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Chem 173: Final Exam Review Short Answer and Problems 1. Provide the missing name, elemental symbol, or chemical formula for each of the following: Tc __________________________________ ____________________________ ____________________________ Y(NO3)3 ______________________________ manganese Ag2SO4 ______________________________ ammonium chloride ____________________________ SnI2 __________________________________ CCl4 __________________________________ 2. copper (I) oxide potassium bromate ____________________________ nitrous acid ____________________________ Balance the following chemical equation: ______ IBr (s) + ______ NH3 (g) Ÿ ______ NI3 (s) + ______ NH4Br (s) 3. Write the balanced chemical equation that describes the combustion of acetylene gas, C2H2. 4. Consider the following elements: Cr Be N Co Ar Rb Br In Identify which of these elements correspond to the following descriptions. NOTE: Each of the following descriptions is written in the plural form for consistency only. Some blanks will have only one answer, others may have more than one answer. 5. Those which are alkali metals: ______________________ Those which are non-metals: ______________________ Those which are in the 4th period of the periodic table: ______________________ Those which are noble gases: ______________________ Those which exist as diatomic molecules in their elemental state: ______________________ Those which are transition metals: ______________________ Provide the missing pieces of information in the following table: Symbol (including mass number and ion number of protons number of neutrons charge) number of electrons 69 Ga3+ 34 6. 36 Perform the following unit conversions: (1 mi = 1.6093 km, 1 h = 3600 s) 72 mi/h = __________ m/s 7. 46 3.68 mg/mm3 = ____________ µg/m3 10.449 dm = ___________ nm Give the answer to the following calculation with the proper number of significant figures: [(14.7 – 13.645)/2.04](26.93) = _______________________ 8. Consider the following statement; then circle the physical properties and underline the chemical properties of mercury (II) oxide. Mercury (II) oxide is an orange-red solid with a density of 11.1 g/cm3. It decomposes when heated to give mercury and oxygen. Mercury (II) oxide does not dissolve in water. 9. Consider the following list of compounds, and then identify which of these compounds correspond to the following descriptions. NOTE: All descriptions are written in the plural form. Some blanks will require only one answer, others more than one. If none of the compounds match the description, write NONE. P 4O10 LiOH CH3OH NH3 H2SO4 MgCl2 HC2H3O2 10. those which are strong electrolytes _______________________________ those which are weak electrolytes _______________________________ those which are non-electrolytes _______________________________ those which are strong acids _______________________________ those which are strong bases _______________________________ those which are salts _______________________________ 1.40 mol of yttrium metal and 1.40 mol of molecular oxygen are allowed to react according to the following equation: 4 Y (s) + 3 O2 (g) Ÿ 2 Y2O3 (s) In this reaction the limiting reactant is _______________ . If 0.475 mol of Y2O3 is actually recovered in an experiment, the percent yield is ________________. 11. Predict possible products (include physical states) for the following reactions. If no reaction will occur, write NONE. a. LiOH (aq) + LiHSO4 (aq) Ÿ ___________________________________________________ b. AgNO3 (aq) + CdCl2 (aq) Ÿ ___________________________________________________ c. CO2 (g) + KOH (aq) Ÿ ___________________________________________________ 12. Write the balanced chemical, full ionic, and net ionic equations that describe the reaction that occurs when an aqueous solution of calcium hydroxide reacts with chlorous acid (HClO2). Be sure to indicate the charges and phases throughout. balanced chemical equation: full ionic equation: net ionic equation: 13. Indicate the oxidation number for: a. S in HSO3– __________ 14. b. Na in NaClO2 __________ c. Se in K2Se __________ Write the balanced redox equation for the following reaction that occurs in basic solution. Use the method of balancing the half reactions. Circle the reducing agent. Zn (s) + NO3– (aq) Ÿ NH3 (aq) + Zn(OH)42– (aq) oxidation half reaction: reduction half reaction: overall equation: 15. Match each one of the following species with the appropriate descriptive chemical property: Na+ __________ a. a proton donor BaO __________ b. an oxidizing agent P 2O5 __________ c. an acid anhydride HC2H3O2 __________ d. a spectator in precipitation reactions MnO4– __________ e. a basic oxide 16. Sodium chloride dissolves in water according to the following thermochemical equation: NaCl (s) Ÿ Na+ (aq) + Cl– (aq); 17. ∆H = +3.90 kJ a. Is the dissolution of NaCl an endothermic or exothermic process? _______________________ b. Is energy released or absorbed during this process? _______________________ 2.50 mol of NO2 (g) is placed in a 15.0 L container at 35°C. The average speed of the NO2 molecules in this container is __________________ . If the temperature is changed to 100°C, the average speed of the NO2 molecules will _________________ (increase, decrease or remain constant), and the pressure exerted by the gas will ____________________ (increase, decrease or remain constant). 18. a. For a given compound, DH°vap is the enthalpy change that corresponds to the physical change from the _________________ state to the _________________ state. b. The phase change described above in a is _____________________ (endothermic, exothermic). c. For mercury, DH°fus = 2.292 kJ/mol and DH°vap = 59.30 kJ/mol. What is the value of DH°sub for mercury? ____________________ d. Which of the following will not have DHf ° equal to zero? K (l), F2 (g), Ar (g), O2 (g) ____________________ 19. Write the chemical equation of formation for methanol, CH3OH (l). 20. There are 2 compounds of the formula Pt(NH3)2Cl2. Cl NH3 Cl Pt Cl Cl Pt NH3 NH3 NH3 The compound on the right, cisplatin, is used in cancer therapy. Both compounds have square planar geometry. Circle the compound that has a non-zero dipole moment. 21. Consider the acetonitrile molecule: H H C C N H a. What are the bond angles around each carbon? __________ __________ b. What is the hybridization on each C atom? __________ __________ c. Determine the total number of s and p bonds. s __________ p __________ 22. 23. a. Draw the Lewis structure of XeF2 b. What is the electron pair geometry? _________________________ c. What is the molecular shape? _________________________ d. What orbitals on Xe and F participate in the bonding? _________________________ e. What orbitals on F house nonbonding electron pairs? _________________________ f. Is this molecule polar or nonpolar? _________________________ a. Will potassium nitrate be more soluble in water or hexane (C6H14)? ______________________ b. Will propane (C3H8) be more soluble in water or carbon tetrachloride (CCl4)? ______________________ c. Urea - a nonvolatile, molecular compound - is dissolved in pyridine. Will the vapor pressure of the solution increase, decrease, or remain constant relative to pure pyridine? d. ______________________ If NH4NO3 is dissolved in pyridine, would you expect the ∆Psoln to be larger or smaller than the ∆Psoln observed in the solution prepared in part c above? e. ______________________ A solution is prepared by dissolving 1 mol of ammonium phosphate (NH4)3PO4 in water. What is the theoretical value of i (# mol particles in solution) for this reaction? f. ______________________ Will the vapor pressure of the solution described in a be higher or lower than the vapor pressure of pure water? ______________________ g. Will the vapor pressure of the solution described in e be higher or lower than a solution prepared by dissolving 1 mol of potassium chloride (KCl) in water? 24. 25. 26. ______________________ A solution is prepared by dissolving 30.0 g of propanol (C3H7OH, molar mass = 60.10 g/mol) in 80.0 g of water (molar mass = 18.02 g/mol) resulting in a solution with a total volume of 108 mL. a. Calculate the molarity of this solution. ______________________________ b. Calculate the mole fraction of propanol in this solution. ______________________________ c. Calculate the mass % propanol in this solution. ______________________________ d. Calculate the molality of this solution. ______________________________ a. Draw the Lewis structure (including any resonance structures) for carbonate anion. b. The bonding in this structure is best described as {localized, delocalized, ionic, or nonpolar}. ________________ c. What obritals would be involved in p bonding between C and O? ____________ on C and ____________ on O a. Write the balanced chemical equation for the dissociation of benzoic acid,HC7H5O2. b. What is the conjugate base of benzoic acid? ______________________ c. Write the Ka expression for this dissociation. ______________________ d. If pKa of benzoic acid = 4.19, what is Ka of benzoic acid? ______________________ 27. The reaction 2 A + B Ÿ C obeys the following rate equation: Rate = k[A][B]1/2. The order of the reaction with respect to B is ___________________, and the total order of the reaction is ____________________. The units of k are _____________________. If the rate of formation of C is 2.00 mol•L—1•s—1, then the rate of consumption of A = ____________________. 28. The rate constant (k) for a reaction was studied as a function of temperature (T). The activation energy for this reaction can be found from the slope in a linear plot of _________________ vs. ___________________. 29. Write the expression for the equilibrium constant KC for the following reaction: C (s) + 1/2 O2 (g) ÷ CO (g) 30. Consider the following equilibrium: Ca2+ (aq) + H2O (l) + CO2 (g) ÷ CaCO3 (s) + 2 H+ (aq); DH = +16 kJ/mol Each of the following changes are made separately. Circle the predicted effect of each change on the quantity listed (i = increase, d = decrease, nc = no change). Change Add catalyst Quantity Amount of H+ Effect i d nc Add more CaCO3 Amount of H+ i d nc + Increase Volume Amount of H i d nc Decrease temperature K i d nc Increase temperature Amount of CaCO3 i d nc Add Ar Amount of CO2 i d nc K i d nc 2+ Add more Ca 31. KP = 0.250 at 1100 K for the following equilibrium: 2 SO2 (g) + O2 (g) ÷ 2 SO3 (g) For this reaction K = ________________________, and for the following reaction: SO2 (g) + 1/2 O2 (g) ÷ SO3 (g) KP = __________________________ . 32. Write chemical formulas for three strong acids: ______________ , ______________ , _____________ 33. Identify each of the following species as acid, base, conjugate acid, or conjugate base: H2PO4– (aq) + SO32– (aq) ÷ HPO42– (aq) + HSO3– (aq) ___________ 34. ___________ ___________ ___________ Suppose that QP < KP for the reaction 2A + B ÷ C at some time. As the reaction proceeds, will the partial pressures of the substances change? If so, how? _______________________________________________________________ 35. The pH of a solution at the stoichiometric point in a titration of a ____________________ acid and a ___________________ base is less than 7.00. 36. Write the Ksp expression for Cr2(CO3)3 Ksp = ____________________________ 37. Write the balanced net ionic equation ( include physical states) that describes each of the following: a. the dissolution/precipitation equilibrium for copper (I) sulfate, Cu2SO4: b. the neutralization reaction that occurs when KOH is added to a buffer solution composed of acetic acid and potassium acetate c. the equilibrium that determines the pH of a solution of sodium hypochlorite, NaOCl (aq) 38. The pH of a 0.20 M solution of propylamine, (C3H7NH2 - a weak base) is 12.0. If 0.10 M propylammonium chloride, C 3H7NH3Cl, is added to the solution, will the pH increase, decrease, or remain constant? _______________________ 39. You need to prepare a buffer solution with pH = 11.0. Which of the following components, together with its conjugate is the best choice for this buffer? Circle your choice. (CH3)2NH, Kb = 5.4 x 10-4 HF, Ka = 3.5 x 10-4 NH3, Kb = 1.8 x 10-5 40. Which of the following has the greatest molar solubility? Circle your choice: Al(OH)3, Ksp = 1.9 x 10-33 Fe(OH)3, Ksp = 2.6 x 10-39 Cr(OH)3, Ksp = 6.7 x 10-31 41. A reaction will occur spontaneously at constant temperature and pressure when its ________________ is negative. 42. Circle the Lewis acid in the following reaction: Ag+ (aq) + 2 CN– (aq) ÷ Ag(CN)2– (aq) 43. Aniline (C6H5NH2) is a weak base, Kb = 3.8 x 10–10. What is pKa for C6H5NH3+? 44. Consider the following reaction: AgI (s) + 2 CN– (aq) ÷ Ag(CN)2– (aq) + I– (aq). What is the equilibrium constant _________________________ for this reaction given that for AgI, Ksp = 1.5 x 10–16 and Kf = 5.6 x 10+8 for the following reaction: Ag+ (aq) + 2 CN– (aq) ÷ Ag(CN)2– (aq). 45. Classify each of the following solutions as either acidic (A), basic (B), or neutral (N): a. CH3NH3I (aq) ______________ 46. K = _________________________ b. KI (aq) ______________ Identify the reducing agent in the following chemical reaction: 14 H+ (aq) + Cr2O72–(aq) + 3 Ni (s) Ÿ 3 Ni2+ (aq) + 2 Cr3+ (aq) + 7 H2O (l) 47. c. NaH2PO4 (aq) _______________ _________________________ Consider a Galvanic cell based on the following cell reaction: Mg (s) + 2 Ag+ (aq) Ÿ Mg2+ (aq) + Ag (s) The individual half reactions are: Mg2+ + 2 e– Ÿ Mg; E° = –2.37 V Ag+ + e– Ÿ Ag; E° = 0.80 V Does this reaction proceed spontaneously in the direction written? ___________________________ How many moles of electrons are transferred in this reaction? 2+ ___________________________ 2+ 48. Consider the following Galvanic cell: Co(s) | Co (aq) || Fe (aq) | Fe (s) Write the half-reaction that takes place at the cathode: 49. Consider a Galvanic cell based on the following chemical reaction: Zn (s) + Cu2+ (aq) Ÿ Zn2+ (aq) + Cu (s). This cell operates at a potential of 0.95 V when [Cu2+] = 1.0 x 10–5 M and [Zn2+] = 1.0 M. Calculate E° for this cell at 25°C. 50. Consider the following half-reactions: Cu+ (aq) + e- › Cu (s); E° = 0.52 V NO3- (aq) + 4 H+ (aq) + 3 e- › NO (g) + 2 H2O (l); E° = 0.96 V a. Write the overall cell reaction for this Galvanic cell. b. Write the line notation that describes this cell. c. Calculate DG° for the chemical reaction that occurs in this Galvanic cell. (1 V = 1 J/C) 51. An electrolytic cell is used to plate zinc metal from Zn2+ ions in solution. What current is required to produce 25.0 g of Zn in 30.0 min? (1 amp = 1 C/s) 52. a. What type of emission occurs when Mn-49 decays according to the following equation: 49 Mn Ÿ 49Cr + ??? _________________________________ b. 53. Write the balanced nuclear equation that describes a-emission of Pt-170. Consider a 25.0 g sample of glucose, C6H12O6 (molar mass = 180.18 g/mol). a. How many glucose molecules are there in this sample? b. How many hydrogen atoms are there in this sample? 54. Styrene is a compound containing only carbon and hydrogen. Combustion analysis performed on a 2.78 g sample of styrene yielded 9.40 g of CO2 and 1.92 g of H2O. The molar mass of styrene is 104.2 g/mol. Determine the empirical and molecular formula of styrene. 55. Limestone is composed of calcium carbonate (CaCO3) as well as other compounds. In an analysis, a chemist takes a sample of limestone which has a mass of 413 mg and treats it with oxalic acid (H2C 2O4). A chemical reaction occurs between the calcium carbonate and the acid producing calcium oxalate and other products. CaCO3 (s) + H2C 2O4 (aq) Ÿ CaC2O4 (s) + H2O (l) + CO2 (g) The mass of CaC2O4 obtained is 472 mg. What is the percent by mass of calcium carbonate in the original sample of limestone? 56. Calcualte the volume of a 2.50 M potassium sulfate solution that would be required to react completely with 50.0 mL of a 1.48 M barium chloride solution. 57. A 28.2 g sample of a metal is heated to 95.2˚C and dropped in a calorimeter which contains 100.0 g of water at 25˚C. The final temperature in the calorimeter is 31.0˚C. Assuming no heat is lost to the surroundings or the calorimeter, calculate the specific heat of the metal. The specific heat of H2O (l) is 4.184 J/g•°C. 58. A sample of neon in a rigid 5.00 L flask exerts a pressure of 0.100 atm at 250 K. Argon is then added until the mole fraction of argon is 0.333. The temperature remains at 250 K. a. Calculate the number of mol of Ne in the flask. b. Calculate the mass (in g) of argon that was added c. Calculate the total pressure in the flask after the argon was added. 59. A student in lab titrates an unknown cobalt (II) oxalate dihydrate sample with 23.4 mL of 0.0203 M KMnO4 (aq). The balanced net ionic equation for the titration is: 2 MnO4– (aq) + 5 C2O42– (aq) + 16 H+ (aq) Ÿ 2 Mn2+ (aq) + 10 CO2 (g) + 8 H2O (l) Calculate the mass, in g, of oxalate ion present in the unknown sample. 60. 50.0 mL of HCl (g) at 25°C and 800 torr is dissolved in enough water to produce 125.0 mL of a hydrochloric acid solution. a. What is the molar concentration of the resulting HCl (aq)? b. 61. Suppose that coal, of density 1.5 g/cm3, is carbon. The combustion of carbon is described by the equation C (s) + O2 (g) Ÿ CO2 (g) ∆H = –394 kJ a. Calculate the heat produced when a lump of coal of size 7.0 cm x 6.0 cm x 5.0 cm is burned. b. 62. 63. 64. This HCl (aq) is used in a titration to determine the molar concentration of a sol’n of Ca(OH)2. 31.44 mL of HCl (aq) are required to reach the stoichiometric pt in the titration of 15.0 mL Ca(OH)2. Calculate [Ca(OH)2]. Estimate the mass of water that can be heated from 15°C to 100°C with this piece of coal. For the following elements write the ground state electron configuration and show how the atomic orbitals would be populated for n > 3. Be sure that your atomic orbitals are clearly labeled. (An example of what I want: For krypton, the ground state electron configuration would be [Ar] 4s2 3d10 4p6, and the atomic orbitals for n > 3 would be populated as follows: 4s ↑Ø 3d ↑Ø ↑Ø ↑Ø ↑Ø ↑Ø 4p ↑Ø ↑Ø ↑Ø ) a. cobalt b. bromine a. Using the Bohr model of the atom determine the energy associated with the transition of an electron from n = 3 to n = 7. b. Calculate the wavelength of light associated with this transition. c. What type of eletromagnetic radiation does this wavelength correspond to? ___________________________ d. Is light absorbed or emitted during this transition? ___________________________ l = ___________________________ Draw the orbitals of the l = 1 subshell. Be sure that their orientation with respect to an x, y, z axis system is clear. 65. Consider the following equilibrium: H2 (g) + I2 (g) ÷ 2 HI (g); K = 54.1 at 700 K. 2.00 mol of H2 and 1.50 mol of I2 were put in a previously evacuated 10.0 L vessel, and the temperature was maintained at 700 K. Determine the equilibrium concentrations of all species in the vessel. [H2]eq _________________ [I2]eq _____________________ —7 —1 —1 at 283°C and 5.12 x 10—4 M—1s—1 at 413°C. 66. The rate constant for the decomposition of HI is 3.52 x 10 Determine the activation energy for this reaction. 67. The pH of a 0.100 M solution of iodic acid, HIO3, is 1.91 at 25°C. Determine Ka for iodic acid. 68. Calculate the pH of a 0.0100 M solution of cocaine, a weak base (pKb = 5.59). 69. Calculate the molar solubility of AgI (Ksp = 1.5 x 10–16) in 0.200 M MgI2 (aq). 70. Consider the following reaction at 40°C: H2O (l) ÷ H+ (aq) + OH– (aq). At 40°C, DG°f = –237.1 kJ/mol for H2O (l); DG°f = –156.0 kJ/mol for OH– (aq); and DG°f = 0 kJ/mol for H+ (aq). a. Calculate DG° for this reaction. b. Calculate the equilibrium constant (Kw) at 40°C. c. Calculate the pH of pure water at 40°C. M s [HI]eq __________________ 71. Determine DH for the following reaction: C2H4 (g) + 6 F2 (g) Ÿ 2 CF4 (g) + 4 HF (g). Use the following information and apply Hess's Law of Heat Summation: H2 (g) + F2 (g) Ÿ 2 HF (g); DH = –537 kJ C (s) + 2 F2 (g) Ÿ CF4 (g); DH = –680 kJ 2 C (s) + 2 H2 (g) Ÿ C2H4 (g); DH = +52.3 kJ 72. Determine the standard enthalpy change for the following reaction using DH°f data: 4 NH3 (g) + 5 O2 (g) Ÿ 4 NO (g) + 6 H2O (g) for NH3 (g); DH°f = –46 kJ/mol for NO (g); DH°f = + 90 kJ/mol for H2O (g); DH°f = –242 kJ/mol 73. How much energy will be released when 10.0 g of H2O (l) initially at 80°C is cooled to H2O (s) at –10°C? For H2O (l); s = 4.18 J/g•°C. For H2O (s); s = 2.03 J/g•°C. For H2O at 0°C; DH°fus = 6.01 kJ/mol 74. A biochemical engineer isolates a bacterial gene fragment and dissolves a 10.0 mg sample in enough water to make 30.0 mL of sol'n. The osmotic pressure of the sol'n is 0.340 Torr at 25°C. What is the molar mass of the bacterial gene fragment? 75. a. What volume of ethylene glycol (C2H6O2), a nonelectrolyte, must be added to 15.0 L of water to produce an antifreeze solution with a freezing point of –30.0°C? For C2H6O, d = 1.11 g/cm3; for H2O, d = 1.00 g/cm3, Kf = 1.86 °C•kg•mol–1, Kb = 0.51 °C•kg•mol–1, fp = 0°C, bp = 100°C. b. What is the boiling point of the solution described in a above? 76. 77. Consider an ideal solution of 200.0 mL of pentane (C5H12, d = 0.63 g/mL) and 200.0 mL of hexane (C6H14, d = 0.66 g/mL). The vapor pressures of the pure liquids are 511 Torr for pentane and 150 Torr for hexane. a. b. Which compound is more volatile: pentane or hexane? Determine the total vapor pressure of this solution. ______________________ c. Calculate the mole fraction of pentane in the vapor. a. Which of the following has the highest boiling point? PCl3, NH3, or He ______________________ b. Which of the following has the lowest vapor pressure at 25°C? Cl2, Br2, or I2 ______________________ c. Which of the following is the most polarizeable? F, Cl, Br, or I ______________________ d. Rank the following in order of increasing strength of intermolecular forces: C2H4, LiCl, H2CO, He, CH3OH ____________ < ____________ < ____________ < ____________ < ____________ e. Put the following bonds in order of increasing bond polarity: O – Cl, Cl – Cl, Na – Cl __________ < __________ < __________ 78. 79. The normal boiling point of benzene is 80.1°C. a. What is the vapor pressure of acetone at 80.1°C? ______________________ b. Calculate the vapor pressure of acetone at 25°C. For benzene, DHvap = 30.8 kJ/mol. a. b. Draw the molecular orbital diagram for N2–. Show the atomic orbitals on the left and right sides of the diagram and the molecular orbitals in the middle. Calculate the bond order for this ion. __________________ c. Is this species paramagnetic or diamagnetic? d. e. f. g. __________________ Write the molecular electron configuration for N2 . ______________________________________ Which species would MO theory predict to be more stable, N2– or C2+? __________________ – + Would N2 or C2 have a longer bond? __________________ Would N2– or C2+ have a greater bond dissociation energy? __________________ – 80. Indicate whether each of the following statements is true or false: a. The C–O bonds in the carbonate ion are best described as delocalized p bonds. ____________ 3 b. When one 2s orbital and three 2p orbitals mix, three equivalent sp hybrid orbitals are formed. c. If the central atom in a molecule is sp3d2 hybridized, the bond angles in the molecule are ª109°. ____________ d. sp hybrid orbitals have two lobes with one lobe significantly larger than the other. e. The extent of parallel overlap in p bond formation is typically less than the orbital overlap in a s bond. ____________ Therefore a p bond is weaker than a s bond. 81. ____________ Match the description on the left with the one best term on the right. i. hybridization that results in one unhybridized p orbital ________ a. p bond b. sp2 ii. hybridization resulting in bond angles of 120 and 90° ________ c. double bond d. sp3 iii. geometry around each C in C2H4 ________ e. triple bond f. sp3d iv. one s bond + two p bonds ________ g. s bond h. tetrahedral v. e– density concentrated symmetrically about i. trigonal planar j. linear the internuclear axis 82. 83. ____________ ________ 2 2 The electron configuration for an element is [Ne]3s 3p . a. Identify this element. ____________________________ b. What does [Ne] represent? ____________________________ c. How many unpaired electrons are there in this atom? ____________________________ – d. Give the 4 quantum numbers for one of the e ‘s in the 3s orbital. e. For this element a 3p e– is best described as a(n) {core, valence, noble gas, or paramagnetic} e–. _____________ f. Identify the electron-occupied orbital of highest energy in this atom. a. n = ____; l = ____; ml = ____; ms = _____ ____________________________ – The energy required to remove an e from a gas phase atom or ion is the {ionization energy, electron affinity, enthalpy of formation, or lattice energy}. 84. b. Write the balanced chemical equation that corresponds to the 2nd ionization energy for aluminum. c. For Al, the 2nd ionization energy will be {less than, greater than, or equal to} the 1st ionization energy. a. How many valence electrons does P3– have? b. 85. 86. 87. ____________________________ 3– Give one example of an atom or ion that is isoelectric with P . ____________________________ 12 a. Which electromagnetic radiation has the longer wavelength? n = 4.5 x 10 Hz or n = 6.4 x 1015 Hz b. Which one of the following has the longest wavelength? visible light, UV light, microwaves, IR c. Which light has the greater speed? i. l = 20 nm, ii. l = 800 nm, or iii. neither – both are the same d. Which electromagnetic radiation has greater energy per photon? i. l = 2 x 10–2 nm, or ii. l = 450 nm a. Give all possible values for ml for an electron in a 2p orbital. b. Is the following combination of quantum numbers allowed? n = 3, l = 3, ml = 1, ms = +1/2 c. How many electrons can have the following quantum numbers? n = 3 and ml = 0. Consider the phase diagram for substance Z shown below. a. Estimate the normal melting temperature of Z. b. What phase change (if any) occurs when Z is held at 200 K and undergoes a DP from 1 to 0.001 atm? c. What phase change (if any) occurs when Z is held at 200 atm and undergoes a DT from 300 to 100 K? ____________________ ___________ ________________ Chem 173: 1. Final Exam Review Short Answer and Problems technetium Cu2O yttrium nitrate Mn silver sulfate NH4Cl tin (II) iodide KBrO3 carbon tetrachloride HNO2 (aq) 2. 3 IBr (s) + 4 NH3 (g) Ÿ 1 NI3 (s) + 3 NH4Br (s) 3. 2 C2H2 (g) + 5 O2 (g) Ÿ 4 CO2 (g) + 2 H2O (g) 4. Rb ANSWERS N, Ar, Br Cr, Co, Br Ar N, Br Cr, Co 5. 31, 38, 28 80 Se2– 6. 32 m/s, 3.68 x 1012 mg/m3, 1.0449 x 109 nm 7. 15 8. physical properties: orange-red solid, d = 11.1 g/cm3 chemical properties: decomposes when heated, does not dissolve in water 9. LiOH, H2SO4, MgCl2 NH3, HC2H3O2 P 4O10, CH3OH H2SO4 LiOH LiOH, MgCl2 10. Y, 67.9 11. a. Li2SO4 (aq) + H2O (l) b. AgCl (s) + Cd(NO3)2 (aq) c. K2CO3 (aq) + H2O (l) 12. Ca(OH)2 (aq) + 2 HClO2 (aq) Ÿ Ca(ClO2)2 (aq) + 2 H2O (l) Ca2+ (aq) + 2 OH– (aq) + 2 HClO2 (aq) Ÿ Ca2+ (aq) + 2 ClO2– (aq) + 2 H2O (l) (*because HClO2 is a weak acid) OH– (aq) + HClO2 (aq) Ÿ ClO2– (aq) + H2O (l) 13. a. +4, b. +1, c. –2 14. ox: 4 OH– (aq) + Zn (s) Ÿ Zn(OH)42- (aq) + 2 e– red. 8 e– + 6 H2O (l) + NO3– (aq) Ÿ NH3 (aq) + 9 OH– (aq) net: 7 OH– (aq) + 6 H2O (l) 4 Zn (s) + NO3– (aq) Ÿ NH3 (aq) + 4 Zn(OH)42– (aq); 15. top to bottom: d, e, c, a, b 16. a. endothermic, b. absorbed 17. 409 m/s, increase, increase * Zn is the reducing agent 18. a. l Ÿ g (or g Ÿ l), b. endothermic for l Ÿ g (exothermic for g Ÿ l), c. 61.59 kJ/mol, d. K(l) 19. C (s) + 2 H2 (g) + 1/2 O2 (g) Ÿ CH3OH (l) 20. compound on right – cis-Pt(NH3)2Cl2 is polar (has non-zero dipole moment) 21. a. 109.5°, 180° b. sp3, sp c. s 5, p 2 a. should be 22 electrons or 11 e– pairs b. trigonal bipyramidal c. linear d. sp3d on Xe with sp3 on F e. sp3 f. nonpolar 22. 23 a. water, b. CCl4, c. decrease, d. larger, e. 4, f. lower, g. lower 24. a. 4.62 M, b. 0.101, c. 27.3%, d 6.24 m 25. a. CO32–, 24 electrons or 12 e– pairs; there are 3 total resonance structures b. delocalized c. 2p on C with 2p on O a. HC7H5O2 (aq) + H2O (l) ÷ C7H5O2– (aq) + H3O+ (aq) b. C 7H5O2– c. Ka = [C7H5O2–][H3O+]/[HC7H5O2] d. 6.5 x 10–5 26. 27. 0.5, 1.5, M–.5•s–1, 4.00 M•s–1 28. ln k vs. 1/T 29. K = [CO]/[O2]1/2 30. nc, nc, d, d, i, nc, nc 31. 22.5, 0.500 32. HCl, HBr, HI, HNO3, HClO4, H2SO4 33. H2PO4– acid, SO32– base, HPO42– conjugate base, HSO3– conjugate acid 34. reaction proceeds forward toward equilibrium therefore PC increases and PA and PB decrease 35. strong, weak 36. Ksp = [Cr3+]2[CO32–]3 37. a. Cu2SO4 (s) ÷ 2 Cu+ (aq) + SO42– (aq) b. HC2H3O2 (aq) + OH– (aq) Ÿ C2H3O2– (aq) + H2O (l) c. OCl– (aq) + H2O (l) ÷ HOCl (aq) + OH– (aq) 38. decrease (the equilibrium in question is C3H7NH2 (aq) + H2O (l) ÷ C3H7NH3+ (aq) + OH– (aq)) 39. (CH3)2NH 40. Cr(OH)3 41. DG 42. Ag+ (aq) 43. 4.58 44. K = Ksp • Kf = 8.4 x 10–8 45. a. acidic, b. neutral, c. acidic 46. Ni 47. yes, 2 48. Fe2+ + 2 e– Ÿ Fe 49. 1.1 V 50. a. 3 Cu (s) + NO3– (aq) + 4 H+ (aq) Ÿ NO (g) + 2 H2O (l) + 3 Cu+ (aq), E° = 0.44 V b. Cu (s) | Cu+ (aq) || NO (g) | H+ (aq), NO3– (aq) | inert conductor like Pt (s) or C (gr) c. –127 kJ 51. 41.0 A 52. a. positron emission (b+), b. 53. a. 8.36 x 1022 molecules, b. 1.00 x 1024 H atoms 54. CH, C8H8 55. 89.3% 56. 29.6 mL 57. 1.39 J/g°C 58. a. 0.0244 mol Ne, b. 0.487 g Ar, c. 0.150 atm 59. 0.105 g C2O42– 60. a. 0.0172 M, b. 0.0180 M 61. a. 1.03 x 104 kJ released, b. 2.5 x 104 g 62. a. b. [Ar]4s23d7, 4s ↑Ø 2 10 170 Pt Ÿ a + 166Os 3d ↑Ø ↑Ø ↑ ↑ ↑ 5 [Ar]4s 3d 4p , 4s ↑Ø 3d ↑Ø ↑Ø ↑Ø ↑Ø ↑Ø 63. a. 1.98 x 10–19 J, b. 1000 nm, c. infra-red, d. absorbed 64. see p. 310, figure 7.14(b) in your text 65. [H2] = 0.068 M, [I2] = 0.018 M, [HI] = 0.264 M 66. 177 kJ/mol 67. 1.73 x 10–3 68. 10.20 69. 3.8 x 10–16 mol/L 70. a. +81.1 kJ, b. 2.82 x 10–14, c. 6.77 71. –2486 kJ 72. –908 kJ 73. 6.88 kJ released 74. 1.82 x 104 g/mol 75. a. 1.35 x 104 cm3 or 13.5 L b. 108.2°C 4p ↑Ø ↑Ø ↑ 76. a. pentane, b. 343 Torr, c. 0.796 77. a. NH3, b. I2, c. I, d. He < C2H4 < H2CO < CH3OH < LiCl, e. Cl–Cl < O–Cl < Na–Cl 78. a. 760 Torr, b. 109 Torr 79. a. see me b. paramagnetic c. (s2s)2(s2s*)2(p2p)4(s2p)2(p*2p)1 d. N2– e. C 2+ f. N2– 80, a. T, b. F, c. F, d. T, e. T 81. i. b, ii. f, iii. i, iv. e, v. g 82. a. silicon b. electron configuration of neon: 1s2 2s2 2p6 c. 2 d. n = 3, l = 0, ml = 0, ms = +1/2 or – 1/2 e. valence f. 3p 83. a. ionization energy, b. Al+ (g) Ÿ Al2+ (g) + e–, c. greater than 84. a. 8, b. S2–, Cl–, Ar, K+, Ca2+, Sc3+ 85. a. n = 4.5 x 1012 Hz b. microwaves c. neither – both are the same d. l = 2 x 10–2 nm 86. a. –1, 0, or +1, b. no, c. 6 87. a. 175 K, b. l Ÿ g, c. l Ÿ s