Stoichiometry Calculations with Chemical Formulas and Equations

1. StoichiometryCalculations with Chemical Formulas and Equations Chapter 3 BLB 12th

2. Expectations • Balance chemical equations. • g ↔ moles ↔ molecules ↔ atoms • Find empirical and molecular problems. • Calculate amounts of reactants and products. • Calculate theoretical and percent yield.

3. Stoichiometry Quantity relationships based on chemical equations 3 Main Concepts: • Chemical formula – molar ratio of atoms • Chemical equations – molar ratio of compounds • Law of Conservation of Mass: mass of reactants = mass of products

4. 3.1 Chemical Equations Components: • reactants → products • Physical states (s, l, g, aq) • Reaction conditions (heat Δ, light, solvents, etc.) • Coefficients determine molar ratios. The number of moles of each type of atom must be the same on each side. Balancing: • By inspection • Use coefficients; don’t change chemical formulas

5. Coefficients vs. Subscripts

6. Fe2S3(s) + HCl(aq) → FeCl3(s) + H2S(g) KClO3(s) → KCl(s) + O2(g) HNO3(l) + P4O10(s) → (HPO3)3(l) + N2O5(g)

7. 3.2 Some Simple Patterns of Chemical Reactivity Combination and Decomposition combination: A + B → C 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) decomposition: C → A + B 2 NaN3(s) → 2 Na(s) + 3 N2(g)

8. 3.2 Some Simple Patterns of Chemical Reactivity Combustion • burning of a fuel in the presence of oxygen • products of complete combustion: CO2, H2O • exothermic (produces heat)

9. Molecular view of methane combustion

10. 3.2 Some Simple Patterns of Chemical Reactivity Combustion C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) 2 CH3OH(g) + 3 O2(g) → 2 CO2(g) + 4 H2O(g) • Each C atom in fuel produces 1 mol CO2 • Each H atom in fuel produces ½ mol H2O

11. 3.3 Formula Weights Formula and Molecular Weights (amu) formula weight – general molecular weight – molecules formula unit weight – ionic compound - sum of the atomic masses of each atom in chemical formula

12. % Composition • % composition by mass • Mass of one type of atoms over mass of all atoms

14. 3.4 Avogadro’s Number and the Mole • Word association: pair – dozen – case – ream –

15. 3.4 Avogadro’s Number and the Mole • amu impractical for lab use (too small) • Avogadro’s number: 6.0221421 x 1023 mol-1 • The number of atoms in exactly 12 g of 12C • For conversions: 6.022 x 1023 ?/mol, where ? can equal atoms, molecules, ions, etc. • 1 mole = Avogadro’s number of anything • molar mass – mass in grams of one mole of a substance, which is equal to the atomic mass in amu; g/mol

16. 3.4 Avogadro’s Number and the Mole • Atoms & compounds have different masses, thus the mass of 1 mole of atoms & compounds are different.

17. 3.4 Avogadro’s Number and the Mole Conversions: • g → mol divide by molar mass • mol → g multiply by molar mass • Abbreviations: mole – mol molarity – M

18. Practice with Avogadro’s # & the Mole

19. 3.5 Empirical Formulas from Analyses • Empirical formula – smallest whole number ratio of atoms • Molecular formula – actual ratio of atoms in a compound; multiple of the empirical formula; must know molar mass of compound • Use % composition to find formula Problems

20. A once-used gasoline additive contains 49.5% C, 3.2% H, 22.0% O, and 25.2% Mn. Determine the emipirical formula of this compound.

21. Azulene, a hydrocarbon, contains 93.71% C. Its molar mass is ~128 g/mol. Determine the emipirical and molecular formulas for azulene.

22. 3.5 Empirical Formulas from Analyses Summary: • % data → grams • Grams → moles • Moles → molar ratio → empirical formula • Empirical formula → molecular formula

23. 3.5 Empirical Formulas from Analyses • Combustion analysis • 1 mol C in fuel → 1 mol CO2 • 2 mol H in fuel → 1 mol H2O Problems

24. The combustion of propane, a hydrocarbon, produces 2.641 g CO2 and 1.442 g H2O. Determine the emipirical formula of propane.

25. 3.6 Quantitative Information from Balanced Equations 3 Main Concepts: • Chemical formula – molar ratio of atoms • Chemical equations – molar ratio of compounds • Law of Conservation of Mass: mass of reactants = mass of products

26. What’s balanced in a balanced equation?

27. Stoichiometry Problems Use these 4 steps as a guide: (p. 97) • Write & balance chemical equation. • Convert to moles. • Apply molar ratio. • Convert from moles to quantity desired (mass, volume, etc.)

28. Stoichiometry Problems

29. How many grams of CaCl2 is produced from taking 2 antacid tablets, each containing 500. mg of CaCO3?CaCO3(s) + 2 HCl(aq) → CO2(g) + H2O(l) + CaCl2(s)

30. From Sample Exercise 3.16, p. 98

31. How many grams of HCl are needed to react with 1000 mg of CaCO3?CaCO3(s) + 2 HCl(aq) → CO2(g) + H2O(l) + CaCl2(s)

33. 3.7 Limiting Reactants • Limiting reactant – reactant that is completely consumed; limits the amount of product that can be formed • Theoretical yield – calculated yield of a product based on limiting reactant • Percent yield

34. 78. Calculate the theoretical yield (in grams) of NO when 2.00 g of NH3 react with 2.50 g of O2. NH3(g) + O2(g) → NO(g) + H2O(g)

35. Silver metal reacts with elemental sulfur according to the reaction below. If 2.0 g each of silver and sulfur react, what is the theoretical yield (in grams) of silver(I) sulfide? How many grams are left over? 16 Ag(s) + S8(s) → 8 Ag2S(s)

36. How many grams are left over? 16 Ag(s) + S8(s) → 8 Ag2S(s)

37. 84. When hydrogen sulfide gas is bubbled into a solution of sodium hydroxide, the reaction forms sodium sulfide and water. How many grams of sodium sulfide are formed if 1.25 g of hydrogen sulfide is bubbled into a solution containing 2.00 g of sodium hydroxide, assuming that the sodium sulfide is made in 92.0% yield?

39. When 0.750 g iron(III) chloride hydrate is heated, 0.300 g of steam is produced. What is the value of x ? FeCl3·xH2O(s) → FeCl3(s) + x H2O(g) Δ