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

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

3. Stoichiometry Quantity relationships based on chemical equations 3 Main Concepts: 1. 1. Chemical formula – molar ratio of atoms 2. 2. Chemical equations – molar ratio of compounds 3. 3. 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. Fe 2 S 3 (s) + HCl(aq) → FeCl 3 (s) + H 2 S(g) KClO 3 (s) → KCl(s) + O 2 (g) HNO 3 (l) + P 4 O 10 (s) → (HPO 3 ) 3 (l) + N 2 O 5 (g)

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

8. 3.2 Some Simple Patterns of Chemical Reactivity Combustion burning of a fuel in the presence of oxygen products of complete combustion: CO 2, H 2 O exothermic (produces heat)

9. Molecular view of methane combustion

10. 3.2 Some Simple Patterns of Chemical Reactivity Combustion C 3 H 8 (g) + 5 O 2 (g) → 3 CO 2 (g) + 4 H 2 O(g) 2 CH 3 OH(g) + 3 O 2 (g) → 2 CO 2 (g) + 4 H 2 O(g) Each C atom in fuel produces 1 mol CO 2 Each H atom in fuel produces ½ mol H 2 O

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 % composition by mass Mass of one type of atoms over mass of all atoms Mass of one type of atoms over mass of all atoms

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

15. amu impractical for lab use (too small) Avogadro’s number: 6.0221421 x 10 23 mol -1 The number of atoms in exactly 12 g of 12 C For conversions: 6.022 x 10 23 ?/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 3.4 Avogadro’s Number and the Mole

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

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

18. Practice with Avogadro’s # & the Mole

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

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 Combustion analysis 1 mol C in fuel → 1 mol CO 2 2 mol H in fuel → 1 mol H 2 OProblems

24. The combustion of propane, a hydrocarbon, produces 2.641 g CO 2 and 1.442 g H 2 O. Determine the emipirical formula of propane.

25. 3.6 Quantitative Information from Balanced Equations 3 Main Concepts: 1. 1. Chemical formula – molar ratio of atoms 2. 2. Chemical equations – molar ratio of compounds 3. 3. 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 CaCl 2 is produced from taking 2 antacid tablets, each containing 500. mg of CaCO 3 ? CaCO 3 (s) + 2 HCl(aq) → CO 2 (g) + H 2 O(l) + CaCl 2 (s)

30. From Sample Exercise 3.16, p. 98

31. How many grams of HCl are needed to react with 1000 mg of CaCO 3 ? CaCO 3 (s) + 2 HCl(aq) → CO 2 (g) + H 2 O(l) + CaCl 2 (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 NH 3 react with 2.50 g of O 2. NH 3 (g) + O 2 (g) → NO(g) + H 2 O(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) + S 8 (s) → 8 Ag 2 S(s)

36. How many grams are left over? 16 Ag(s) + S 8 (s) → 8 Ag 2 S(s) 16 Ag(s) + S 8 (s) → 8 Ag 2 S(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 ? FeCl 3 ·x H 2 O(s) → FeCl 3 (s) + x H 2 O(g) Δ