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Molar Ratios From Empirical Formulas

Molar Ratios From Empirical Formulas. Empirical formula the smallest whole number ratio of atoms (or moles of atoms) of each element in a substance H 2 O 2 H atoms for every 1 O atom 2 moles of H atoms for every mole of O atoms. Molar Ratios From Empirical Formulas.

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Molar Ratios From Empirical Formulas

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  1. Molar Ratios From Empirical Formulas • Empirical formula • the smallest whole number ratio of atoms (or moles of atoms) of each element in a substance • H2O • 2 H atoms for every 1 O atom • 2 moles of H atoms for every mole of O atoms

  2. Molar Ratios From Empirical Formulas • The relative number of moles of each element in a substance can be used as a conversion factor called the molar ratio. • Molar ratio = moles element A mole of substance • Molar ratio = moles element A moles element B or

  3. Molar Ratios From Empirical Formulas • Fe2O3: • Molar Ratio = 2 moles of Fe mole Fe2O3 • Molar Ratio = 3 moles O mole Fe2O3 • Molar Ratio = 2 moles Fe 3 moles O

  4. Molar Ratios From Empirical Formulas • Molar ratios can be used to determine the number of moles of a particular element in a given substance. Moles A Molar Ratio Moles B

  5. Molar Ratios From Empirical Formulas Example: How many moles of Na+ ions are present in 2.5 moles of Na2SO4 ?

  6. Grams B Moles B Molar Ratios From Empirical Formulas • Remember, once you find the number of moles of a substance present, you can use: • Molar mass to find the number of grams • Avogadro’s number to find the number of atoms, ions, or molecules Moles A Molar Ratio N Molar mass Atoms B

  7. Molar Ratios From Empirical Formulas Example: What is the mass of iron present in 4.00 moles of Fe2O3?

  8. 2 molecules 1 molecules 2 molecules 2(6.02x1023) molecules 6.02x1023 molecules 2(6.02x1023) molecules 2 moles 1 mole 2 moles Quantitative Information from Chemical Equations • Coefficients in a balanced equation • number of molecules (formula units, etc) • number of moles 2 H2 + O2 2 H2O

  9. Quantitative Information from Chemical Equations • In industry, aspirin is prepared by: C7H6O3 + C4H6O3 C9H8O4 + HC2H3O2 Salicylic acid Acetic anhydride Acetyl-salicylic acid Acetic acid • Chemists want to use the right amount of reactants in order to minimize “left over” reactants that contaminate the product.

  10. Quantitative Information from Chemical Equations • The coefficients in a balanced chemical equation can be used to relate the number of moles of each substance involved in a reaction. Molar ratios mol reactantmol reactantmol product mol reactant mol product mol product

  11. Molar Ratios from Chemical Equations For the reaction: N2 + 3 H2 2 NH3 • 3 different molar ratios (and their inverses) can be written: 1 mol N2 1 mol N2 3 mol H2 3 mol H2 2 mol NH3 2 mole NH3

  12. Molar Ratios from Chemical Equations MOLAR RATIOS = CONVERSION FACTORS • Moles of product that can be formed from a certain number of moles of reactant(s) • Moles of reactants needed to form a certain number of moles of product • Moles of reactant 2 needed to completely react with reactant 1

  13. Quantitative Information from Chemical Equations Example: If you have 1.0 mole of H2, how many moles of NH3 can you produce? N2 + 3 H2 2 NH3 Note: Make sure your equation is balanced!

  14. Quantitative Information from Chemical Equations Example: If you have 1.0 mole of H2, how many moles of N2 will be required to completely react all of the H2? N2 + 3 H2 2 NH3

  15. Quantitative Information from Chemical Equations Example: How many moles of N2 are needed to produce 0.50 moles of NH3? N2 + 3 H2 2 NH3

  16. Quantitative Information from Chemical Equations • Finding # of moles is great BUT • You don’t measure out moles in the lab! • Chemists use a balance to measure the mass of a substance used or produced in a reaction. • How can you determine the mass of reactants or products?

  17. Quantitative Information from Chemical Equations • Use the molar mass to convert from moles to grams • The number of grams of a substance per mole

  18. Molar mass Mass (g) Compound A Moles Compound A Molar ratio Molar mass Mass (g) Compound B Moles Compound B Quantitative Information from Chemical Equations “The MAP”

  19. Quantitative Information from Chemical Equations Example: How many grams of water will be produced by the complete combustion of 10.0 g of propane?

  20. Combustion Reactions – Revisited! • You should be able to write a balanced equation for the combustion of an organic compound or a metal. • Organic compounds: • Metals: Not bal.

  21. Quantitative Information from Chemical Equations Example: Hydrofluoric acid can’t be stored in glass because it attacks the silicates in the glass: Na2SiO3 (s) + 8 HF (aq)  H2SiF6 (aq) + 2 NaF (aq) + 3 H2O How many grams of HF are needed to dissolve 55.0 g of Na2SiO3?

  22. Quantitative Information from Chemical Equations

  23. Making Bologna Sandwiches • Suppose you were going to make bologna sandwiches: + + 2 Bread + 1 Bologna  1 sandwich

  24. Making Bologna Sandwiches 2 Bread + 1 Bologna  1 sandwich + We can only make 4 sandwiches because we don’t have enough bologna! Bologna = limiting reagent or limiting reactant

  25. Limiting Reagent or Limiting Reactants • Similar situations occur in chemical reactions when one of the reactants is used up before the others. • No further reaction can occur • The excess reactant(s) are “leftovers.”

  26. Limiting Reagent or Limiting Reactants 2 H2 (g) + O2 (g)  2 H2O (l) If we react 10 moles of H2 with 7 moles of O2, not all of the O2 will react because we will run out of H2 first!

  27. Limiting Reagent or Limiting Reactants 10 H2 7 O2 10 H2O + 2 O2

  28. Limiting Reagent or Limiting Reactants • Limitingreagent(limiting reactant): • the reactant that is completely consumed in a reaction • determines or limits the amount of product formed.

  29. Limiting Reagent or Limiting Reactants • Three approaches to identifying the limiting reactant: • Compare the number of moles of each reactant needed with the number of moles of each reactant available • Calculate the number of grams of product that each reactant could form • Reactant that forms the least amount of product will be the limiting reagent. OR

  30. Limiting Reagent or Limiting Reactants • Three approaches to identifying the limiting reactant (cont.) • Pick one of the reactants and calculate the grams of the other reactant needed to exactly react with the one you picked OR

  31. Limiting Reagent or Limiting Reactants First Method: 1) Convert the mass of each reactant to moles. 2) Pick one of the reactants (it doesn’t matter which one) and calculate the number of moles of the other reagent needed to completely react with the one chosen. 3) Compare the # moles needed vs. # moles available

  32. Limiting Reagent or Limiting Reactants Second Method: 1) Calculate the number of grams of product that each reactant could form. 2) Limiting reagent = reactant that forms the smallest mass of product.

  33. Limiting Reagent or Limiting Reactants Third Method: • Pick one reactant. • Calculate the grams of the other reactant needed. • Compare the grams needed with the grams available.

  34. Limiting Reagent or Limiting Reactants Example: If 10.0 grams of H2 are mixed with 75.0 grams of O2, which reactant is the limiting reagent? 2 H2 (g) + O2 (g)  2 H2O (l)

  35. Limiting Reagent or Limiting Reactants Method 1

  36. Limiting Reagent or Limiting Reactants Method 2

  37. Limiting Reagent or Limiting Reactants Method 3

  38. Limiting Reagent or Limiting Reactants • Once you find the limiting reagent, you can also find the amount of product that can be formed in the reaction. • If you used the second method for identifying the limiting reagent, you’ve already done this! Molar mass Molar ratio Moles limiting reagent Moles product grams product

  39. Limiting Reagent or Limiting Reactants Example: How many grams of water are produced by burning 5.00 g of methane in the presence of 5.00 g of oxygen? CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)

  40. Example • Since the mass of both reactants is given, you need to check to see if one of them is a limiting reagent.

  41. Example

  42. Yield • TheoreticalYield: • the quantity of product that is calculated to form when all of the limiting reagent reacts • Actualyield: • the amount of product actually obtained in a reaction

  43. Yield • Often, the actual yield is less than the theoretical yield: • reactants may not react completely • i.e. the reaction does not go to completion • “by-products” may form • unwanted side reactions (competing reactions) • difficulty isolating and purifying the desired product

  44. Yield • Percent Yield: • relates the actual yield and the theoretical yield • % Yield = actual yield x 100% theoretical yield

  45. Yield Example: In a certain reaction between H2 and CO to form methanol, the theoretical yield is 83.3 g of CH3OH. If the actual yield of the reaction was 81.5 g, what was the % yield?

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