Study Guide for Chapters 4-6

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Study Guide for Chapters 4 -6 Chapter 4: Chemical Reactions and Chemical Quantities  State the meaning of each part of a chemical equation o What do the coefficients in a chemical equation stand for  Translate word problems into chemical equations o memorize the 7 diatomic elements: Br 2 , I 2 , N 2 , Cl 2 , H 2 , O 2 , F 2  Balance chemical equations  Calculate the amount of product obtained from a given amount of reactant o Convert grams of reactant to moles using molar mass o Convert moles of reactant to moles of product using mole ratios o Converts moles of product to grams using molar mass  Identify the limiting reactant  Solve problems involving a limiting reactant  Calculate the theoretical yield in a reaction  Calculate the percent yield of a reaction, given the actual yield Chapter 5: Introduction to Solutions and Aqueous Reactions  Draw the structure of water, and label partial charges  Describe: o electronegativity o polar o nonpolar o like dissolves like rule  Define: o solute o solvent o aqueous solution o strong electrolyte (give examples) o weak electrolyte (give examples) o nonelectrolyte (give examples) o strong acids (give examples) o weak acids (give examples) o strong bases (give examples) o weak bases (give examples) o stock solution o standard solution  Know the molarity equation o determine the molarity of a solution given the mass of solute and volume of solution o determine the molarity of each ion in a solution o determine the mass and/or volume of reagents necessary to prepare a solution of a given molarity o Describe the process of forming a standard aqueous solution:  from a solid  by dilution  memorize the dilution equation: M 1 V 1 = M 2 V 2

Precipitation Reactions:  Use solubility rules to predict determine which ionic compounds will be soluble in water o Write dissolution equations for ionic compounds, using the correct charge and number of each ion in solution o calculate the concentration of ions in solution  Use solubility rules to predict the products of reactions in aqueous solution  Write molecular, ionic, and net ionic equations to describe precipitation reactions in solution  Solve problems involving precipitation reactions o use the molarity equation in stoichiometry problems Acid-Base Reactions:  Define o Acid o Base o Analyte o Titrant o Endpoint o Equivalence point o monoprotic, diprotic, polyprotic  Predict the products of acid-base reactions o Write molecular, ionic, and net ionic equations to describe acid-base reactions in water  strong acid, strong base, weak acid, weak base  Use the molarity equation to solve acid-base titration problems o Calculate the molarity and/or volume of the analyte o calculate the molar mass of the analyte Oxidation-Reduction Reactions:  Define: o oxidation o reduction o oxidizing agent o reducing agent  Assign oxidation states to atoms in a compound  Identify which elements are oxidized and which are reduced in a reaction  Identify which elements are acting as oxidizing agents and reducing agents in a reaction Chapter 6: Gases  Use Kinetic Molecular Theory (KMT) to describe the behavior of gases on the molecular and macroscopic scale  Know the Ideal Gas Law: PV = nRT o Solve problems using the ideal gas law o solve problems involving changing conditions by modifying the ideal gas law o Use the ideal gas law to solve for moles of a gas in a stoichiometry problem  Know STP conditions: 1 atm pressure, 0 o C  Use the relationship 1 mole of a gas = 22.4L at STP in stoichiometry calculations  Calculate the density of a gas at STP

 Calculate the density of a gas under non-STP conditions  Know Dalton’s Law of Partial Pressures o Solve problems involving mixtures of gases o Solve problems involving collecting a gas over water  Know the mole fraction equation o solve for the partial pressure of a gas within a mixture of gases  Calculate o kinetic energy of a gas depending on its temperature o root mean square velocity of a gas depending on its temperature and molar mass  Compare the rate of effusion of two gases o qualitatively o quantitatively  Compare the distance of diffusion of two gases o qualitatively o quantitatively  Real Gases o Describe why and how gases deviate from ideal behavior o under what conditions do you see the most deviation from ideal behavior? Chapter 4 Questions: 1. Balance the following chemical equations: a. ___AgI + ___Na 2 S  ___Ag 2 S + ___NaI b. ___Mg(OH) 2 + ___HCl  ___MgCl 2 + ___H 2 O c. ___TiCl 4 + ___H 2 O  ___TiO 2 + ___HCl d. ___HNO 2  ___HNO 3 + ___NO + ___H 2 O 2. The following reaction is used to form lead(II) iodide crystals. What mass of lead(II) iodide could be formed from 250.0g of lead(II) acetate? Pb(C 2 H 3 O 2 ) 2 + 2KI  PbI 2 + 2KC 2 H 3 O 2 3. A reaction combines 64.81g of silver nitrate with 92.67g of potassium bromide: AgNO 3 + KBr  AgBr + KNO 3 a. Which reactant is limiting? Which is excess? b. What is the theoretical yield of silver bromide? c. If the actual yield of silver bromide is 14.77g, what was the percent yield? Chapter 5 Questions: 1. Write the dissociation equation for K 3 PO 4 in water. 2. Which of the following substances would be soluble in water? a. CH 3 CH 2 OH b. C 6 H 6 c. LiBr 3. Classify the following as strong, weak, or nonelectrolytes: a. CH 3 CH 2 OH b. HCl c. NH 3 4. What is the molarity of 49.74g of H 2 SO 4 dissolved in enough water to make 500.mL of solution? 5. Calculate the concentration of sodium ions in 0.25M Na 3 PO 4 6. Describe how you would prepare 250.mL of 1.0M Na 2 SO 4 7. Describe how you would prepare 500.mL of 1.0N H 2 SO 4 from 17.8M H 2 SO 4

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8. A stock solution of sodium hydroxide is prepared by dissolving 120.0g of NaOH in 500.0mL of water. What is the molarity of the stock solution? 9. How many milliliters of 0.50M Ca(OH) 2 are required to react with 30.mL of a 0.12M HCl solution? 10. Predict whether a precipitate will form in each reaction below. Write the balanced molecular equation for each: a. NaCl(aq) + Hg 2 (NO 3 ) 2 (aq)  b. Ca(OH) 2 (aq) + Na 2 CO 3 (aq)  c. Na 2 S(aq) + FeCl 3 (aq)  11. A solution contains Ag + , Pb 2+ , and Fe 3+ . If you want to precipitate the Pb 2+ selectively, what anion would you choose? 12. Predict the precipitate and write the molecular, ionic, and net ionic equation for the following reaction: a. aqueous sodium sulfide reacts with aqueous copper(II) nitrate 13. What mass of Mg(OH) 2 is produced when 100.mL of 0.42M Mg(NO 3 ) 2 is added to excess NaOH solution? 14. Calculate the mass of CaSO 4 produced when 10.mL of 6.0M H 2 SO 4 is added to 100.mL of 0.52M Ca(NO 3 ) 2 . 15. What volume of 0.1379M HCl is required to neutralize 10.0mL of 0.2789M NaOH solution? 16. A titration is done using 0.1302M NaOH to determine the molar mass of a monoprotic acid. If 1.863g of the acid requires 70.11mL of the NaOH solution, what is the molar mass of the acid? 17. Predict the products and write the molecular, ionic, and net ionic equation for the following: a. H 2 SO 4 (aq) + NaOH(aq)  b. HClO 4 (aq) + NH 3 (aq)  c. HC 2 H 3 O 2 (aq) + LiOH(aq)  18. What volume of 0.2M NaOH is required to neutralize 50.mL of 0.1M H 2 SO 4 ? 19. Assuming the stoichiometry is 1:1, what is the concentration of an unknown acid if a 20.0mL of acid is neutralized by 33.4mL of 0.250M base? 20. Determine the oxidation number of each atom in the following compounds: a. MgBr 2 b. Na 2 SO 4 c. Cr 2 O 7 2- d. CaCO 3 e. NaClF 4 f. HNO 3 21. In each of the following equations, identify which element is being oxidized, which is being reduced, which is the oxidizing agent, and which is the reducing agent: a. H 3 AsO 4 + Zn  AsH 3 + Zn 2+ b. MnO 2 + Hg + Cl -  Mn 2+ + Hg 2 Cl 2 c. HXeO 4 - + Pb  Xe + HPbO 2 - d. ClO 4 - + I -  ClO 3 - + IO 3 - Chapter 6 Questions: 1. A diver at a depth of 100 ft (pressure = 3 atm) exhales a small bubble of air with a volume equal to 100mL. What will the volume of the bubble be at the surface of the water (1 atm)? 2. What would the volume of gas contained in an expandable 1.0L cylinder at 15 MPa be at 1.0atm. (Assume constant temperature) 3. The volume of a gas (at constant pressure) is to be used as a “thermometer”. If the volume at 0.0 o C is 75.0cm 3 , what is the temperature when the measured volume is 56.7cm 3 ? 4. An 11.2 L sample of gas is determined to contain 0.50 moles of N 2 . At the same temperature and pressure, how many moles of gas would there be in a 20.L sample? 5. What is the volume of 16g of sulfur dioxide at 20 o C and 740 torr pressure?

6. A sample of gas occupies 3.8L at 15 o C and 1.00 atm. What does the temperature need to be for the gas to occupy 8.3L at 1.00 atom? 7. A flask which can withstand an internal pressure of 2500 torr, but no more, is filled with a gas at 21.0 o C and 758 torr and heated. At what temperature will it burst? 8. The density of liquid nitrogen is 0.808 g/mL at -196 o C. What volume of nitrogen gas at STP must be liquefied to make 10.0L of liquid nitrogen? 9. Calculate the volume occupied by 2.5 mol of an ideal gas at STP. 10. An unknown gas has a density of 7.06 g/L at a pressure of 1.50 atm and 280K. Calculate the molar mass of the gas. 11. A 27.7mL sample of CO 2 (g) was collected over water at 25.0 o C and 1.00 atm. The vapor pressure of water at 25.0 o C = 23.8 torr. a. What is the partial pressure (in torr) of the CO 2 ? b. How many moles of CO 2 gas are collected? 12. A gaseous mixture of oxygen, hydrogen, and nitrogen has a total pressure of 1.50 atom and contains 8.20g of each gas. Find the partial pressure of each gas in this mixture. 13. Assume the mole fraction of nitrogen in air is 0.8902. Calculate the partial pressure of nitrogen in the air when the atmospheric pressure is 820 torr. 14. Which gas would effuse faster, Ne or CO 2 ? How much faster? Chapter 4 Answers: 1. a. 2 AgI + 1 Na 2 S  1 Ag 2 S + 2 NaI b. 1 Mg(OH) 2 + 2 HCl  1 MgCl 2 + 2 H 2 O c. 1 TiCl 4 + 2 H 2 O  1 TiO 2 + 4 HCl d. 3 HNO 2  1 HNO 3 + 2 NO + 1 H 2 O 2. 354.3g of PbI 2 3. a. AgNO 3 is limiting, KBr is in excess b. Theoretical yield = 71.64g AgBr c. 20.62% Chapter 5 Answers: 1. K 3 PO 4 (s)  3K + (aq) + PO 4 3- (aq) 2. a. yes b. no c. yes 3. a. nonelectrolyte b. strong electrolyte c. weak electrolyte 4. 1.01 M 5. 0.75 M Na + 6. 35.5g of Na 2 SO 4 dissolved in enough water to make 250.mL 7. 28mL of 17.8M H 2 SO 4 diluted to 500.mL 8. 6.000M NaOH 9. 3.6mL Ca(OH) 2 10. a. 2NaCl(aq) + Hg 2 (NO 3 ) 2 (aq)  Hg 2 Cl 2 (s) + 2NaNO 3 (aq) b. Ca(OH) 2 (aq) + Na 2 CO 3 (aq)  CaCO 3 (s) + 2NaOH(aq) c. 3Na 2 S(aq) + 2FeCl 3 (aq)  Fe 2 S 3 (s) + 6NaCl(aq) 11. SO 4 2- 12. molecular: Na 2 S(aq) + Cu(NO 3 ) 2 (aq)  CuS(s) + 2NaNO 3 (aq) ionic: 2Na + (aq) + S 2- (aq) + Cu 2+ (aq) + 2NO 3 - (aq)  CuS(s) + 2Na + (aq) + 2NO 3 - (aq) net ionic: S 2- (aq) + Cu 2+ (aq)  CuS(s) 13. 2.5g Mg(OH) 2

14. 7.1g CaSO 4 15. 20.2mL HCl 16. 204.1 g/mol 17. a. molecular: H 2 SO 4 (aq) + 2NaOH(aq)  2H 2 O(l) + Na 2 SO 4 (aq) ionic: 2H + (aq) + SO 4 2- (aq) + 2Na + (aq) + 2OH - (aq)  2H 2 O(l) + 2Na + (aq) + SO 4 2- (aq) net ionic: H + (aq) + OH - (aq)  H 2 O(l) b. molecular: HClO 4 (aq) + NH 3 (aq)  NH 4 ClO 4 (aq) ionic: H + (aq) + ClO 4 - (aq) + NH 3 (aq)  NH 4 + (aq) + ClO 4 - net ionic: H + (aq) + NH 3 (aq)  NH 4 + (aq) c. molecular: HC 2 H 3 O 2 (aq) + LiOH(aq)  H 2 O(l) + LiC 2 H 3 O 2 (aq) ionic: HC 2 H 3 O 2 (aq) + Li + (aq) + OH - (aq)  H 2 O(l) + Li + (aq) + C 2 H 3 O 2 - (aq) net ionic: HC 2 H 3 O 2 (aq) + OH - (aq)  H 2 O(l) + C 2 H 3 O 2 - (aq) 18. 50. mL NaOH 19. 0.418M acid 20. a. Mg = +2 Br = -1 b. Na = +1 S = +6 O = -2 c. Cr = +6 O = -2 d. Ca = +2 C = +4 O = -2 e. Na = +1 Cl = +3 F = -1 f. H = +1 N = +5 o = -2 21. a. As changes from +5 to -3: As is being reduced and is the oxidizing agent Zn changes form 0 to +2: Zn is being oxidized and is the reducing agent b. Mn changes from +4 to +2: Mn is being reduced and is the oxidizing agent Hg changes from 0 to +1: Hg is being oxidized and is the reducing agent c. Xe changes from +6 to 0: Xe is being reduced and is the oxidizing agent Pb changes from 0 to +2: Pb is being oxidized and is the reducing agent d. Cl changes from +7 to +5: Cl is being reduced and is the oxidizing agent I changes from -1 to +5: I is being oxidized and is the reducing agent Chapter 6 Answers: 1. 300.mL 2. 150L 3. 206K, -67 o C 4. 0.89 moles 5. 6.2L 6. 629K, 356 o C 7. 970.15K, 697 o C 8. 6460L 9. 56L 10. 108 g/mol 11. a. 736 torr b. 0.001096 moles 12. O 2 = 0.0831 atmH 2 = 1.32 atm N 2 = 0.0952 atm 13. 730 torr 14. Ne would diffuse 1.48 times faster than CO 2

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Study Guide for Chapters 4-6

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