2017 SUMMER ASSIGNMENT AP CHEMISTRY

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Berenice Goodwin
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2017 SUMMER ASSIGNMENT AP CHEMISTRY - This summer work packet is essentially a review of all material covered during Honor Chemistry. It must be completed prior to the first day of class.

1 2017 SUMMER ASSIGNMENT AP CHEMISTRY - This summer work packet is essentially a review of all material covered during Honor Chemistry. It must be completed prior to the first day of class. The problems presented in this packet consist of free response type problems similar to those found on the AP Chemistry Examination. You will also find a solution and grading rubric after each question. After you complete the problem, you must grade yourself using the approved solution and the grading rubric. If you achieve a score less than a five on any one problem, you should rework that problem until you achieve a score greater than a five. Use your own paper for each problem. You are not authorized to work with any other student on these problems. However, you may review the key concepts associated with each problem with another student before attempting the problem. You may use your textbook, notes, and calculator in completing the problems. You can expect that AP Chemistry Examination questions will be in essentially the same format and will be graded in much the same way as the rubrics in this packet. During the examination, you will have probably no more than 20 minutes for each problem. You will only have access to an examination specific reference sheet which includes a periodic table and various formulas. In addition to these ten problems you must memorize the atomic symbol, name, atomic number, and group number for the first 54 elements on the periodic table, all of the polyatomic ions listed on your Honors Chemistry reference data sheet, and the colors of ions in solution. You will be quizzed on these on the first day of class. If you discover that you are struggling significantly with grasping the concepts reflected in these problems, have trouble achieving a score of at least a five, take over an hour to work any one of these problems, or are unwilling to memorize the listed information, you may want to reconsider if AP Chemistry is appropriate for you. Approved: K. Miller Updated:

2 Objectives In completing the problems in this packet, students will: 1. Translate among macroscopic observations of change, chemical equations, and particle views. (LO.3.1, SP.1.5 and 7.1) 2. Translate an observed chemical change into a balanced chemical equation and justify the choice of equation type (molecular, ionic, or net ionic) in terms of utility for the given circumstances. (LO.3.2, SP.1.5 and 7.1) 3. Use stoichiometric calculations to predict the results of performing a reaction in the laboratory and/or to analyze deviations from the expected results. (LO.3.3, SP.2.2 and 5.1) 4. Relate quantities to identify stoichiometric relationships for a reaction including situations involving limiting reactants and situations in which the reaction has not gone to completion. (LO.3.4, SP.2.2, SP.5.1, and SP.6.4) 5. Design and/or interpret data from an experiment that uses gravimetric analysis to determine the concentration of an analyte in a solution. (LO.1.19, SP.4.2, SP.5.1, SP.6.4) 6. Design a plan and use data collected on the synthesis or decomposition of a compound to confirm the conservation of matter and the law of definite proportions. (LO.3.5, SP.2.1, SP.4.2, and SP.6.4) 7. Justify the observation that the ratio of the masses of the constituent elements in any purse sample of that compound is always identical on the basis of the atomic molecular theory. (LO.1.1, SP.6.1) 8. Select and apply mathematical routines to mass data to identify or infer the composition of pure substances and/or mixtures. (LO.1.2, SP.2.2) 9. Select and apply mathematical routines to mass data in order to justify a claim regarding the identity and/or estimated purity of a substance. (LO.1.3, SP.2.2, and SP.6.1) The alpha-numeric codes identified with each objective are taken from the objectives and science practices required by the College Board for an AP Chemistry course.

3 Problem 1 You are given three unknown hydrocarbon compounds (unknowns A, B, and C), each containing only carbon and hydrogen. Using the data provided determine the identities of the unknowns. 1. Complete combustion of 1.00 grams of unknown A resulted in the formation of 1.39 liters of CO2 (at STP) and 2.23 grams of H2O. What is the most probable molecular formula of unknown A? 2. Determine the molecular weight/mass of unknown B if the vapor density of B at STP is 3.13 grams per liter. 3. The empirical formula for unknown B is CH2. Using the results from question 2, determine the molecular formula. 4. You have 1.00 grams of unknown C which contains grams carbon. The complete combustion of 1.00 mole of unknown C requires 5 moles of O2. What is the most likely molecular formula of unknown C?

4 Problem 2 The following questions refer to a laboratory activity designed to determine the concentration of lead ions in a solution of lead (II) nitrate using gravimetric analysis. In the experiment, a 50.0mL sample of Pb(NO3)2 solution was added to an excess of potassium iodide solution. The resulting precipitate, PbI2, was collected via decanting and vacuum filtration, dried in a drying oven, and weighed. You may assume that: - Both solutions were pure - No decomposition took place in the drying oven - The filter did not allow any PbI2 to pass through - All lead is lead (II) - An analytical balance was used to weigh the product 1. If g PbI2 were collected, how many grams of lead (II) nitrate must have been in the 50.0 ml solution? 2. What was the molarity of the lead (II) nitrate solution? 3. How many grams of lead (II) ions were in the lead (II) nitrate solution? 4. What is the minimum number of moles of potassium iodide that must be present in solution to assure a complete precipitation of all lead from solution? 5. If the solid PbI2 on the filter paper was not thoroughly rinsed with distilled water prior to drying, what implications would that have on your answer to the first question? Defend/explain your response.

5 Problem 3 Write the formulas to show the reactants and the products for the following three laboratory situations. In all cases a reaction occurs. Assume that solutions are aqueous unless otherwise indicated. Represent substances in solution as ions if the substances are extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction, in other words, prepare a net ionic equation for reactions involving ionic compounds. 1. Calcium oxide powder is added to distilled water. 2. Methylamine gas is bubbled into distilled water. 3. A 0.2M barium nitrate solution is added to an alkaline 0.2M potassium chromate solution. Problem 4 NH3 (aq) + H2O (l) NH4 + (aq) + OH - (aq) In an aqueous solution, ammonia reacts as represented about. In M NH3 (aq) at 25 o C, the hydroxide ion concentration, [OH - ] is 5.60 X 10-4 M. In answering the following, assume that temperature is constant at 25 o C and that volumes are additive. 1. Write the equilibrium-constant expression for the reaction represented above. 2. Determine the ph of M NH3 (aq). 3. Determine the value of the base ionization constant, Kb, for NH3 (aq). 4. In an experiment, a 20.0 ml sample of M NH3 (aq) was placed in a flask and titrated to the equivalence point and beyond using M HCl (aq). Determine the volume of M HCl (aq) that was added to reach the equivalence point.

6 Problem 5 A titration experiment was conducted to determine the ph of a known volume of a strong monoprotic acid solution. A hydroxide solution of known concentration was poured into a buret and then titrated into the acid solution. The volume of hydroxide solution titrated into the acid was measured at the equivalence point and used to calculate the concentration of the acid solution. Which of the following would cause an error in the calculated value of the ph of the acid solution? Explain each answer. 1. The buret was rinsed with the hydroxide solution before the solution was added to the buret. 2. The experimenter did not notice that a few drops of hydroxide solution spattered outside the acid solution container during the titration. 3. The buret was rinsed with distilled water before the hydroxide solution was added. 4. The experimenter read the hydroxide solution level from the top of the fluid instead of the bottom of the meniscus both before and at the equivalence point. 5. Some hydroxide solution was spilled while the experimenter was pouring it into the buret. Problem 6 How many grams of NaCl are required to precipitate most of the Ag + ions from a 2.50X10 2 ml of M AgNO3 solution? The only source of Cl - ions available to combine with the Ag + ions to form the precipitate is the NaCl. 1. Write the net ionic equation for the reaction. 2. Calculate the percent by mass of Ag+ and Cl- ions in AgCl. 3. Calculate the number of grams of Ag+ ions in the solution of AgNO3. 4. Determine the number of grams of NaCl. Problem 7 The following three mixtures have been prepared: 1. CaO plus water: 2. SiO2 in water 3. CO2 plus water. For each mixture, predict whether the ph is less than 7, equal to 7, or greater than 7. Justify your answers.

7 Problem 8 A 0.562g sample of an unknown substance was dissolved in 17.4g benzene. The freezing point of the solution was o C. The freezing point of pure benzene is o C. For benzene, Kf = o C/m and Kb = 2.61 o C/m. Assume that the solute is a non-electrolyte. 1. What is the molality of the solution? 2. What is the molar mass of the unknown? 3. If the boiling point of pure benzene is 80.2 o C, what is the boiling temperature of the solution? Problem 9 A 15.0 gram sample of solid sodium carbonate reacts with 23.1 grams sulfuric acid to form an aqueous solution of an ionic salt, carbon dioxide gas, and liquid water. When the reaction is complete and the liquids evaporated, 16.9g of the ionic salt is recovered. 1. Determine the identity of the ionic salt. 2. Identify the limiting reactant. 3. Determine the percent yield. Problem 10 Use your understanding of atomic molecular theory to prove that a gram sample of pure lead (II) nitrate has the same ratio of the masses of the constituent elements as a gram sample that contains 94% lead (II) nitrate and 6% other compounds.

8 Approved Solution and Grading Rubric for Problem 1 The most important thing to recognize in this problem is that you are dealing with a hydrocarbon consisting of only carbon and hydrogen. This means that the only source of carbon in the CO2 and hydrogen in the H2O must be from the hydrocarbon. Question 1: Determine the number of moles of carbon in Unknown A: Using either the volume of CO2 or the amount of H2O, determine the number of moles of carbon using stoichiometry and dimensional analysis. Assume one mole of the hydrocarbon is used L CO2 1 mol CO2 1 mol C 22.4 L CO2 1 mol CO2 = mol C If you started with H2O, you should find that you get the same number of moles. Determine the number of moles of hydrogen in Unknown A: Since the only source of hydrogen is the hydrocarbon, use the H2O to determine the number of moles of hydrogen g H2O 1 mole H2O g H2O = mol H2O Since the mole ratio of hydrogen to oxygen in H2O is 2:1, you must have mol H. Determine the mole ratio between carbon and hydrogen in the unknown..248 mol H/.0621 mol C = 4 In other words, there is a 4:1 ratio of hydrogen atoms to carbon atoms in the unknown. The empirical formula for Unknown C is, therefore, CH4. You do not know the molecular mass of Unknown A, but you do know the amounts analyzed, so you can determine how many moles of a product would be produced by 1.00g CH g CH4 1 mol CH4 1 mol CO g CH4 1 mol CH4 = mol CO2, since this matches the number of moles determined earlier, the multiplier must have been 1, so the molecular formula is the same as the empirical formula, CH4.

9 Question 2: Use the ideal gas law or molar volume to determine the molecular weight/mass of Unknown B g Unk B 22.4 L Unk B 1 L Unk B 1 molunk B = 70.1 g/molunk B Question 3: Determine the multiplier for the empirical formula g Unk B (molecular) 1 molunk B 1 molunk B 14.0 g Unk B (empirical) = 5, therefore the molecular formula is C5H10. Question 4: 1.00 g of Unknown C contains g carbon or 81.7%. It would also have 18.3% hydrogen. Assuming a 100 g sample: 81.7 g C 1 mol C g C = 6.80 mol C 6.80 = 1 mol C X 3 = 3 mol C 18.3 g H 1 mol H 1.01 g H = mol H 6.80 = 2.66 mol H X 3 = 8 mol H The empirical formula must be C3H8. Since you don t know the molecular weight/mass of the unknown, determine the mole ratio from the balanced chemical equation: C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (g) The mole ratio between C3H8 and O2 is 1:5; notice that if you multiplied the subscripts in the empirical formula by 2, the mole ration would increase to 1:10 and this would continue increasing. This means that the empirical and molecular formulas are the same: C3H8.

10 Approved Solution and Grading Rubric for Problem 2 Before answering any of the questions, you must determine the balanced chemical equation for the reaction: Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq) Question 1: g PbI2 1 mol PbI2 1 molpb(no3) g Pb(NO3) g PbI2 1 mol PbI2 1 molpb(no3)2 = g Pb(NO3)2 Question 2: 50.0 ml solution =.05 L solution g Pb(NO3)2 1 molpb(no3) g Pb(NO3)2.05L solution =.100 M Pb(NO3)2 Question 3: % mass Pb = mass Pb/mass Pb(NO3)2 = g/331.2 g = or 62.56% g Pb(NO3)2 X.6256 = g Pb Question 4: g Pb(NO3)2 1 molpb(no3)2 2 mol KI g Pb(NO3)2 1 molpb(no3)2 =.010 mol KI or 1.00 X 10-2 mol KI Question 5: Failure to adequate rinse the lead (II) iodide will increase the likelihood of contaminants. Upon drying this will increase the mass of the solid which will, in turn, result in an overestimation of the amount of lead (II) iodide formed. This will affect all calculations depicted above.

11 Approved Solution and Grading Rubric for Problem 3 Each answer earns 3 points, for the reactants and for the products. All products must be correct to earn both product points. Equations do not need to be balanced and phases do not need to be indicated. Any spectator ions on the reactant side nullify the 1 possible reactant point, but if they appear again on the product side, there is no product-point penalty. A fully molecular equation, when it should be ionic, earns a maximum of. Ion charges must be correct (1 additional point if all ion charges in all questions are correct). Question 1: CaO (s) + H2O (l) Ca(OH)2 (aq) 3 points Question 2: CH3NH2 (g) + H2O (l) CH3NH3 + (aq) + OH - (aq) 3 points Question 3: Ba 2+ (aq) + CrO4 2- (aq) BaCrO4 (s) 3 points All ion charges are correct!

12 Approved Solution and Grading Rubric for Problem 4 1. K = [NH4 + ][OH - ]/[NH3] 2. [OH - ] = 5.60 X 10-4, therefore poh = -log [OH - ] = ph + poh = 14 ph = 14 poh = = Kb = [NH4 + ][OH - ]/[NH3] since this is a weak base, you will need to use the ICE equation. NH3 (aq) H2O NH 4+ OH - Initial (M) Change (M) -x +x +x Equilibrium (M) x x x x = [OH - ] = [H + ] = 5.60 X 10-4 M Kb = [NH4 + ][OH - ]/[NH3] = x 2 / x = (5.60 X 10-4 )2/ X 10-4 = 1.80 X NH3 + H + NH 4+ Mol NH3 = mol/L X.0200 L = 3.60 X 10-4 mol = mol H+ ions needed Volume HCl (aq) = mol needed/molarity = 3.60 X 10-4 mol/0.0120m =.0300 L or 30 ml

13 Approved Solution and Grading Rubric for Problem 5 1. This is the proper experimental procedure and will have no adverse impact on the calculated value of the acid solution. 2. This will make the measured volume of the hydroxide solution larger than the actual amount added, the calculated value for the moles of OH - and H + will be too large, the calculated value of [H + ] will be too large, and the calculated ph will be too small. 3. This will dilute the hydroxide solution, which means that too large a volume of hydroxide solution will be added, the calculated value for OH - and H + will be too large, the calculated value of [H + ] will be too large, and the calculated ph will be too small. 4. The levels were read consistently although they were read from the wrong spot. The two errors should cancel, and the calculated ph should be correct. 5. This will not affect the concentration of the hydroxide solution or the measurement of the volume poured into the acid solution the acid solution, so the calculated ph should not be affected.

14 Approved Solution and Grading Rubric for Problem 6 1. Ag + (aq) + Cl - (aq) AgCl (s) 2. mass of element/mass of compound X 100% = % mass of element % mass Ag + = g Ag + /143.4 g AgCl =.752 or 75.2% % mass Cl - = 35.45g Cl - /143.4 g AgCl =.248 or 24.8 % 3. molarity solution X volume of solution X molar mass = mass.0113 mol Ag +.25L solution g Ag + 1 L solution 1 mol Ag + =.305 g A + 4. Since the only source of Ag+ ions was the solution of AgNO3, the mass of Ag+ ions in the AgCl must be.305 g Ag+, therefore the mass of the compound (AgCl) can be found using the formula for % mass. % mass element = mass of element/mass of compound mass of compound = mass of element/% mass element mass AgCl =.305 g Ag+/.752 =.405g AgCl This can now be used to find the mass of the Cl - Mass Cl - =.248 X.405g AgCl =.100g Cl - The % mass of Cl - in NaCl can be found using the same basic procedure. % mass Cl - = mass of element/mass of compound = g Cl - /58.44 g NaCl =.607 or 60.7% mass of compound = mass of element/% mass element mass NaCl =.100g Cl - /.607 =.168 g NaCl

15 Approved Solution and Grading Rubric for Problem 7 1. CaO dissolves into a weak base so its ph should be 8-10 (greater than 7). 3 points 2. SiO2 is insoluble so the ph will remain unchanged (equal to 7). 4 points 3. CO2 dissolves into a strong acid so its ph should be 1-3 (less than 7). 3 points

16 Approved Solution and Grading Rubric for Problem 8 1. molality = moles solute/kilograms solvent Since you don t know the moles of the unknown substance you need to use the freezing point depression equation for benzene to find the molality. Tf = Kfm or m = Tf/Kf Tf = o C o C = 1.38 o C = 1.38 o C/5.065 o C/m = m unknown solution 2. Use the molality calculated in previous step and the basic definition of molality in a stoichiometric set up to find the molar mass of the unknown substance..562 g Unk 1 Kg benezene.0174 Kg benzene.272 molunk = 199 g/mol unknown substance 3. Tb = Kbm = 2.61 o C/m (.272 m) =.710 o C 80.2 o C o C = 80.9 o C

17 Approved Solution and Grading Rubric for Problem 9 You must determine the balanced chemical equation before beginning this problem. Na2CO3 (s) + H2SO4 (aq) Na2SO4 (aq) + CO2 (g) + H2O (l) 1. Since this is a double replacement reaction, the ionic salt would have to be Na2SO4. 2. The balanced chemical reaction shows that there is a one-to-one molar ratio of the two reactants g Na2CO3 1 mol Na2CO g Na2CO3 =.14 mol Na2CO g H2SO4 1 mol H2SO g H2SO4 =.24 mol H2SO4 molar ratio =.24 mol H2SO4/.14 mol Na2CO3 = 1.7 molh2so4to 1 mol Na2CO3, therefore there is an excess of H2SO4. 3. Determine the theoretical yield using stoichiometry and compare to actual yield to determine the percent yield g Na2CO3 1 mol Na2CO3 1 mol Na2SO g Na2SO g Na2CO3 1 mol Na2CO3 1 mol Na2SO4 = 20.1 gna2so4 % yield = actual yield/theoretical yield = 16.9 g Na2SO4/20.1 gna2so4 =.841 or 84.1 % yield

18 Approved Solution and Grading Rubric for Problem 10 Atomic molecular theory indicates that atoms of elements combine in small, whole number ratios to form distinct molecules. In this particular case, lead (II) nitrate would have the following ratio: 5 points Pb(NO3)2 Or 1 atom of lead to 2 atoms of nitrogen to 6 atoms of oxygen or 1Pb 2+ ion to 2 NO3 - ions If you were to determine the percent composition of each element (using their molar masses) in the pure sample, you would discover that the ratio of these percentages would simplify to the ratio identified above. Likewise, if you did the same thing using the amount of lead (II) nitrate in the impure sample, you would discover the same ratio. 5 points

Chapter 4 Reactions in Aqueous Solutions. Copyright McGraw-Hill

Chapter 4 Reactions in Aqueous Solutions Copyright McGraw-Hill 2009 1 4.1 General Properties of Aqueous Solutions Solution - a homogeneous mixture Solute: the component that is dissolved Solvent: the component