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Practice Problems Quantitative Aspects

Practice Problems Quantitative Aspects. How to keep things straight when solving quantitative problems:. First identify what you are being asked to find If it is not clear to you what the problem asks, list down the data or information you are provided with

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Practice Problems Quantitative Aspects

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  1. Practice ProblemsQuantitative Aspects

  2. How to keep things straight when solving quantitative problems: • First identify what you are being asked to find • If it is not clear to you what the problem asks, list down the data or information you are provided with • Connect this with a formula, a definition or a topic you have learned about • Find out if you have all the information you need to use the formula. If not, find other information or other formulas that can be used to supply the missing information • Check out the units. Find conversion factors or equivalences if necessary

  3. Typical quantitative problems • Questions based on defined terms • Atomic mass, molar mass • Average atomic mass of an element • Avogadro’s number • Molar volume • Questions based on formulas • Density, volume, mass, # of moles • % composition • Empirical and molecular formulas • Balancing and Interpreting Chemical Equations

  4. Questions based on the Chemical formula • Given the formula, find the amount of one of its components • Given a component in a formula, find the amount • Questions based on the balanced chemical equation: Proportions, mole ratios

  5. FYI: Terms that can be used as factors for dimensional analysis • Conversion units(equivalences) • Percentages (%) • Density • Molar mass • Molar volume • Avogadro’s number • Mass ratios, mole ratios

  6. 3.2 Avogadro’s Number & the Molar Mass of an Element • How many moles of magnesium (Mg) are there in 87.3 g of Mg?

  7. 3.2 Avogadro’s Number & the Molar Mass of an Element • Calculate the number of grams of lead (Pb) in 12.4 moles of lead.

  8. 3.2 Avogadro’s Number & the Molar Mass of an Element • Calculate the number of atoms in 0.511 g of potassium (K).

  9. 3.3 Molecular Mass • What is the molecular mass of methanol (CH4O)?

  10. 3.3 Molecular Mass • Calculate the number of moles of chloroform (CHCl3) in 198 g of chloroform.

  11. 3.3 Molecular Mass • How many H atoms are in 72.5 g of isopropanol (rubbing alcohol), C3H8O?

  12. 3.5 Percent Composition of Compounds • Calculate the percent composition by mass of each of the elements in sulfuric acid (H2SO4).

  13. 3.5 Percent Composition of Compounds • Determine the empirical formula of a compound having the following percent composition by mass: K: 24.75%; Mn: 34.77%; O: 40.51%.

  14. 3.5 Percent Composition of Compounds • Calculate the number of grams of Al in 371 g of Al2O3.

  15. 3.6 Experimental Determination of Empirical Formulas • A sample of a compound containing boron (B) and hydrogen (H) contains 6.444 g of B & 1.803 g of H. The molar mass of the compound is about 30 g. What is its molecular formula?

  16. 3.7 Chemical Reactions & Chemical Equations • Balance the equation representing the reaction between iron(III) oxide, Fe2O3, & carbon monoxide (CO) to yield iron (Fe) & carbon dioxide (CO2).

  17. 3.10 Amounts of Reactants & Products (Intro to Chapter 4) Methanol (CH3OH) burns in air according to the equation 2CH3OH + 3O2 → 2CO2 + 4H2O If 209 moles of methanol are used up in a combustion process, what is the number of moles of H2O produced?

  18. 3.11 Practice Writing and Balancing Equations; Identifying Type of Rxn • Worksheet posted • See also Notes on Types of Chemical Equations (Discussed in Lab March 4) • END Test 2 Practice

  19. CHAPTER 4 Start: Amounts of Reactants & Products The reaction between nitric oxide (NO) and oxygen to form nitrogen dioxide (NO2) is a key step in photochemical smog formation: 2NO(g) + O2(g) → 2NO2(g) How many grams of O2 are needed to produce 2.21 g of NO2?

  20. Limiting Reagents CHAPTER 4 The reaction between aluminum and iron(III) oxide can generate temperatures approaching 3000°C and is used in welding metals: 2Al + Fe2O3 → Al2O3 + 2Fe In one process, 124 g of Al are reacted with 601 g of Fe2O3. (a) Calculate the mass (in grams) of Al2O3 formed. (b) How much of the excess reagent is left at the end of the reaction?

  21. 3.10 Reaction Yield Industrially, vanadium metal, which is used in steel alloys, can be obtained by reacting vanadium(V) oxide with calcium at high temperatures: 5Ca + V2O5 → 5CaO + 2V In one process, 1.54 x 103 g of V2O5 react with 1.96 x 103 g of Ca. (a) Calculate the theoretical yield of V. (b) Calculate the percent yield if 803 g of V are obtained. CHAPTER 4

  22. Solutions and Reactions • The important thing to remember is that substances exist in different forms in solution CHAPTER 4

  23. 4.2 Precipitation Reactions • Classify the following ionic compounds as soluble or insoluble: a) CuS b) Ca(OH)2 c) Zn(NO3)2 CHAPTER 4

  24. 4.2 Precipitation Reactions • Predict the precipitate produced by mixing an Al(NO3)3 solution with a NaOH solution. Write the net ionic equation for the reaction. CHAPTER 4

  25. 4.3 Acid-Base Reactions • Classify each of the following species as a Bronsted acid or base: a) SO42- b) HI CHAPTER 4

  26. 4.4 Oxidation-Reduction Reactions • Assign oxidation numbers to all of the elements in the following compound and ion: a) PF3 b) MnO4- CHAPTER 4

  27. 4.4 Oxidation-Reduction Reactions • Identify the following redox reactions by type: a) Fe + H2SO4 → FeSO4 + H2 b) S + 3F2 → SF6 c) 2CuCl → Cu + CuCl2 d) 2Ag + PtCl2 → 2AgCl + Pt CHAPTER 4

  28. 4.5 Concentration of Solutions • What is the molarity of an 85.0 mL ethanol (C2H5OH) solution containing 1.77 g of ethanol? CHAPTER 4

  29. 4.5 Concentration of Solutions • What volume (in milliliters) of a 0.315 M NaOH solution contains 6.22 g of NaOH? CHAPTER 4

  30. 4.5 Concentration of Solutions • How would you prepare 2.00 x 102 mL of a 0.866 M NaOH solution, starting with a 5.07 M stock solution? CHAPTER 4

  31. Solution Stoichiometry • Stoichiometry of solutions relies on the same pattern of using the mole ratios from balanced chemical equations • When given grams moles can be calculated as usual from molar masses • When given solutions, moles can be calculated from M and V CHAPTER 4

  32. 4.6 Gravimetric Analysis • A sample of 0.3220 g of an ionic compound containing the bromide ion (Br-) is dissolved in water and treated with an excess of AgNO3. If the mass of the AgBr precipitate that forms is 0.6964 g, what is the percent by mass of Br in the original compound? CHAPTER 4

  33. 4.7 Acid-Base Titrations • How many grams of KHP (KHC8H4O4 ) are needed to neutralize 18.64 mL of a 0.1004 M NaOH solution? KHC8H4O4(aq) + NaOH(aq) → KNaC8H4O4(aq) + H2O(l) KHC8H4O4 molar mass = 204.2 g CHAPTER 4

  34. 4.7 Acid-Base Titrations • How many milliliters of a 1.28 M H2SO4 solution are needed to neutralize 60.2 mL of a 0.427 M KOH solution? CHAPTER 4

  35. 4.8 Redox Titrations • How many milliliters of a 0.206 M HI solution are needed to reduce 22.5 mL of a 0.374 M KMnO4 solution according to the following equation: 10HI + 2KMnO4 + 3H2SO4 → 5I2 + 2MnSO4 + K2SO4 + 8H2O CHAPTER 4

  36. According to the following balanced reaction, how many moles of NO are formed from 8.44 moles of NO2 if there is plenty of water present?3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) CHAPTER 4

  37. A 14.01 g sample of N2 reacts with 3.02 g of H2 to form ammonia (NH3). If ammonia is the only product, what mass of ammonia is formed? CHAPTER 4

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