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Advanced Placement Chemistry 1st Semester Review 1. Answer the following questions for each of the listed elements. a. Charge of each possible ion that forms in aqueous solution. b. Formula for any possible polyatomic ion that forms, if any at all. c. Formula of common oxides that form. Identify which oxides yield acidic solution in water and which yield basic solution in water. Element Ion(s) Ion Formula(s) Oxide formula Oxide Rxn w/ Acid Oxide Rxn w/ Base (A) F (B) S (C) Mg (D) Ar (E) Mn 2. 3. List common uses for each of the following substances. Also, list common properties (oxidizer, reducer, amphoteric properties, acidic, basic, neutral) for each. (A) Hydrofluoric acid _____________________________________________________________ (B) Carbon dioxide _____________________________________________________________ (C) Aluminum hydroxide _____________________________________________________________ (D) Ammonia _____________________________________________________________ (E) Hydrogen peroxide _____________________________________________________________ Define the following terms. Give examples of each type of solid. (A) A network solid ______________________________________________ (B) A molecular solid with zero dipole moment ______________________________________________ 4. (C) A molecular solid with hydrogen bonding ______________________________________________ (D) An ionic solid ______________________________________________ (E) A metallic solid ______________________________________________ Assume that you have an "unknown" consisting of an aqueous solution of a salt that contains one of the ions listed above. a. Provide the color of the ion in solution. b. Identify any characteristic odor when mixed with a base then with an acid. c. Identify any ppt that forms from the reaction with the SO42- ion and with the OH1- ion. Formula Name Color of ion (aq) Odor PPT (color) w/ PPT (color) SO42w/ OH12 (A) CO3 ¯ (B) Cr2O72¯ (C) NH4+ (D) Ba2+ (E) Al3+ ©T. M. Chipi, chemistry instructor 1 5. The liquefied hydrogen halides have the normal boiling points given below. Explain the relatively high boiling point of HF as compared to the other halogens. Hydrogen Halide Normal Boiling Point, °C HF +19 HCl - 85 HBr - 67 HI - 35 6. Provide examples below representing a pair of isotopes. Atomic Number Mass Number Element Name I. (A) II. I. (B) II. I. (C) II. I. (D) II. 7. What is the effect observed on each of the following when a gas in a closed container is heated at constant volume until its absolute temperature is doubled? (A) The density of the gas ______________________________________________ (B) The pressure of the gas ______________________________________________ (C) The average velocity of the gas molecules (D) The number of molecules per cm 3 (E) The potential energy of the molecules ______________________________________________ ______________________________________________ ______________________________________________ 8. Which element is represented by the electronic configuration shown, 1s2 2s2 2p6 3s2 3p3? __________________ What is the formula of the compound that forms when this element reacts with Group 1A, 2A, & 3A elements? 9. The density of an unknown gas is 4.20 grams per liter at 3.00 atmospheres pressure and 127 °C. What is the molecular weight of this gas? (R = 0.0821 liter-atm / mole-K) 10. The critical temperature of a substance is the _____________________________________________________ 11. Write the nuclear equation when Po-214 decays; the emission consists consecutively of an alpha particle, then two beta particles, and finally another alpha particle. __________________________________________________________________________________________ 12. What is the color of each of the following ions in solution? (A) Fe(H2O)42+ ____________________ 2+ (B) Cu(NH3)4 ____________________ (C) Cr2O72____________________ + (D) K ____________________ (E) Br¯ ____________________ (F) NO3¯ ____________________ (G) I¯ ____________________ 2+ H) Co ____________________ ©T. M. Chipi, chemistry instructor 2 13. Write the net ionic equation for the reaction between calcium carbonate and hydrochloric acid. _________________________________________________________________________________________ 14. A(s) + B(g) <==> C(s) + D(g); ΔH = (-) How can the substances in the equation above at equilibrium, at pressure P and temperature T, be shifted to favor the products? _________________________________________________________________________________________ 15. What is needed in order to convert from molarity to molality? ____________________________________ 16. What process of radioactive decay changes C-14 to N-14? _______________________________________ 17. Equal masses of (3) different ideal gases, X, Y, and Z, are mixed in a sealed rigid container at constant temperature. What is true about the partial pressure of gas X? How can it be calculated? _________________________________________________________________________________________ 18. Draw the geometry of the SO3 molecule. 19. Give the type of covalent bond found in each of the following molecules and rank them from the shortest to the longest bond length, and from the strongest to the weakest bond. Element Covalent Bond (S,D,T) Bond Length (1-5) Bond Strength (1-5) (A) N2 (B) O2 (C) Cl2 (D) Br2 (E) I2 20. Metallic copper is heated strongly with concentrated sulfuric acid. Write the reaction. _________________________________________________________________________________________ 21. What is the observed trend for atomic radius across the periods for the transition metals? _________________________________________________________________________________________ 22. Write the reaction producing tetraphosphorus decoxide from the burning of white phosphorus. How many moles of O2 is needed to produce 17.0 grams of P4O10? _____________________________________________ 23. What is the general formula of an alkene? If 0.3.00 gram of any alkene is burned in excess oxygen, what number of moles of H2O is formed? ____________________________________________________________ CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l); ΔΗ° = - 889.1 kJ ΔHf° H2O(l) = - 285.8 kJ / mole ΔHf° CO2(g) = - 393.3 kJ / mole 24. What is the standard heat of formation of methane, ΔHf° CH4(g), as calculated from the data above? 25. Two flexible containers for gases are at the same temperature and pressure. One holds 0.50 gram of hydrogen and the other holds 8.0 grams of oxygen. What can be said regarding each of the following for each gas? (A) The volume of the containers. ________________________________ (B) The number of molecules in each container. ________________________________ (C) The density of each sample of gas. ________________________________ (D) The average kinetic energy of the molecules of each gas. ________________________________ (E) The average speed of the gas molecules for each gas. ________________________________ ©T. M. Chipi, chemistry instructor 3 26. Where does Pi bonding occur in a species? ______________________________________________________ 3 Ag(s) + 4 HNO3 <==> 3 AgNO3 + NO(g) + 2 H2O 27. The reaction of silver metal and dilute nitric acid proceeds according to the equation above. If 0.30 mole of powdered silver is added to 30.mL of 3.0-molar nitric acid, how many moles of NO gas can be formed? 28. What is always true about the phase diagram of any one-component system with respect to a. the slope of the curve representing equilibrium between phases ______________________________ b. the temperature at the triple point vs. that at the normal freezing point ______________________________ c. the pressure at the triple point compared to normal pressure? ______________________________ 29. At 20. °C, the vapor pressure of toluene is 22 millimeters of mercury and that of benzene is 75 millimeters of mercury. An ideal solution, equimolar in toluene and benzene, is prepared. At 20. °C, what is the mole fraction of benzene in the vapor in equilibrium with this solution? 30. Ice is added to some hot water in a Thermos-bottle, which is then sealed. What has happened to the total energy and the total entropy when the system reaches equilibrium? 31. Write the ground state electron configuration for the Mn3+ ion. _______________________________________ 32. When 70.mL of 3.0-molar Na2CO3 is added to 30.mL of 1.0-molar NaHCO3 the resulting [Na+] is 33. State whether each of the following has a dipole moment or not (a zero dipole moment). (A) HCN (B) NH3 (C) SO2 (D) NO2 (E) PF5 _________________ _________________ _________________ _________________ _________________ 34. A solution of potassium dichromate is added to an acidified solution of iron (II) sulfate. Write the redox reaction. _________________________________________________________________________________________ 35. What is the proper procedure used in the titration of an acid with a base and vise-versa? _________________________________________________________________________________________ 36. Write the net ionic equation for the reaction that occurs during the titration of nitrous acid with sodium hydroxide ________________________________________________________________________________ 37. Determine which of the following species functions as an oxidizing agent and which as a reducing agent. (A) Cr2O72¯ __________________________ (B) MnO4¯ __________________________ (C) NO3¯ __________________________ (D) S __________________________ (E) I¯ __________________________ 38. Determine which of the following elements are paramagnetic and which are diamagnetic. a. Ca __________________________ b. V __________________________ c. Co __________________________ d. Zn __________________________ e. As __________________________ ©T. M. Chipi, chemistry instructor 4 39. What is the proper procedure in the preparing of a molar solution of a soluble solid in water? _________________________________________________________________________________________ 40. A 20.0-milliliter sample of 0.200-molar K2CO3 solution is added to 30.0 milliliters of 0.400-molar Ba(NO3)2 solution. Barium carbonate precipitates. The concentration of barium ion, Ba2+, in solution after reaction is 41. What is the mole fraction of ethanol, C2H5OH, in an aqueous solution in which the [C2H5OH] = 4.6 molal? 42. Use a set of four quantum numbers to describe one of the outermost electrons in a barium atom in the ground state. ___________________________________________ 43. What is needed to calculate the molecular weight of a gas collected by displacing water in a water-filled flask inverted in a trough of water, but does NOT need to be measured during the experiment? _________________________________________________________________________________________ 44. A 27.0-gram sample of an unknown hydrocarbon was burned in excess oxygen to form 88.0 grams of carbon dioxide and 27.0 grams of water. What is a possible molecular formula of the hydrocarbon? 45. How many moles of NaF must be dissolved in 1.00 liter of a saturated solution of PbF2 at 25 °C to reduce the [Pb2+] to 1 x 10¯6 molar? (Ksp of PbF2 at 25 °C = 4.0 x 10¯8) 46. Consider the equilibrium: HgO(s) + 4 I¯ + H2O <==> HgI42¯ + 2 OH¯; ΔH < 0. What is the effect of each of the following changes on the concentration of HgI42¯? (A) Increasing the concentration of OH¯ ______________________________________________ (B) Adding 6 M HNO3 ______________________________________________ (C) Increasing the mass of HgO present ______________________________________________ (D) Increasing the temperature ______________________________________________ (E) Adding a catalyst ______________________________________________ 47. In a titration experiment based on the equation 5 Fe2+ + MnO4¯ + 8 H+ <==> 5 Fe3+ + Mn2+ + 4 H2O, 25.0mL of an acidified Fe2+ solution requires 14.0 milliliters of standard 0.050-molar MnO4¯ solution to reach the equivalence point. The concentration of Fe2+ in the original solution is 48. Draw structures for each of the following, including resonance structures necessary to describe the bonding satisfactorily. (A) H2S (B) SO2 (C) CO2 (D) OF2 (E) PF3 49. Organize the following aqueous solutions from the highest boiling point to the lowest boiling point. (A) 0.10 M potassium sulfate (B) 0.10 M hydrochloric acid (C) 0.10 M ammonium nitrate (D) 0.10 M magnesium sulfate (E) 0.20 M sucrose ________________ ________________ ________________ ________________ ________________ 50. Calculate the volume of hydrogen gas produced at standard temperature and pressure when 2.00 grams of aluminum metal is added to an excess of hydrochloric acid. ©T. M. Chipi, chemistry instructor 5 Review of AP Chemistry topics: 2nd year course in chemistry, first semester Equilibrium: Equilibrium tables Kc, Kp, ksp, keq Reaction Quotient (Qc, Qp, Qip) Le Chatelier’s principle Solutions: Molarity, molality, % by mass, % by volume, mass/volume % Electrolytes (strong and weak) and non-electrolytes Mole fraction Colligative properties (vapor pressure, osmotic pressure, freezing point depression and boiling point elevation) van’t Hoff factor Ideal and non-ideal solutions Solubility, solubility of gases vs. solubility of solids, solubility curves Henry’s Law (solubility vs. pressure); S1P1 = S2P2 Thermochemistry: Enthalpy Heat of formation Entropy (definition): solid < liquid < gas First Law of Thermo Review of Chemistry 1 topics: 1st year course in chemistry Atomic structure Ionization energy 3- -1 +1 3+ Isoelectronic series [N -F , Ne, Na -Al ] Electron affinity Electron configuration (s, p, d, f) Electronegativity Orbital filling Ionic and atomic radii Paramagnetism Mole concept Chemicals and their use Dimensional analysis Bonding: ionic & covalent, σ & π Empirical formula Dipole moment Colors of ppt (solids) Hydrogen bonding (H-F, H-N, H-O) Bond length/strength Types of crystalline solids Net-ionic equations Molecular geometry Neutralization based on % composition stoichiometric coefficients Kinetic Molecular Theory (gases) Conservation of mass Density-Molar Mass-ideal gas equation Solubility Names & formulas Chemical reactions (redox, etc.) Isotopes Colors of ions in solution Charge of ion in complex ion Naming/formula complex ion Periodicity Normal boiling point Temperature = Average KE of particles Density from PV = nRT Quantum numbers Resonance structures Stoichiometry and general problem solving Organic compounds Phase diagrams: critical point, triple point, critical temperature, phase change Behavior of gases: temperature, pressure, volume, amount or concentration Nuclear Chemistry (alpha, beta, positron, e- capture) Lab procedures (titration, making molar soln, collection of gas over water) ©T. M. Chipi, chemistry instructor 6