1. Determine the mass of water that can be produced when 10.0g of hydrogen is combined with excess oxygen. 2 H 2 + O 2 2 H 2 O

Transcription

1 Pre-AP Chemistry Spring 2016 Final Review Objective 6.1: Students will recognize indicators of chemical change write balanced chemical equations to describe them based on common reactivity patterns. [S.12.C.1, S.14.C.1, S.15.C.6] 1. Balance and give the reaction type for each of the reactions below. (a) _ K (s) + _ H 2 O (l) _ KOH (aq) + _ H 2 (g) (b) _ AlCl 3 (s) _ Al (s) + _ Cl 2 (g) (c) _ C 5 H 12 (g) + _ O 2 (g) _ CO 2 (g) + _ H 2 O (g) 2. Complete and balance the following reactions. (a) Sodium carbonate is heated and decomposes. (b) Hexane is burned in excess oxygen. (c) Iron(III) chloride solution is poured over a piece of silver. (d) Lead(II) nitrate and sodium iodide solutions are mixed. (e) Nitrogen gas is passed over a hot piece of lithium metal. Objective 7.1 Students will use Avogadro s number and the concept of the mole to convert chemical quantities, including the determination of empirical and molecular formulas. [S.12.C.2, S.13.C.1, S.13.C.2, GL.16.C.2] 1. Answer the following questions concerning moles and mass: (a) How many grams are in mol of MgCl 2? (b) How many grams are in molecules of (NH 4 ) 2 SO 4? (c) How many grams of lithium are there in 4.5mol of Li 3 P? 2. Answer the following questions concerning moles and volume at STP: (a) Find the number of nitrogen molecules in 45.2L of nitrogen gas. (b) What volume will 65.3g of helium occupy? (c) Determine the number of oxygen atoms in 450L of air if O2 comprises approximately 21% of normal atmospheric air. 3. Fumaric acid, C 4 H 4 O 4, is a molecule found in most organisms. The percent composition of C, H and O in fumaric acid by mass is? (b) Write the empirical formula for fumaric acid. (c) If I have 13.5g of fumaric acid, then what mass of carbon do I have? 4. A compound was found to be comprised of % iron, % sulfur and % oxygen by mass. What is its empirical formula? Objective 7.2: Students will use stoichiometry to determine quantities formed and consumed in chemical reactions including limiting reactants, excess reactants, theoretical yields, and percent yields. [S.12.C.3, S.12.C.4, S.15.C.1, S.15.C.3, S.15.C.4] 1. Determine the mass of water that can be produced when 10.0g of hydrogen is combined with excess oxygen. 2 H 2 + O 2 2 H 2 O 2. What volume of carbon dioxide can be produced when excess vinegar (acetic acid, HC 2 H 3 O 2 ) is reacted with 45.3g of calcium hydrogen carbonate? 2 HC 2 H 3 O 2 + Ca(HCO 3 ) 2 2 Na + +2 C 2 H 3 O H 2 O +2 CO 2 3. Answer the following questions concerning what would happen if I mixed 10.0g of lithium metal with 30.0L of nitrogen gas: (a) write the BCE; (b) determine the theoretical yield of the product; (c) determine the theoretical yield if 2.0g of product is recovered; and (d) explain whether or not his can be accounted for if I dropped some of the lithium metal as I was preparing to introduce the nitrogen gas. N Li 2 Li 3

2 Objective 8.1: Students will understand the Kinetic Molecular Theory (KMT) for gases and use it to explain the following relationships: Boyle s Law, Charles Law, Gay-Lussac s Law, Avogadro s Law, and Grahams Law. [GL.16.C.1, GL.17.C.1] 1. Explain how molecules and/or atoms in the gaseous phase can exert pressure. 2. Consider a balloon with a volume of 3235 ml and a temperature of 22 C. (a) What would the volume of the balloon become if the temperature were increased to 72 C? (b) Whose law is in operation in this scenario? (c) Explain how this happens using KMT. (d) Why is it necessary to use kelvin temperatures instead of Celsius ones for these types of problems? 3. If argon gas is known to effuse from a container at a rate of molecules per minute, then what would be the rate of effusion of boron trifluoirde gas from the same container under the same conditions? State the gas law used and explain the reasoning of this phenomena. Objective 8.2: Students will apply the combined gas law, ideal gas law, and Dalton s law to calculate the effects of pressure (P), temperature (T), volume (V), and number of moles (n) on a gas system. [GL.16.C.2, GL.17.C.1] 1. A certain bag of chips has a volume of 400mL on the ground, where the temperature is 18 C and the pressure is measured as 94kPa. What is the volume of the bag on bottom of the ocean, where the temperature is 120 C and the pressure is 65atm? 2. Use the ideal gas law to answer the following questions. (a) What volume would 18.5g of helium gas occupy at a temperature and pressure of 100K and 8.5atm, respectively? (b) What is the density of neon gas in a room where the pressure is 185kPa and it is currently 65 C? (c) If an unknown gas is known to have a density of 3.2g/L at a temperature of 300K and a pressure of 110kPa, then what is the molar mass of the gas? 3. Answer the following concerning the pressures of gas mixtures. (a) I have a container with a total pressure of 177kPa that contains argon, xenon, and chlorine gases. If the pressure of xenon and chlorine are equal at 50kPa, then what is the pressure of argon in the tank? (b) If a certain 10.0L container with a temperature of 20 C is loaded with 1.2mol Cl 2, 2.5mol of Br 2, and some F 2 and the pressure is then recorded as 500kPa, then how many moles of F 2 are in the tank? Objective 8.3: Students will apply the ideal gas law to systems not at standard temperature and pressure (STP) to allow stoichiometric calculations. [GL.18.C.1, S.12.C.3, S.12.C.4] 1. How many grams of chlorine gas can be produced when I decompose 45L of BrCl 3 gas at a temperature of 50 C and a pressure of 80kPa? 2. I place 15g of aluminum in a 25.0L container and then add bromine gas so that the temperature and pressure become 60 C and 210kPa, respectively. What is the maximum mass of product that could be formed? Al + 3 Br 2 AlBr 3 3. Explain, in sufficient detail, how you could determine the molar mass of an unknown elemental diatomic gas by reacting it with a known mass of solid sodium. Be sure to mention all pertinent measurements to be made. Objective 10.2: Students will calculate solution concentrations in various units and perform conversions between them, including the following: molarity (M) and molality (m). [S.15.C.3] 1. Calculate the molarity of a solution that contains 20g of NH 4 Cl in exactly 500mL of solution. (b) How many moles of HNO 3 are present in 25.0mL of a 1.20M solution of nitric acid? (c) How many milliliters of a 2.50M KOH solution is needed to obtain 45.0g of KOH? 2. What would the new volume be of a solution if I take a.35m concentration of HCl in a 50ml beaker and would like to reduce my concentration by half? (M1V1=M2V2)

3 Objective 11.1 Students will understand the process of equilibrium, write equilibrium expressions, and calculate the value of equilibrium constants. [E.24.C.1, E.24.C.2] 1. Write an equilibrium expression for the following reactions: a) 3NO(g) N2O(g) + NO2 (g) b) CH4 (g) + 2 H2S (g) CS2 (g) + 4 H2 (g) c) Ni(CO4) (g) Ni (s) + 4 CO(g) d) HF(aq) H+ (aq) + F- (aq) e) 2 Ag (s) + Zn+2 (aq) 2 Ag+(aq) + Zn (s) 2. Gaseous hydrogen iodide is place in a closed container at 425 C, where it partially decomposes to hydrogen and iodine: 2 HI (g) H2 (g) + I2 (g) At equilibrium it is found that [HI] = 3.53 x 10-3 M, [H2] = 4.79 x 10-4 M, and [I2] = 4.79x 10-4 M. What is the Kc at this temperature? 3. The following equilibrium process is being studied at 230 C. 2 NO (g) + O2 (g) 2 NO2 (g) In one experiment, the concentrations of the reacting species at equilibrium are found to be [NO] = M, [O2] =.127 M, and [NO2] = 15.5 M. Calculate the equilibrium constant of the reaction at this temperature. 4.For the Haber process, N2 (g) + 3 H2 (g) 2 NH3 (g), Kc 1.45 x 10-5 at 500 C. In an equilibrium mixture of the three gases at 500 C, the concentration of H2 is.928 M and N2 is.432 M. What is the concentration of NH3 at this mixture? 5. Phosgene, COCl2, decomposes according to the reaction : COCl2(g) <-> CO(g) + Cl2(g) Kc = 8.3 x 10-4 at 360ºC. Calculate the equilibrium concentrations of reactant & products when 5.00 mol of COCl2 decomposes in a 10.0 L flask. Objective 11.2 Students will use Le Chaltelier s Principle to predict the effect of system changes on the equilibrium position and will use ICE diagrams to solve for equilibrium concentrations. [E.24.C.3, E.24.C.4] 1. Explain Le Chatelier s principle. How can this principle help us maximize the yields of reactions? 2. List four factors that can shift the position of an equilibrium. Only one of these factors can alter the value of the equilibrium constant. Which one is it? 3. Consider 4 NHs (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g), ΔH= kj. How does each of the following changes affect the yield of NO at equilibrium? Answer increase, decease, or no change. a. increase [NH3] b. increase [H2O] d. decrease the volume of the container in which the reaction occurs e. add a catalyst 4. Heating solid sodium bicarbonate in a closed vessel establishes the following equilibrium: 2 NaHCO3 (s) Na2CO3 (s) + H2O (g) + CO2 (g) What would happen to the equilibrium position if a) some of the CO2 were removed from the system; b) some solid Na2CO3 were added to the system ; c) some of the solid NaHCO3 were removed from the system? The temperature remains constant

4 Objective 12:1: Students will understand and apply the 1stlaw of thermodynamics in the context of a chemical system[ke.23.c.5] 1.Define the following: system, surroundings, potential energy, kinetic energy, heat, exothermic, endothermic, 1stlaw of thermodynamics. 2.Where is energy in a chemical system stored? 3.Why is energy released or gained in a chemical reaction? 4.Which way does heat flow when exchanged? 5.List three common forms of energy Objective 12.2:Students will define and apply the concept of enthalpy to calculate heat exchange using constant pressure calorimetry[ke.23.c.1, KE.23.C.4 3.A certain calorimeter is constructed and used to measure the specific heat of a metal. The 1 10g of water in the calorimeter was initially at 20 C. If placing 250g of the metal, initially at 140 C, in the water caused its temperature to rise to 31 C, then what is the specific heat of the metal? 4.The specific heat of copper is J/g C. How many joules of heat are necessary to raise the temperature of a 1.42kg block of copper from 25.0 C to 88.5 C? Objective 11.1: Students will name acids and bases and compare and contrast the following acid-base theories: Arrhenius, Brønsted-Lowry, and Lewis. [AB.19.C.1, AB.20.C.1] 1. Write the formulas for the following acids/bases: a. sodium hydroxide b. hydrochloric acid c. sulfuric acid d. hypochlorous acid 2. Explain why water is an amphoteric substance. 3. Compare and contrast the Arrhenius and Bronsted-Lowry acid/base theories. 4. Determine the most specific acid/base classification for the following: a. H + + OH - H 2 O b. HBr + H 2 O H 3 O + + Br - c. HClO + BrO - HBrO + ClO - Objective 11.2: Students will compare and contrast acid-base properties and perform calculations to quantify the relationships between ph, poh, [H+], and [OH ] as they relate to acids and bases. [AB.21.C.1, AB.21.C.3] 1. Compare and contrast properties of acids and bases. 2. Determine the ph of 0.2M HCl. 3. Determine the ph of a solution with [OH ] = M 4. A solution is produced by dissolving 35.2g LiOH into 2000mL of solution. Determine the ph, [H + ], [OH - ], and poh. Objective 11.3: Students will use titration as an analysis method for determining the concentration of unknown solutions via solution stoichiometry. [AB.22.C.1, AB.22.C.2, S.15.C.4] 1. Explain the difference between an endpoint and an equivalence point in a titration.

5 2. Complete and balance the following equation: Ba(OH) 2 (aq) + HCl (aq) + 3. You load a buret with 0.5 M HCl and record the initial volume as 6.70 ml. You prepare a 10.0 ml sample of barium hydroxide by adding an additional 25 ml of water and a few drops of phenolphthalein. You titrate to the phenolphthalein end point. The final reading of the buret was 22.2 ml. a. How much acid was used? b. Find the molarity of the barium hydroxide solution.