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Stoichiometry

Stoichiometry. Chemical Equations. Short hand way to represent chemical reactions H 2 + Cl 2 → HCl Symbols + = reacts with → = produces, yields Δ = adding energy (usually heat) ↔ = reversible aq - aqueous s - solid l - liquid g - gas. Balancing Reactions.

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Stoichiometry

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  1. Stoichiometry

  2. Chemical Equations • Short hand way to represent chemical reactions H2 + Cl2 → HCl • Symbols + = reacts with → = produces, yields Δ = adding energy (usually heat) ↔ = reversible aq - aqueous s - solid l - liquid g - gas

  3. Balancing Reactions • Law of Conservation of mass • Change coefficients not subscripts why? • Start balancing with elements other than H and O • Next balance H and then O last • Trial and Error

  4. Balancing Practice • __Al +__HCl → __AlCl3 + __H2 • __H3PO4 +__ HCl → __PCl5 + __ H2O • __ CO(g) + __H2(g) → __C8H18(l) + __H2O • __C2H6 + __O2 → __H2O + __CO2 • __(NH4)3PO4 + __Pb(NO3)4 → __Pb3(PO4)4 + __NH4NO3

  5. Types of Reactions • Combination – A + B → C N2(g) + 3H2(g) → 2NH3(g) • Decomposition – C → A + B 2KClO3(s) → 2KCl(s) + 3O2(g) Combustion – oxidation reaction that is exothermic Most often involves O2 CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) • Single Replacement – AB + C → A + BC Use activity series 2HCl + F2 → 2HF + Cl2 • Double Replacement – AB + CD → AD + CB 3 Driving Factors AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

  6. Formula Weights • Definition – sum of the atomic weights of each atom • Molecular weight – formula weight of a molecule • Calculate the formula weight of: • Ca(MnO4)2 • KH2PO4

  7. Percent Compostion • Definition - percent of mass contributed by each element • % Comp = # of a particular atom x atomic mass Total mass of the compound • Calculate the percentage of phosphorus in P4O10.

  8. Mole • Avogadro's # - 1 mole = 6.02x1023 objects • Representative Particles • Formula Units • Molecules • Atoms • Calculate the number of oxygen atoms in 1.5 moles of sodium carbonate.

  9. Mole • Molar mass – mass of one mole of a compound • Check if you are working with an isotope • Find the sum of the masses of the atoms • Unit is g/mol • Grams ← use g/mol → Moles ← use Avogadro's # → Representative particles

  10. Mole Practice • What is the mass, in grams of 6.33 mol of NaHCO3? • How many nitric acid molecules are in 4.20 grams of HNO3? How many oxygen atoms?

  11. Empirical and Molecular Formula • Finding Empirical Formula • Assume 100 g • Convert from grams to moles • Divide by smallest # of moles • Use results as subscripts • Convert Empirical to Molecular • Divide molar mass by empirical mass • Multiply empirical formula by result

  12. Molecular Formula Example • What is the molecular formula of a compound that is 71.65% Cl, 24.27% C, 4.07% H and has a molar mass of 98.96g/mol?

  13. Molecular Formula Practice • Caffeine contains 49.48% C, 5.15% H, 28.27% N, 16.49% by mass and has a molar mass of 194.2 g/mol. Determine the molecular formula of caffeine.

  14. Molecular Formula Example • Isopropyl alcohol is made up of C,H, and O. Combustion of 0.255g of this alcohol produces 0.521g CO2 and 0.306g H2O. What is the molecular formula if the molar mass is 60 g/mol.

  15. Molecular Formula Practice • Caproic acid, which is responsible for the foul odor of dirty socks, is composed of C, H, and O atoms. Combustion of a 0.255g sample of this compound produces 0.512 g CO2 and 0.209g H2O. What is the molecular formula is the molar mass is 116 g/mol?

  16. Calculations from Balances Equations • First ensure that equation is balanced. • Use mole ratios • Grams A → Moles A → Moles B → Grams B • When potassium chlorate is heated it decomposes in potassium chloride and oxygen. How many grams of oxygen can be prepared from 4.50g of potassium chlorate.

  17. Calculations practice • Milk of magnesia (Mg(OH)2) is often used as an antacid. It neutralizes excess hydrochloric acid secreted by the stomach. How many grams of milk of magnesia are needed to remove 150 mg of hydrochloric acid?

  18. Limiting Reagents • Limiting Reagents – reactant that runs out • Excess Reagents – reactant that is left over • Limiting reagent gives you the least amount of product

  19. Limiting Reagent Example • A strip of zinc metal weighing 2.00g is placed in an aqueous solution containing 2.50g of silver nitrate. Silver and zinc nitrate are produced. Which reactant is limiting? How many grams of each product will form? How grams of excess reagent will be left over?

  20. Limiting Reagent Practice • Nitrogen gas can prepared by passing gaseous ammonia over solid copper (II) oxide at high temperatures. The other products of the reaction are solid coper and water vapor. If a sample containing 18.1g NH3 is reacted with 90.4g CuO, which is the limiting reagent? How many grams of N2 are made? How much excess reagent is left?

  21. Yield • Theoretical Yield – amount of product formed when all of the LR is used up • Actual Yield – amount of product formed during an experiment • Percent Yield – actual yield x 100% theorectical yield

  22. Yield Example • When iron (III) oxide react with carbon monoxide to form iron and carbon dioxide. If you start with 150g of iron (III) oxide as the limiting reagent, what is the theoretical yield? If the actual yield of iron was 87.9g, what was the percent yield?

  23. Yield Practice • Methanol (CH3OH) is used as a fuel in race cars. It can be manufactured by a combination of gaseous carbon monoxide and hydrogen. If 68.5 kg of CO is reacted with 8.60 kg H2. Calculate the theoretical yield of methanol. If 3.57x104g of methanol is produced, what is the percent yield of methanol?

  24. Homework • 6, 8, 18, 30, 46, 48, 56, 58, 60, 70, 74, 76, 78

  25. Homework • 6, 8, 18, 30, 46, 48, 56, 58, 60, 74, 76, 78

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