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Unit 4 – The Language of Chemistry: Part Deux

Unit 4 – The Language of Chemistry: Part Deux. Amadeo Avogadro. II.B.2(f) – Describe Avogadro’s hypothesis and use it to solve stoichiometric problems III.A.2(a) – Explain the meaning of mole and Avogadro’s number. Avogadro’s Hypothesis.

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Unit 4 – The Language of Chemistry: Part Deux

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  1. Unit 4 – The Language of Chemistry: Part Deux

  2. Amadeo Avogadro II.B.2(f) – Describe Avogadro’s hypothesis and use it to solve stoichiometric problems III.A.2(a) – Explain the meaning of mole and Avogadro’s number

  3. Avogadro’s Hypothesis • Equal amounts of gases at the same temperature contain equal numbers of molecules • Leads to definition of the “mole” • Mole Def: the number equal to the number of atoms in 12.01 grams of carbon • We will return to this idea later

  4. Atomic Mass Units (amu) • Since the mass of one atom is so tiny, it is more practical to use relative atomic masses • Carbon has been arbitrarily assigned a mass of exactly 12 atomic mass units • One atomic mass unit is defined as exactly 1/12 the mass of a carbon-12 atom

  5. Avogadro’s Number • Experimentally determined that there are 6.022 x 1023 atoms in exactly one mole • Remember this number! • 6.022 x 1023

  6. Molar Mass • The mass of one mole of a pure substance • Units: grams/mol • Equal to the mass of 6.022 x 1023 atoms of a pure substance • Molar mass of an element is numerically equal to the atomic mass of the element in atomic mass units

  7. Converting Mass, Moles, & Atoms/Molecules/Ions III.A.2(b) -Interconvert between mass, moles, and number of particles

  8. Gram/Mole/Atom Conversions

  9. Example: What is the mass in grams of 3.50 mol of copper?

  10. More Examples • What is the mass in grams of 2.25 mol of Fe? • What is the mass of 0.375 mol of K? • What is the mass of 16.3 mol of Ni? OR • How many moles are in 5.00 g of calcium? • How many moles are in 3.60 x 10-10 g Au?

  11. More Complex Examples • How many moles of Ag are in 3.01 x 1023 atoms of Ag? • How many atoms of Na are in 36.0 grams of Na? • How many oxygen atoms are in 180.18 g of glucose?

  12. Determining Chemical Formulas III.A.1(e) - Calculate the percent composition of a substance, given its formula or masses of each component element in a sample III.A.1(f) - Determine the empirical formulas and molecular formulas of compounds, given percent composition data or mass composition data III.A.2(c) - Distinguish between formula mass, empirical mass, molecular mass, gram molecular mass, and gram formula mass

  13. Empirical Formulas III.A.1(f) - Determine the empirical formulas and molecular formulas of compounds, given percent composition data or mass composition data

  14. Calculating an Empirical Formula Empirical Formula – consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole-number ratio of the different atoms in the compound • Always assume 100.0 g sample • Change percents to grams • Calculate moles of each element • Divide moles by smallest mole amount to determine ratio • Multiply to get whole numbers (if necessary)

  15. Determine the Empirical Formula of Aspirin • What is the empirical formula of aspirin? It has 4.48 % H, 60.00 % C, and 35.52 % O by mass.

  16. Example #1 • An oxide of aluminum is formed by the reaction of 4.151 g of aluminum with 3.692 g of oxygen. Calculate the empirical formula for this compound.

  17. Example #2 • When a 0.3546 g sample of vanadium metal is heated in air, it reacts with oxygen to achieve a final mass of 0.6330 g. Calculate the empirical formula of this vanadium oxide.

  18. Example #3 • A sample of lead arsenate, an insecticide used against the potato beetle, contains 1.3813 g of lead, 0.00672 g of hydrogen, 0.4995 g arsenic, and 0.4267 g of oxygen. Calculate the empirical formula for lead arsenate

  19. MolecularFormulas III.A.1(f) - Determine the empirical formulas and molecular formulas of compounds, given percent composition data or mass composition data III.A.2(c) - Distinguish between formula mass, empirical mass, molecular mass, gram molecular mass, and gram formula mass

  20. Calculating a Molecular Formula Molecular Formula – the actual formula of a molecular compound Example: C2H4 – ethylene C3H6 – cyclopropane Both have a 2H:1C ratio

  21. Determining Molecular Formula • Must know formula mass to calculate molecular formula • Divide experimental formula mass by empirical mass • Multiply subscripts by quotient

  22. Example #4 • Find the molecular formula of a compound with an empirical formula of CH and a formula mass of 78.110 amu

  23. Answer C – 12.011 H – 1.008 Total = 13.019 78.110 ÷ 13.019 ≈ 6 6(CH) = C6H6

  24. Example #5 • A white powder is analyzed and found to have an empirical formula of P2O5. The compound has a molar mass of 283.88g. What is the compounds molecular formula?

  25. Finding Molecular Formula from Empirical Formula • An unknown compound is found in tree sap. It has been shown to be composed of 40.0% carbon, 6.7% hydrogen and 53.3% oxygen by mass. It was also discovered that 5.00 moles of the material has a mass of 900 grams. • What is the molecular formula of this compound?

  26. Combustion Analysis

  27. Example #6 • What is the empirical formula of a hydrocarbon that produces 2.703 g CO2 and 1.108 g H2O when combusted?

  28. Example #7 • What is the empirical formula of a substance containing carbon, hydrogen, and oxygen if 1.000 g of substance produces 1.467 g CO2 and 0.6003 g H2O upon combustion? • The molar mass of the substance is 120 g/mol. What is the molecular formula?

  29. Example #8 • What is the molecular formula of a substance containing carbon, hydrogen, and oxygen if it has a molar mass of 234 g/mol and 0.360 g of substance produces 0.406 g CO2 and 0.250 g H2O upon combustion?

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