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Unit 9

Unit 9. “the Mole". How do we measure how much of something we have?. Mass (g) - how much stuff Volume (L) – space the stuff occupies Count pieces or particles. Representative particles. The smallest pieces of a substance For an element it is an atom (Fe) Unless it is diatomic

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Unit 9

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  1. Unit 9 “the Mole"

  2. How do we measure how much of something we have? • Mass (g) - how much stuff • Volume (L) – space the stuff occupies • Count pieces or particles

  3. Representative particles • The smallest pieces of a substance • For an element it is an atom (Fe) • Unless it is diatomic • For a molecular compound it is a molecule (O2 or CO2). • For an ionic compound it is a formula unit (NaCl).

  4. But we have a problem counting those pieces! Most samples have LOTS of atoms, molecules, or formula units We want things to be simple! Scientists created a new unit

  5. The MOLE Defined as the number of carbon atoms in exactly 12 g of carbon-12. Avogadro’s number = the number of particles in one mole of a substance Avogadro’s number = 6.022 x 1023 particles

  6. The Mole 1 mole Fe = 6.022 x 1023 atoms 1 mole CO2 = 6.022 x 1023 molecules 1 mole NaCl = 6.022 x 1023 formula units

  7. Avogadro as a conversion factor 0.30 mol F atoms F 6.022 x 1023 1 mol F 0.30 mol F = _________ atoms F

  8. Calculation question • How many molecules of CO2 are there in 4.56 moles of CO2 ?

  9. Calculation question • How many moles of water is 5.87 x 1022 molecules?

  10. Calculation question • How many atoms of carbon are there in 1.23 moles of C6H12O6 ? • Umm…this one isn’t as easy… • How many atoms of C are in one molecule of C6H12O6 ? • 1 molecule = 6 atoms molecules 6 1 1.23 moles 6.022 x 1023 1 atoms C mole molecule

  11. Molar Mass • The mass of one mole (g/mol) • One mole is always 6.022 x 1023 pieces • But those pieces can have different masses • 1 dozen eggs vs. 1 dozen elephants • AKA • gram atomic mass (atoms) • gram molecular mass (molecules) • gram formula mass (compounds)

  12. Molar Mass/ Gram Atomic Mass • The mass of 1 mole of an element in grams • Equal to the average atomic mass found on the periodic table • We can write this as 12.01 g C = 1 mole C • And since 1 mole = 6.022 x 1023 atoms, we can also say that 6.022 x 1023 atoms C = 12.01 g C

  13. Molar mass/gram molecular mass • The mass of one mole of a molecule (formed by covalent bonds) • In 1 mole of H2O molecules there are 2 moles of H atoms and 1 mole of O atoms • To find the mass of one mole of a molecule, add the masses of the parts of that molecule

  14. Example • What is the mass of 1 mole of CH4? • 1 mole of C = 12.01 g • 4 mole of H x 1.01 g = 4.04g • 1 mole CH4 = 12.01 + 4.04 = 16.05g • The Gram Molecular mass or Molar Mass of CH4 is 16.05g

  15. Molar mass/ gram formula mass • The mass of one mole of an ionic compound (formed by ionic bonds) • Calculated the same way • What is the molar mass of Fe2O3? • 2 moles of Fe x 55.85 g = 111.70 g • 3 moles of O x 16.00 g = 48.00 g • Molar mass = 111.70 g + 48.00 g = 159.70g

  16. Molar mass as a conversion factor 3.50 mol Cu 63.55 g Cu 1 mol Cu 3.50 mol Cu = _________ g Cu

  17. Molar mass as a conversion factor • How many moles is 5.69 g of NaOH?

  18. Gases and the Mole

  19. Gases • It’s difficult to find the mass of a sample of gas, so we usually measure volume. • Two things affect the volume of a gas • Temperature and pressure • Must compare samples of gases at the same temperature and pressure

  20. Standard Temperature and Pressure • Avogadro's Hypothesis - at the same temperature and pressure, equal volumes of gas have the same number of particles • 0ºC and 1 atm pressure • abbreviated STP • At STP, 1 mole of gas occupies 22.4 L • Called the molar volume • Same volume but can still have different masses! (remember the density blocks…)

  21. Examples • Remember - 1 mole = 22.4 L • What is the volume of 4.59 moles of CO2 gas at STP? • How many moles is 5.67 L of O2 at STP? • What is the volume of 8.8g of CH4 gas at STP?

  22. Density of a gas • D = m /V • For a solid, the units were g/cm3 • For a gas, the units will be g / L • To find the density, we need to know the mass and the volume. • Assume you have 1 mole • At STP the volume is 22.4 L • The mass is the molar mass

  23. Examples • Find the density of CO2at STP. • Find the density of CH4 at STP.

  24. Going the other way • Given the density, we can find the molar mass of the gas • Again, assume you have 1 mole at STP, so V = 22.4 L. • What is the molar mass of a gas (at STP) with a density of 1.964 g/L? • D = m/V • What is the molar mass of a gas with a density of 2.86 g/L?

  25. Percent Composition • Describes the composition by mass of a substance • Like all percents • Part x 100 whole • To calculate - • Find the mass of each component • Divide by the total mass

  26. Example • Calculate the percent composition of silver in a compound that is made of 29.0 g of Ag and 4.30 g of S.

  27. % Composition from the Formula • If given the formula, assume you have 1 mole. 1. Calculate the total molar mass 2. Calculate the total mass of each element present 3. Calculate the percent composition: = (total mass of element present ÷ total molar mass) x 100

  28. Examples • Calculate the percent composition of each element in C2H4 • Calculate the percent composition of aluminum in aluminum carbonate.

  29. Percent to Mass • Multiply % composition by the total mass to find the mass of that component • How much aluminum is in 450 g of aluminum carbonate?

  30. Empirical Formula • The lowest whole number ratio of elements in a compound • Molecular formula - the actual ratio of elements in a compound • C2H4 molecular formula • CH2 empirical formula • C3H6 molecular formula • H2O molecular formula • H2O empirical formula

  31. Finding Empirical Formulas • Just find the lowest whole number ratio if given the molecular formula • C6H12O6 = • CH4N2 = • It is not just the ratio of atoms, it is also the ratio of moles

  32. Calculating Empirical Formulas • We can find empirical formula from the percent composition • Assume you have a 100 g sample • The percentages become grams. • Convert grams to moles. • Find lowest whole number ratio by dividing everything by the smallest number of moles. • Use the whole number ratio to write the formula

  33. Example • Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N. • Assume 100 g so… • 38.67 g C x 1mol C = 3.220 mole C 12.01 g C • 16.22 g H x 1mol H = 16.1 mole H 1.01 g H • 45.11 g N x 1mol N = 3.220 mole N 14.01 g N

  34. Example • The ratio is 3.220 mol C = 1 mol C 3.220 mol N 1 mol N • The ratio is 16.1 mol H = 5 mol H 3.220 mol N 1 mol N • 1C: 5H: 1N • C1H5N1 or CH5N

  35. Empirical to Molecular Formula • Since the empirical formula is the lowest ratio, the actual molecule could have the same mass or a higher mass • If given the empirical formula & the molar mass, calculate the molecular formula by… • Dividing the given molar mass by the mass of one mole of the empirical formula. • You will get a whole number. • Multiply the empirical formula by this whole number

  36. Example • A compound has an empirical formula of ClCH2 and a molar mass of 98.96 g/mol. What is its molecular formula? • Find the mass of one mole of the empirical formula or “1 empirical mole” • Divide the molar mass given by the mass of one empirical mole. • Multiply the empirical formula by that whole number

  37. Percent comp  molecular formula • Multiply the percent by the molar mass • This gives you the mass of that element in 1 mole of the compound • Change this to moles • You will get whole numbers • These are the subscripts

  38. Percent comp  Molecular Formula Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula? 49.48% C x 194 g = Change that to moles 5.15% H x 194 g = 28.87% N x 194 g = 16.49% O x 194 g =

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