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Chemistry Final Exam Study Guide: Unit 1: 1. Which compound(s) is/are molecular? 1. Cu(NO3)2 2. SF6 3. Phosphorous Trichloride 4. Sodium Sulfate a) 1 only b) 2 only c) 1 and 2 only d) 3 and 4 only 2. In the laboratory, Austin is given a mixture of iron filings, sand and salt. To separate this mixture, Austin uses a magnet, boiling water and a filter. Which statement is true about the process Austin uses to separate this mixture? a) Austin uses only physical changes to separate the mixture’s component. b) Austin uses only chemical changes to separate the mixture’s components. c) Austin uses both physical and chemical changes to separate the mixture’s components. d) Austin uses neither physical nor chemical changes to separate the mixture’s components. 3. Which of the following atoms contains the fewest number of protons? a) 86 Rb b) 79 Br c) 91 Zr d) 106 Pd e) 79 Se 4. What is the mass in grams of 3.0 x 10^23 molecules of CO2 (molar mass of CO2 =44.0 g) a) 22 g b) 44 g c) 66 g d) 88 g 5. Which substance is expected to contain both ionic and covalent bonds? a) NaCl (s) b) C2H5OH (l) c) NH4NO3 (s) d) H2O (l) 6. An unused flashbulb contains magnesium and oxygen. After use, the contents are changed to magnesium oxide but the total mass does not change. This observation can be best explained by the a) Law of constant composition b) Law of multiple proportions c) Avogadro’s law d) Law of conservation of mass 7. Which group of elements is most likely to lose one electron one forming metals? a) Transition metals b) Halogens c) Alkaline earth metals d) Alkali metals e) Noble gases 8. What is the percent by mass of oxygen in magnesium oxide, MgO? a) 20 % b) 40 % c) 50 % d) 60 % 9. The number 10.00 has how many significant figures? a) 1 b) 2 c) 3 d) 4 e) 5 10. Which of the following formulas is incorrect? a) KBr b) CaSO4 c) KNO3 d) AL2O e) Li2O 11. Properties such as color and density, which can be observed or measured without changing the composition of a substance are called _________________ properties. 12. Substances like hydrogen (H2) and oxygen (O2) that are composed of only one type of atom are classified as ________________. 13. A science article refers to a temperature of 300.0 K. The equivalent Celsius temperature is _______. 14. ______ is the atomic symbol for an element that has 24 neutrons and a mass of 45. 15. The energy due to the position of an object is classified as _______________ energy. 16. Those numbers with an unlimited number of significant figures are called ________________________. 17. _______________________ is the numerical relationship between chemical quantities in a balanced equation or formula. 18. The molar mass of cobalt (II) phosphate hexahydrate is ________________. 19. __________________ are atoms of an element with the same number of protons, but a different number of neutrons. 20. The substances on the left side of the chemical equation are called _________________. 21. The products of the reaction between gaseous ammonia (NH3) and copper (II) oxide at high temperature are nitrogen gas, solid copper and water vapor. a) What is the balanced equation for this reaction? Include the state of matter b) If a sample containing 18.1 g of ammonia is reacted with 90.4 g of copper (II) oxide. Which is the limiting reagent? c) How many grams of nitrogen can be produced? d) If the reaction actually gave 6.63 g of nitrogen instead of the calculated in part C, what will be the percent yield of nitrogen? e) How many grams of the excess reactant remains? 22. In a typical sample of magnesium, 78.99 % is magnesium -‐24 (with atomic mass 23.986 amu), 10.00 % magnesium-‐25 (24.985 amu) and 11.01 % magnesium-‐26 (25.985 amu). Calculate the average atomic mass of magnesium. 23. An infant ibuprofen suspension contains 101 mg/ 5.0 L suspension. The recommended dose is 1.2 x 10^ 10 pg/kg body weight. How many mL of this suspension should be given to an infant weighing 18 lbs? (1 lb= 453.59 g) 24. Determine the number of protons, neutrons and electrons in each of the following a) 208/ 82 Pb 4+ _______ protons ________ neutrons ________ electrons b) 34/16 S2-‐ __________ protons _________ neutrons ________ electrons c) 27/13 Al _____________ protons _________ neutrons _________ electrons 25. Fill in the blanks in the following table Cation Anion Formula Name Lithium phosphate Fe2 (SO4) 3 Ca2+ OH-‐ Ammonium chlorate KMnO4 26. Name the following compounds a) H2SO4 b) NI3 c) S2F4 d) BeCl2 x 5H20 27. A compound is decomposed and the masses of its constituent elements are 1.245 g of Ni and 5.381 g of I. Calculate the empirical formula of the compound. 28. Answer the following questions a) Perform the calculation with the correct number of significant figures: 22.81 + 2.2457 b) Express 0.0003711 in scientific notation with 1 significant figure c) How many significant figures are in the quantity 1200.0 29. In a reaction, 34.0 g of chromium (III) oxide reacts with 12.1 g of aluminum to produce chromium and aluminum oxide. If 23.3 g of chromium is produced, what mass of aluminum oxide is produced? 30. Which ion has the same number of electrons as an atom of helium? a) S2-‐ b) P3-‐ c) Be 2+ d) Ca 2+ 31. Which of these are molecular or ionic compounds? a) P4O10 b) SrCl2 c) MgCO3 d) H2SO4 32. Give the name or chemical formula of each of the following: a) K3PO4 b) Copper (II) sulfate tetrahydrate c) NF3 d) Hydrocyanic acid 33. The density of gold is 19,320 kg/ m^3. What is its density in g/cm^3? 34. How many significant figures do the following numbers have? a) 1934 b) 890. c) 0.00120 35. Determine the answer for each of the following a) 17.34 + 4.900 + 23.1 b) 3.9 x 6.05 x (4.2 x 10^2) c) 9.80 – 4.762 d) 14.1/ 5 36. A measurement was taken three times. The correct measurement was 68.1 mL. Circle whether or not the set of measurements is accurate, precise, both or neither a) 78.1 mL, 43.9 mL, 2 mL accurate precise both neither b) 68.1 mL, 68.2 mL, 68.0 mL accurate precise both neither c) 98.0 mL, 98.2 mL, 97.9 mL accurate precise both neither 37. Which of the following are exact numbers? a) There are fifteen books on a shelf b) The weight of a sample is .825 g c) The height of the empire state building d) The number of seconds in a day 38. Check in the appropriate box if the following properties are chemical or physical Properties Physical chemical Water boils at 100 C UV light converts O3 to O2 Sodium metal is soft and can easily be cut Chlorine is a green gas 39. When performing the calculation 34.530 g + 12.1 G + 1,222.34 g, the final answer must have: a) Units of g/cm^3 b) Only one decimal place c) Three sig figs d) Three decimal places 40. The average daytime temperature of Jupiter is 40 C. Calculate its temperature in F? 41. If the amount of mercury in a polluted lake is 0.40 ug/ mL, what is the total mass in kilograms of mercury in the lake if the lake has a surface area of 100 mi^2 and a depth of 20 ft? (volume = surface area x depth) ( 1 mi = 1.6909 km) (1 ft= 0.3048 m) Unit 2 1. What is the oxidation number of each atom in sodium phosphate? a) Na = +1, P= -‐3, O= -‐2 b) Na = +1, P= +5, O= -‐2 c) Na = +1, P = -‐3, O = +2 d) Na = -‐1, P = +5, O = -‐2 e) Na= 0, P = 0, O = 0 2. Which of the following statements is/are correct? i. A system is defined as an object or collection of objects being studied. ii. Surroundings are defined as everything outside of the system being studied. iii. In an exothermic reaction, heat is transferred from the system to the surroundings. a) I only b) ii only c) iii only d) ii and iii e) i, ii, iii 3. A gas mixture contains 2.0 moles of helium and 8.0 moles of carbon dioxide and is held under conditions of standard temperature and pressure. What is the partial pressure of the carbon dioxide? a) 2.0 atm b) 0.2 atm c) 8.0 atm d) 0.8 atm e) none of these 4. Mg (s) + CuSO4 (aq) à MgSO4 (aq) + Cu (s) is an example of: a) Combination reaction b) Decomposition reaction c) Acid-‐ base reaction d) Displacement reaction e) Precipitation reaction 5. At the same temperature and pressure, equal volumes of nitrogen gas and carbon dioxide gas have: a) The same mass b) The same density c) The same number of atoms d) The same number of moles 6. In which one of the following reaction is chlorine oxidized? a) 2Na (s) + Cl2 (g) à 2NaCl (s) b) Cl2 (g) + 2e-‐ à 2Cl-‐ (aq) c) F2 (g) + 2Cl-‐ (aq) à Cl2 (aq) + 2F-‐ (aq) d) 2ClO3 (aq) + 12H+ (aq) + 10e-‐ à Cl2 (g) + 6H2) (l) 7. The reaction 4Al (s) + 3)2 (g) à 2 Al2O3 (s) change in H= -‐3351 kJ is _________ and therefore heat is ___________ by the reaction. a) Exothermic, released b) Exothermic, absorbed c) Endothermic, released d) Endothermic, absorbed 8. If a gas effuses 2.165 times faster than Xe, what is its molar mass? a) 284.3 g/ mol b) 60.65 g/ mol c) 32.00 g/ mol d) 28.01 g/ mol e) 12.94 g/mol 9. In the reaction, NaCl + AgNO3 à AgCl + NaNO3, the white precipitate seen is due to: a) NaCl b) AgNO3 c) AgCl d) NaNO3 10. Which of the following compounds is a nonelectrolyte when dissolved in water? a) NaCl b) MgBr2 c) Cl2 d) Zn(NO3)2 e) KI 11. Consider the following speed distribution curves A and B. If the plots represent the speed distribution of He (g) and Cl2 (g) at STP. Which plot corresponds to Cl2? Curve A _______ Curve B _________ 12. ________ is the energy transferred when a force moves an object 13. _____________ equation shows all reactants and products as if they were intact, undissociated compounds. This gives the least information about the species in solution. 14. A gas absorbs 2.5 J of heat and then performs 13.1 J of work. The change in internal energy of the gas is ______. 15. When a system is absorbing heat from the surroundings its q value should be ___________. 16. _____________ is a substance that produces H+ or H3O + ions when dissolved in H2O. 17. _____________ Law states that at constant pressure and number of moles, the volume occupied by a fixed amount of gas is directly proportional to its absolute (Kelvin) temperature. 18. An exothermic process will have a ___________ delta H value. 19. ________ electrolytes ionize completely and have the capability of conducting electricity. 20. The reducing agent( loses or gains) electrons. 21. It takes 299 cal of energy to heat a sample of pure silver from 12.0 C to 15.2 C. Calculate the mass (in grams) of the sample of silver. The specific heat capacity of silver is 0.24 J/g C 22. At what temperature does 1.00 atm of He gas have the same density as 1.00 atm of Ar gas at 273 K? 23. Given the following choices of reactants, write the balanced molecular, total ionic and net ionic equation (write states of matter). Aqueous sodium sulfide + zinc (II) chloride a) Molecular equation b) Total ionic equation c) Net ionic equation d) Write the spectator ions e) Metals with oxygen gas to produce solid oxides. Write a balanced molecular equation for the reaction of copper to yield copper (I) oxide and then identify the oxidizing reagent. Molecular equation: Oxidizing agent: 24. Categorize each of the following compounds Compound Acid, base, ionic, Electrolyte, non-‐, # of moles of ions molecular weak, strong when dissolved PF3 Sr (OH)2 KMnO4 HI 25. The thermochemical equation for the combustion of benzene is shown below. 2C6H6 (l) + 15 O2 (g) à 12 Co2 (g) + 6 H2O (g) delta H = -‐39093.9 kJ a) Is the reaction absorbing or releasing heat? b) Sketch an enthalpy diagram for the reaction (show reactants, products and direction arrows) c) What is the enthalpy change for the combustion of 12.5 g of C6H6? 26. Hydrogen peroxide, H2O2, can be used as an oxygen source for wastewater treatment. 2H2O2 (aq) à 2H2O (l) + O2 (g) a) Classify the reaction as combination, decomposition or displacement. b) What is the oxidation number of oxygen in the starting material c) Calculate the volume of O2 (g) at 23 C and 726 mmHg that can be liberated from 1.00 L or a 30 % solution of H2O2 in water. The density of the 30 % H2O2 solution is 1.11 g/ cm^3. Unit 3: 1. What is the frequency of the photons emitted by hydrogen atoms when they undergo transitions from n=5 to n=3? 2. Supply the missing quantum numbers and sublevel names N l ml Name of orbital 0 3d 5 1 +1 3 2 -‐2 5s 3. Use Bohr’s equation to calculate how much energy is needed to promote an electron from the H atom ground state to the n=6 level? 4. Calculate the wavelength of a muon (particle with a mass of 1.884 x 10^-‐25 g) traveling at 325 m/s 5. For iron: a) Write the condensed electron configuration b) Draw the partial orbital diagram c) Write a set of quantum numbers for the sixth electron in the d orbital Unit 3: 1. What is the maximum number of electrons in a given atom that can have the quantum number n= 6, l= 3 ? a) 6 b) 7 c) 3 d) 14 e) 10 2. A green color is obtained in the flame test for a barium salt. The energy producing this color is emitted when ________. a) Electrons are raised to higher energy levels by the heat of the flame b) Electrons in the highest levels are expelled c) Electrons drop back to lower energy levels d) Oxidation takes place 3. What is the wavelength of electromagnetic radiation which has a frequency of 4.464 x 10^ 14 s^-‐1? a) 1.338 x 10^ 33 m b) 1.489 x 10^ -‐6 m c) 6.716 x 10^ -‐7 m d) 671.6 nm e) 7.472 x 10^-‐15 nm 4. Which of the following is not isoelectric with a noble gas? a) S^2-‐ b) Ba^+ c) Al^3+ d) Sb^2-‐ e) SC^3+ 5. Which is the electronic configuration of the oxide ion O ^2-‐? a) 1s^2 2s^2 2p^4 b) 1s^2 2s^2 2p^5 c) 1s^2 2s^2 2p^6 d) 1s^2 2s^2 2p^4 3s^2 6. Which of the diatomic elements has a triple bond between its atoms? a) Fluorine b) Oxygen c) Nitrogen d) Hydrogen 7. Which of the following elements has the lowest ionization energy? a) H b) He c) Fr d) Rn e) Rh 8. In the correct Lewis structure for the hydro bromic acid (HBr) molecule, how many unshared electron pairs surround the bromine? a) 3 b) 4 c) 6 d) 8 9. Which of the following is an ionic compound? a) H2S b) NH3 c) I2 d) KI e) CCl4 10. What is a possible set of quantum numbers for the fifth electron filled up in the 4p orbital? [Ar] 3d (all filled up) 4s (filled up) 4p ( five filled up) a) N= 1, l=1, ml= -‐1, ms= +1/2 b) N=3, l=2, ml= -‐1 ,ms = -‐1/2 c) N=4, l=1, ml=0, ms= -‐1/2 d) N=4, l=1, ml=0, ms= +1/2 e) N=4, l=1, ml=1, ms= +1/2 11. ____________ is the energy needed to overcome the attraction between the nuclei and the shared electrons 12. Based on the diagram below, which color of light has the least energy? Violet, blue, blue-‐ green, orange, red 13. An outer-‐ level electron pair that is not involved in bonding is called ___________. 14. According to Heisenberg’s __________ principle, it is impossible to simultaneously measure the exact location and energy of an electron. 15. _______________ is the energy required for the complete removal of 1 mole of electrons from 1 mole of gaseous atoms or ions 16. The d-‐ block elements are also referred to __________________ 17. _____________ principle states that no two electrons in the same atom can have the same four quantum numbers 18. covalent bonding involves ____________ of valence electrons between atoms. 19. The _______________ of a photon is equal to the speed of light divided by the wavelength of the photon. 20. ___________ rule states that the most stable arrangement of electrons is that which contains the maximum number of unpaired electrons, all with the same spin direction. 21. What is the uncertainty in velocity (m/s) of a 25.0 g particle in space that has an uncertainty in its position of 552 pm. 22. How many inner, outer and valence electrons are present in an atom of each of the following elements? -‐ Mn -‐ Se -‐ Ba 23. Based on the periodic table trends: a) Arrange Ca 2+, Cl -‐, K +, P 3 and S 2-‐ in order of increasing ionic size b) Arrange the elements in order of increasing ionization energy: Li, Cs, Cl, Ar, I c) Arrange the elements in order of increasing electronegativity: K, Be, C, K, F, Cs 24. Consider the following two wavelengths: Wave A: Wave B: a) Which wave has the shorter wavelength? b) Which wave has the shorter frequency? c) Which wave has the higher energy? d) Suppose that wave B represents yellow light (570 nm). If you double the wavelength of this light, is it in the visible, IR, or UV region of the electromagnetic spectrum? 25. Below are diagrams for the bright line spectra of four elements and the spectrum of a mixture of unknown gases: Which element (s) are present in the unknown? 26. For the following hydrogen atom transitions: A: n= 1 to n=2 B: n= 2 to n=1 C: n= 2 to n=3 D: n=3 to n=2 E: n= 3 to n=4 F: n=4 to n=3 G: n=1 to n=4 H: n=4 to n=1 I: n=1 to n=3 J: n=3 to n=1 a) Which transition (s) correspond to emissions? b) Which emission’s transition corresponds to the longest wavelength? 27. Write the condensed electron configurations for the following species: a) 1. P 3-‐ 2. Ag + 3. Rh 2+ 4. Ga + b) Fill the partial box diagram for the condensed electronic configuration of the species in part a and state if the atom or ion will be paramagnetic or diamagnetic 1. P 3-‐ 2. Ag + 3. Rh 2+ 4. Ga + c) Write a full set of quantum numbers for the third electron in the d orbital of Ag + d) Write the formula of the oxide Ag + ion forms 28. An electron in a hydrogen atom relaxes to the n=4 level, emitting light of frequency equal to 1.14 x 10^14 Hz. What is the value of n for the level in which the electron originated? 29. X a) Write the Lewis electron dot structure for the following species 1. Cs2 2. MgBr2 b) For Cs2 molecule, answer the following: 1. Bond order 2. Number of total lone pairs electrons in the molecule c) Rank the relative lengths of the N=N, N-‐F, N-‐Br, N-‐Cl and N=-‐ N bonds Chemistry 111 Final Exam Answer Key: Unit 1: 1. C 2. A 3. E 4. A 5. C 6. D 7. B 8. B 9. D 10. D 11. Physical 12. Elements 13. 27 14. Sc 15. Potential 16. Exact numbers 17. Stoichiometry 18. 474.73 g/mol 19. isotopes 20. reactants 21. a) 2NH3 (g) + 3CuO (g) à N2 (g) + 3Cu (s) + 3H2O (aq) b) 1.14 mol c) 10.6 g d) 62.4 % e) x 22. 24.31 amu 23. 4.9 mL 24. a) 82; 126; 78 b) 16; 18; 18 c) 13; 14; 13 25. x Cation Anion Formula Name Li + PO4 3-‐ Li3PO4 Lithium phosphate Fe 3+ So4 2-‐ Fe2(SO4)3 Iron (II) sulfate Ca 2+ OH-‐ CaOH2 Calcium hydroxide NH 4+ CLO3 -‐ NH4ClO3 Ammonium chlorate K+ MnO4 -‐ KMnO4 Potassium permaganate 26. a) sulfuric acid b) nitrogen triodide c) disulfur tetrafluoride d) pentrahydrate 27. N1I2 28. A) 25.05 b) 4 x 10^-‐4 c) 5 29. 22.5 g 30. C 31. A) molecular b) ionic c) ionic d) molecular 32. A) potassium phosphate b) CuSO4 x 4H2O c) Trifluoride d) HCN 33. 19.320 g/cm^3 34. a) 4 b) 3 c)3 35. a) 45.3 b) 5.04 c) 9.9 x 10^3 d) 3 36. a) neither b) both c) precise 37. a) yes b) no c) no d) yes 38. a) physical b) chemical c) physical d) physical 39. b 40. 104 F 41. 6.4 x 10^5 kg/ mL Unit 2: 1. B 2. E 3. D 4. D 5. D 6. C 7. A 8. D 9. C 10. C 11. Curve A 12. Work 13. Molecular 14. -‐10.6 J 15. positive 16. acid 17. charles’s 18. negative 19. strong 20. loses 21. 1.6 x 10^3 g 22. 28.0 K 23. a) Na2S (aq) + ZnCl2 (aq) à 2NaCl (aq) + ZnS (s) b) 2Na + (aq) + S2-‐ (aq) + Zn2+ (aq) + 2Cl-‐ (aq) à 2Na + (aq) + 2Cl-‐ (aq) + ZnS (s) c) S2-‐ (aq) + Zn2+ (aq) à ZnS (s) d) Na +, Cl-‐ e) Cu + O2 à Cu2O; Cu2O molecular: 4Cu (s) + O2 (g) à 2Cu2O (s) oxidizing: O2 24. a) molecular, non, 0 b) base, strong, 3 c) ionic, strong, 2 d) acid, strong, 2 25. a) releasing b) look at diagram c) -‐313 kJ 26. a) decomposition b) -‐1 c) 125 L Unit 3: 1. D; 14 2. C; electrons drop back to lower energy levels 3. D 4. B 5. C 6. C 7. C 8. A 9. D 10. C 11. Bond Energy 12. Red 13. Lone pair 14. Uncertainty 15. Ionization energy 16. Transition metals 17. Pauli’s 18. Sharing 19. Frequency 20. Hund’s 21. 3.82 x 10^ -‐24 m/s 22. a) 18, 2, 7 b) 28, 6, 6 c) 54, 2, 2 23. a) Ca 2+ < K+ < Cl-‐ < S 2-‐ < P 3-‐ b) Cs < Li < I < Cl < Ar c) Cs < K < Be < C < F 24. a) A b) B c) A d) IR 25. He, H 26. A) F, D, B, H, J b) F 27. [Ne] 3s^2 3p^6 [Kr] 4d^10 [Kr] 4d^7 [Ar] 4s^2 3d^10 b) first box: 2 second box: 0 third box: 6; diamagnetic first box: 0 second box: 10 diamagnetic first box: 0 second box: 7paramagnetic first box: 2 second box: 10 diamagentic c) n=4, l= 2, ml=0, ms= +1/2 d) Ag2O 28. n=6 29. a) S (four dots)= C= S (four dots) Mg-‐ Br (six dots) B) 2 double bonds; 4 lone pairs c) N-‐Br > N-‐Cl > N-‐F > N=N > N=-‐N
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