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Stoichiometry

Stoichiometry. Defined. Stoichiometry –The study of quantitative (measurable) relationships that exist in chemical formulas and chemical reactions. You must have: 1. The correct formulas – “criss-cross” or prefixes 2. Balanced Equations – Five types, verify law or conservation of mass.

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Stoichiometry

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  1. Stoichiometry

  2. Defined • Stoichiometry –The study of quantitative (measurable) relationships that exist in chemical formulas and chemical reactions. • You must have: • 1. The correct formulas – “criss-cross” or prefixes • 2. Balanced Equations – Five types, verify law or conservation of mass

  3. Analysis of a chemical reaction/equation • Consider: • N2H4 + 2H2O2 → N2 + 4H2O • Same as saying… • 1mole N2H4 + 2moles H2O2 → 1mole N2 + 4 moles H2O • The mole ratio is the “recipe” for the reaction.

  4. Mole-Mole Problems • Factor-Label Method – problem solving method based upon treating units in calculations as if they were algebraic factors. • N2H4 + 2H2O2 → N2 + 4H2O • Ex.1.4 moles of N2H4 gives how many moles of N2? • 1.4 moles N2H4 x mole N2 = moles N2 • mole N2H4

  5. Another Example • 3Zn + 2H3PO4 → Zn3(PO4)2 + 3H2 • How many moles of Zn3(PO4)2 will be produced from 2.18 moles of H3PO4? • 2.18 moles H3PO4 x moles Zn3(PO4)2 • mole H3PO4 • 1.09 moles Zn3(PO4)2

  6. Verifying the Law of Conservation of Matter • 2H2 + O2 → 2H2O • Calculate the mass of the reactants and the mass of the products • 2 mol H2 x g H2 + 1 mol O2 x g O2 = 36.04g • 1 mol H2 1 mol O2 • Or 2 mol H2O x g H2O = 36.04g • 1 mol H2O

  7. Mass – Mass Problems • Steps • Find molar mass of both substances • change mass to moles of given • use molar bridge (moles of given moles of unknown) • change moles to mass of unknown • ***Remember, coefficients in the balanced equation represent the relative number of moles of reactants and products!***

  8. Mass-Mass Problems • What mass of water is produced from 1.5 grams of glucose? • C6H12O6 + 6 O2 → 6 CO2 + 6 H2O • 1st. change mass to moles using molar mass of glucose. (180 g/mol) • 2nd. Use molar bridge to change from moles of given to moles of unknown. • 3rd change to grams of water

  9. The example of Mass-Mass • C6H12O6 + 6 O2 → 6 CO2 + 6 H2O • 1.5g C6H12O6 to g of water. • 0.9g H2O

  10. Another example • Ex. What mass of aluminum oxide is produced when 2.3 g of aluminum reacts with iron (III) oxide? • 4.3g Al2O3

  11. Mass-Volume problems • 1st mass of given • 2nd moles of given (molar mass) • 3rd moles of unknown (molar ratio) • 4th volume of gaseous unknown • If I have 125 g of Al2O3 how many L of Al do I have @ STP?

  12. Another Example • Find the mass of aluminum required to produce 1.32L of H2 gas @ STP • 2Al + 3 H2SO4 → Al2(SO4)3 + 3 H2

  13. Volume-volume problems • Same as mole-mole just using volumes instead! • If I have 15.5 L of N2 gas, how many L of H2 will react? • N2 + 3 H2 → 2 NH3

  14. 11-3 Limiting Reactants and Percent Yield • When chemicals combine, they are usually in nonstoichiometric proportions. That means there will be a limiting reactant. • Limiting Reactant will be completely used up in the reaction. • The other, leftover amount is said to be in excess. • **The quantities of products formed in a reaction are always determined by the quantity of limiting reactant. **

  15. LR continued • To find the limiting Reactant: • 1st solve 2 separate mass-mass problems. • 2nd see which number is smaller; this is your limiting reactant.

  16. Example of LR • Ex. 3.5 g of Cu is added to 6.0 g silver nitrate. Find the limiting reactant. • (Note: you can calculate the mass of either product, use the easier one to find the molar mass!) • Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag

  17. You can also determine how much of the substance in excess is left over. • 1st determine the LR and then use the amount of the LR and use that to find the mass of the other substance. • Ex. 2.0 g of NaHCO3 reacts with 0.5 g of H3C6H5O7 (citric acid) which is the LR and how much of the other will be left over? What volume of CO2 will be produced?

  18. example • 3 NaHCO3 + H3C6H5O7 → 3 CO2 + 3 H2O + Na3C6H5O7

  19. Percent yield • Percent yield = actual yield x 100 • Expected • Actual = what you got in lab • Expected = what it was supposed to be

  20. Example • A piece of copper with a mass of 5.00g is placed in a solution of AgNO3 . The silver metal produced has a mass of 15.2g. What is the percent yield for this rxn? • 1st calculate the mass of the expected silver. • Cu + 2AgNO3 2 Ag + Cu(NO3)2

  21. Example continued

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