<br />Chemical formula for a compound indicates the types of atoms and the number of each contained in one unit.<br />Ex. CCl2F2 - Freon<br />Ratio of carbon to chlorine to fluorine is 1:2:2<br />Ratios can be written: or <br />In one mole of freon you would have 1 mole of carbon, 2 moles of chlorine and 2 moles of fluorine.<br />
  26. 26. Ex: Mole Relationship from Chemical Formulas<br />Determine the moles of aluminum ions (Al3+) in 1.25 moles of aluminum oxide<br />1.25 molAl2O3 x 2 mol Al3+ ions = 2.50 mol Al3+ ions<br /> 1 mole Al2O3<br />
  27. 27. Practice Problems<br />Determine the number of moles of chloride ions in 2.53 mol ZnCl2.<br />Calculate the number of moles of each element in 1.25 mol glucose (C6H12O6).<br />Determine the number of moles of sulfate ions present in 3.00 mol iron (III) sulfate (Fe(SO4)3).<br />How many moles of oxygen atoms are present in 5.00 mol diphosphoruspentoxide?<br />Calculate the number of moles of hydrogen atoms in 11.5 mol water.<br />
  28. 28. Practice Problems<br />A sample of silver chromate has a mass of 25.8 g.<br />How many Ag+ ions are present?<br />How many CrO42- ions are present?<br />What is the mass in grams of one unit of silver chromate?<br />What mass of sodium chloride contains 4.59 x 1024 units?<br />A sample of ethanol (C2H5OH) has a mass of 45.6 g.<br />How many carbon atoms does the sample contain?<br />How many hydrogen atoms are present?<br />How many oxygen atoms are present?<br />
  29. 29. 11-4 Empirical & Molecular Formulas<br />
  30. 30. Percent Composition<br />the percentage by mass of each element in a compound<br />The percent comp. is found by using the following formula:<br />
  31. 31. Percent Composition Example<br />Ex. Compound XY is 55g element X and 45g element Y<br />55 g of element Xx 100 = 55 % element X<br /> 100 g of compound<br />45 g of element Y x 100 = 45 % element Y<br /> 100 g of compound<br />
  32. 32. Percent Composition from Chemical Formula<br />First find the molar mass of each element and the molar mass of the compound<br />Ex: what is the % composition of H in 1 mole of H2O?<br />Multiply the molar mass of the element by its subscript in the formula.<br />1.01 g/mol H x 2 mol = 2.02 g H<br />% by mass H = 2.02 g x 100 = 11.2% H<br /> 18.02 g H2O<br />
  33. 33. Percent Composition from Chemical Formula<br />Example continued – <br />Molar mass of O for each mole of H2O? <br /> 16.00 g/mol O x 1mol O = 16.00 g O<br />16.00 g x 100 = 88.8 % O<br /> 18.02 g<br />
  34. 34. % Composition Practice<br />What is the percent of C & H in C2H6?<br />What is the percent of each element in Na2SO3?<br />
  35. 35. Empirical Formulas<br />This is the LOWEST whole number ratio of the elements in a compound. For example, the empirical formula for<br />Molecular Formula C6H6<br />Empirical Formula CH<br />What is the empirical formula for each? <br />C2H6 <br />C6H12O6<br />
  36. 36. 11-3: Calculating Empirical Formula<br />Steps for caluculating Empirical Formula give mass or percent composition:<br />If given a percent sign (%), remove the sign & change to GRAMS.<br /><ul><li>You are assuming you have 100 g of the compound.</li></ul>Convert grams ---> moles.<br />Select lowest number of moles<br />Divide each number of moles by this number.<br />If the number divides out evenly, these are the subscripts of the elements in the compound.<br />If any of the numbers have a .5, MULTIPLY them ALLby TWO& then place these numbers as the subscripts.<br />
  37. 37. Example Using Percent Composition<br />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.<br />Convert to moles using molar mass:<br />40.05g S x 1 mol = 1.249 mol S<br /> 32.07g<br />59.95g O x 1 mol = 3.747 mol O<br /> 16.00g<br />
  38. 38. Example Using Percent Composition<br />The mole ratio of S atoms to O atoms in the oxide is 1.249 : 3.747.<br />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.<br />1.249 mol S = 1 mol S<br /> 1.249<br />3.747 mol O = 3 mol O<br /> 1.249<br />The simplest whole number mole ratio of S atoms to O atoms is 1 : 3. The empirical formula for the oxide of sulfur is SO3.<br />
  39. 39. Sample Problems<br />What is the empirical formula for a compound which is 75 % C and 25 % H?<br />What is the empirical formula for a compound which has <br />48.64 % C,<br /> 8.16 % H<br />43.20 % O<br />
  40. 40. Sample Problems #2<br />What is the empirical formula of <br />40.68 % C<br />5.08 % H<br />54.24 % O<br />
  41. 41. Practice Problems<br />A blue solid is found to contain 36.894% N and 63.16% O. What is the empirical formula for this solid?<br />Determine the empirical formula for a compound that contains 35.98% Al and 64.02% S.<br />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?<br />The chemical analysis of aspirin indicates that the molecule is 60.00% C, 4.44% H and 35.56% O. Determine the empirical formula.<br />What is the empirical formula for a compound that contains 10.89% Mg, 31.77% Cl, and 57.34% O?<br />
  42. 42. Molecular Formula<br />Specifies the actual number of atoms of each element in one molecule or formula unit of the substance<br />n= ratio between experimentally determined mass of compound and the molar mass of the empirical formula.<br />
  43. 43. Calculating Molecular Formula<br /> Molar mass of acetylene – 26.04 g/mol<br /> Mass of empirical formula (CH) – 13.02 g/mol<br />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.<br />Experimentally determined molar mass of acetylene = 26.04 g/mol = 2.000<br /> mass of empirical formula CH 13.02 g/mol<br />Molecular Formula = (CH)2<br />Acetylene = C2H2<br />
  44. 44. Determining a Molecular Formula<br />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.<br />Known: Unknown:<br />Percent by mass = 40.68% C empirical formula = ?<br />Percent by mass = 5.08% H molecular formula = ?<br />Percent by mass = 54.24% O<br />
  45. 45. Practice Problems<br />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.<br />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.<br />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?<br />
  46. 46. 11-4 The Formula for a Hydrate<br />
  47. 47. Naming Hydrates<br />Hydrate: a compound that has a specific number of water molecules bound to its atoms.<br />In the formula for a hydrate, the number of water molecules associated with each formula unit of the compound is written following a dot.<br /> ex. Na2CO3·10H2O<br /> called sodium carbonate decahydrate<br /> deca- means 10 and hydrate means water<br /> Therefore there are 10 water molecules are associated with one formula unit of the compound.<br />
  48. 48. Formulas for Hydrates and Examples<br />
  49. 49. Analyzing a Hydrate<br />Must drive off the water by heating the compound<br />Substance remaining after heating is anhydrous (without water)<br />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<br />
  50. 50. Formula for a Hydrate<br />To determine the formula for a hydrate, you must determine the number of moles of water associated with one mole of the hydrate.<br />
  51. 51. Example Problem – Determining the Formula for a Hydrate<br />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.<br />
  52. 52. Example Problem – (cont.)<br />Known:<br />Mass of hydrated compound = 2.50 g CuSO4·xH2O<br />Mass of anhydrous compound = 1.59 g CuSO4<br />Molar mass = 18.02 g/mol H2O<br />Molar mass = 159.6 g/mol CuSO4<br />Unknown:<br />Formula for hydrate = ?<br />Name of hydrate = ?<br />
  53. 53. Example Problem – (cont.)<br />Subtract the mass of the anhydrous copper sulfate from the mass of the hydrated copper sulfate to determine the mass of water lost: <br /> mass of hydrates copper sulfate 2.50 g<br /> mass of anhydrous copper sulfate - 1.59 g<br /> mass of water lost 0.91 g<br />Calculate the number of moles of H2O and anhydrous CuSO4<br />1.59 g CuSO4 x 1 mol CuSO4 = 0.00996 mol CuSO4<br /> 159.6 g CuSO4<br />0.91 g H2O x 1 mol H2O = 0.050 mol H2O<br /> 18.02 g H2O<br />
  54. 54. Example Problem – (cont.)<br />Determine the value of x.<br />x = moles H2O = 0.050 mol H2O = 5.0 mol H2O = 5<br /> moles CuSO4 0.00996 mol CuSO4 1.0 mol CuSO4 1<br />The ratio of H2O to CuSO4 is 5 : 1, so the formula for the hydrate is CuSO4·5H2O, copper(II) sulfate pentahydrate.<br />
  55. 55. Practice Problems<br />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?<br />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?<br />
  56. 56. Uses of Hydrates<br />Drying agents: CaCl2, CaSO4<br />Storage of solar energy: Na2SO4·10H2O<br />