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AP Notes Chapter 4

AP Notes Chapter 4 . Stoichiometry. Atoms are “Letters” Compounds are “Words” Chemical Equations are the “Sentences of Chemistry” Reactants  Products. Experimental Evidence Shows that in chemical reactions. Mass and atoms conserved Moles not conserved. The Mole.

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AP Notes Chapter 4

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  1. AP Notes Chapter 4 Stoichiometry

  2. Atoms are “Letters” • Compounds are “Words” • Chemical Equations are the “Sentences of Chemistry” Reactants  Products

  3. Experimental Evidence Shows that in chemical reactions • Mass and atoms conserved • Moles not conserved

  4. The Mole • The mole is a number. • A very large number, but still, just a number. • 6.022 x 1023 of anything is a mole • A large dozen. • The number of atoms in exactly 12 grams of carbon-12. • Makes the numbers on the table the mass of the average atom.

  5. Balance to conserve atoms: • Mg(s) + O2(g)  MgO(s) • 2. Balance the equation in #1 for one mole of product.

  6. Interpretations of Equation:2Mg + O22MgO • Microscopic • Macroscopic2 moles magnesium and one mole oxygen produce 2 moles magnesium oxide + +

  7. Molar mass • Mass of 1 mole of a substance. • Often called molecular weight. • To determine the molar mass of an element, look on the table. • To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

  8. Find the molar mass of • CH4 • Mg3P2 • Ca(NO3)2 • Al2(Cr2O7)3 • CaSO4 · 2H2O

  9. Chemical Equations • Are sentences. • Describe what happens in a chemical reaction. • Reactants ® Products • Equations should be balanced. • Have the same number of each kind of atoms on both sides because ...

  10. Balancing equations CH4 + O2® CO2 + H2O Reactants Products 1 C 1 4 H 2 2 O 3

  11. Balancing equations CH4 + O2® CO2 + 2 H2O Reactants Products 1 C 1 4 H 2 4 2 O 3

  12. Balancing equations CH4 + O2® CO2 + 2 H2O Reactants Products 1 C 1 4 H 2 4 4 2 O 3

  13. Balancing equations CH4 + 2O2® CO2 + 2 H2O Reactants Products 1 C 1 4 H 2 4 4 4 2 O 3

  14. Abbreviations • (s) ¯ • (g) ­ • (aq) • heat • D • catalyst

  15. Practice • Ca(OH)2 + H3PO4® H2O + Ca3(PO4)2 • Cr + S8® Cr2S3 • KClO3(s) ® Cl2(g) + O2(g) • Solid iron(III) sulfide reacts with gaseous hydrogen chloride to form solid iron(III) chloride and hydrogen sulfide gas. • Fe2O3(s) + Al(s) ® Fe(s) + Al2O3(s)

  16. Write and balance the equation for the reaction in which ammonium nitrate decomposes to form dinitrogen oxide and water.

  17. Write and balance the equation for the reaction of calcium phosphate and phosphoric acid (hydrogen phosphate) to form calcium dihydrogen phosphate.

  18. Meaning • A balanced equation can be used to describe a reaction in molecules and atoms. • Not grams. • Chemical reactions happen molecules at a time • or dozens of molecules at a time • or moles of molecules.

  19. Stoichiometry Calculations based on chemical formulas or equations

  20. Stoichiometry • Given an amount of either starting material or product, determining the other quantities. • use conversion factors from • molar mass (g - mole) • balanced equation (mole - mole) • keep track.

  21. Map of the World of Stoichiometry atoms or molecules atoms or molecules Equation mole 1 mole 2 grams grams

  22. 1. How many grams of dinitrogen oxide are produced from the decomposition of 1.00 x 103 grams of ammonium nitrate?

  23. 2. Calculate the mass of chlorine required to react with 10.0 grams of sodium metal.

  24. 3. Calculate the mass of sodium nitrate that must decompose to produce 128 grams of sodium nitrite.

  25. 4. Calculate the mass of nitrogen needed to make 1000. grams of nitric acid by the process given.

  26. 5. What is the percent yield if 106 grams of sodium chlorite are isolated from the reaction of 202.3 grams of ClO2 with excess sodium hydroxide?

  27. 6. A volume of 3.42 mL of SiCl4 (density = 1.48 g/mL) reacts with excess hydrogen sulfide giving HSSiCl3. The hydrogen chloride produced reacts with 0.449 gram of NaOH. What is the percent yield?

  28. Examples • One way of producing O2(g) involves the decomposition of potassium chlorate into potassium chloride and oxygen gas. A 25.5 g sample of Potassium chlorate is decomposed. How many moles of O2(g) are produced? • How many grams of potassium chloride? • How many grams of oxygen?

  29. Examples • A piece of aluminum foil 5.11 in x 3.23 in x 0.0381 in is dissolved in excess HCl(aq). How many grams of H2(g) are produced? • How many grams of each reactant are needed to produce 15 grams of iron form the following reaction? Fe2O3(s) + Al(s) ® Fe(s) + Al2O3(s)

  30. Examples • K2PtCl4(aq) + NH3(aq)® Pt(NH3)2Cl2(s) + KCl(aq) • What mass of Pt(NH3)2Cl2 can be produced from 65 g of K2PtCl4 ? • How much KCl will be produced? • How much from 65 grams of NH3?

  31. Yield Limiting Reagent How much you get from an chemical reaction • Reactant that determines the amount of product formed. • The one you run out of first. • Makes the least product. • Book shows you a ratio method. • It works. • So does mine

  32. Limiting Reagent • To determine the limiting reagent requires that you do two stoichiometry problems. • Figure out how much product each reactant makes. • The one that makes the least is the limiting reagent.

  33. Limiting Reagent • Controls the amount of product generated in a reaction

  34. 1. What mass of S2Cl2 gas can be prepared from 32.0 grams of sulfur and 71.0 grams of chlorine gas?

  35. 2. What is the percent yield for a reaction in which 878 grams of ammonia gas are made from 1.00 x 103 g of nitrogen gas and 5.00 x 102 grams of hydrogen gas?

  36. 3. What remains in the reaction vessel after 150. grams of carbon tetrachloride liquid reacts with 100. grams of antimony(III) fluoride to form difluoro-dichloromethane (CCl2F2) and antimony(III) chloride solid?

  37. Example • Ammonia is produced by the following reaction N2 + H2® NH3 What mass of ammonia can be produced from a mixture of 100. g N2 and 500. g H2 ? • How much unreacted material remains?

  38. Excess Reagent • The reactant you don’t run out of. • The amount of stuff you make is the yield. • The theoretical yield is the amount you would make if everything went perfect. • The actual yield is what you make in the lab.

  39. Percent Yield • % yield = Actual x 100% Theoretical • % yield = what you got x 100% what you could have got

  40. Examples • Aluminum burns in bromine producing aluminum bromide. In a laboratory 6.0 g of aluminum reacts with excess bromine. 50.3 g of aluminum bromide are produced. What are the three types of yield.

  41. Examples • Years of experience have proven that the percent yield for the following reaction is 74.3% Hg + Br2® HgBr2 • If 10.0 g of Hg and 9.00 g of Br2 are reacted, how much HgBr2 will be produced? • If the reaction did go to completion, how much excess reagent would be left?

  42. Examples • Commercial brass is an alloy of Cu and Zn. It reacts with HCl by the following reaction • Zn(s) + 2HCl(aq) ® ZnCl2 (aq) + H2(g) • Cu does not react. When 0.5065 g of brass is reacted with excess HCl, 0.0985 g of ZnCl2 are eventually isolated.

  43. Pure O2 in Stoichiometry & Combustion Analysis xCO2 is absorbed C x H y Sample is burned completely to form CO2 and H2O y H2O is absorbed 2

  44. What is the balanced equation for the combustion of octane (C8H18)?

  45. Problem 35 p. 142 The aluminum in a 0.764 g sample of an unknown material was precipitated as aluminum hydroxide Al(OH)3, which was then converted to Al2O3 by heating strongly. If 0.127 g of Al2O3 is obtained from the 0.764 g sample, what is the mass percent of aluminum in the sample?

  46. Problem 30 p. 142Analysis of Mixtures A mixture of CuSO4 and CuSO4• 5H2O has a mass of 1.245 g. but, after heating to drive off all the water, the mass is only 0.832 g. What is the mass percent of CuSO4 • 5H2O in the mixture?

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