Updated Chem. Final Review PDF

Matching Match each item with the correct statement below. a. proton d. electron b. nucleus e. neutron c. atom _____c__1. the smallest particle of an element that retains the p...

Matching Match each item with the correct statement below. a. proton d. electron b. nucleus e. neutron c. atom _____c__1. the smallest particle of an element that retains the properties of that element __a__ 2. a positively charged subatomic particle ___d_ 3. a negatively charged subatomic particle __e__ 4. a subatomic particle with no charge ___b_ 5. the central part of an atom, containing protons and neutrons Multiple Choice Identify the choice that best completes the statement or answers the question. ____ 6. The sum of the protons and neutrons in an atom equals the ____. a. atomic number b. nucleus number c. atomic mass d. mass number ____ 7. All atoms of the same element have the same ____. a. number of neutrons b. number of protons c. mass numbers d. mass ____ 8. Isotopes of the same element have different a. numbers of neutrons. b. numbers of protons. c. numbers of electrons. d. atomic numbers. ____ 9. The mass number of an element is equal to ____. a. the total number of electrons in the nucleus b. the total number of protons and neutrons in the nucleus c. less than twice the atomic number d. a constant number for the lighter elements ____ 10. How many protons, electrons, and neutrons does an atom with atomic number 50 and mass number 125 contain? a. 50 protons, 50 electrons, 75 neutrons b. 75 electrons, 50 protons, 50 neutrons c. 120 neutrons, 50 protons, 75 electrons d. 70 neutrons, 75 protons, 50 electrons ____ 11. If E is the symbol for an element, which two of the following symbols represent isotopes of the same element? 1. E 2. E 3. E 4. E a. 1 and 2 b. 3 and 4 c. 1 and 4 d. 2 and 3 ____ 12. In which of the following is the number of neutrons correctly represented? a. F has 0 neutrons. b. As has 108 neutrons. c. Mg has 24 neutrons. d. U has 146 neutrons. ____ 13. How do the isotopes hydrogen-1 and hydrogen-2 differ? a. Hydrogen-2 has one more electron than hydrogen-1. b. Hydrogen-2 has one neutron; hydrogen-1 has none. c. Hydrogen-2 has two protons; hydrogen-1 has one. d. Hydrogen-2 has one proton; hydrogen-1 has none. ____ 14. The atomic mass of an element depends upon the a. mass of each electron in that element. b. mass of each isotope of that element. c. relative abundance of protons in that element. d. mass and relative abundance of each isotope of that element. ____ 15. How does the energy of an electron change when the electron moves closer to the nucleus? a. It decreases. b. It increases. c. It stays the same. d. It doubles. ____ 16. The principal quantum number indicates what property of an electron? a. position b. speed c. energy level d. electron cloud shape ____ 17. What is the shape of the 3p atomic orbital? a. sphere b. dumbbell c. bar d. two perpendicular dumbbells ____ 18. According to the aufbau principle, a. an orbital may be occupied by only two electrons. b. electrons in the same orbital must have opposite spins. c. electrons enter orbitals of highest energy first. d. electrons enter orbitals of lowest energy first. ____ 19. What is another name for the transition metals? a. noble gases b. Group A elements c. Group B elements d. Group C elements ____ 20. Each period in the periodic table corresponds to a(n) ____. a. principal energy level b. energy sublevel c. orbital d. suborbital ____ 21. The modern periodic table is arranged in order of increasing atomic ____. a. mass b. charge c. number d. radius ____ 22. Which of the following categories includes the majority of the elements? a. metalloids b. liquids c. metals d. nonmetals ____ 23. To what category of elements does an element belong if it is a poor conductor of electricity? ____ a. transition elements b. metalloids c. nonmetals d. metals ____ 24. Elements that are characterized by the filling of p orbitals are classified as ____. a. groups 3A through 8A b. transition metals c. inner transition metals d. groups 1A and 2A ____ 25. Of the elements Fe, Hg, U, and Te, which is a representative element? a. Fe b. Hg c. U d. Te ____ 26. How does atomic radius change from top to bottom in a group in the periodic table? a. It tends to decrease. b. It tends to increase. c. It first increases, then decreases. d. It first decreases, then increases. ____ 27. Which of the following factors contributes to the increase in atomic size within a group in the periodic table as the atomic number increases? a. more shielding of the electrons in the highest occupied energy level b. an increase in size of the nucleus c. an increase in number of protons d. fewer electrons in the highest occupied energy level ____ 28. Which of the following statements is true about ions? a. Cations form when an atom gains electrons. b. Cations form when an atom loses electrons. c. Anions form when an atom gains protons. d. Anions form when an atom loses protons. ____ 29. What is the energy required to remove an electron from an atom in the gaseous state called? a. nuclear energy b. ionization energy c. shielding energy d. electronegative energy ____ 30. Compared with the electronegativities of the elements on the left side of a period, the electronegativities of the elements on the right side of the same period tend to be ____. a. lower b. higher c. the same d. unpredictable ____ 31. How many valence electrons are in an atom of phosphorus? a. 2 b. 3 c. 4 d. 5 ____ 32. Which of the following elements forms an ion with a 1– charge? a. fluorine b. hydrogen c. potassium d. sodium ____ 33. What is the formula of the ion formed when tin achieves a stable electron configuration? a. Sn b. Sn c. Sn d. Sn ____ 34. The electron configuration of a fluoride ion, F , is _______. a. 1s 2s 2p b. the same as that of a neon atom c. 1s 2s 2p 3s d. the same as th.at of a potassium ion ____ 35. What is the charge on the cation in the ionic compound sodium sulfide? a. 0 b. 1 c. 2 d. 3 ____ 36. What is the net charge of the ionic compound calcium fluoride? a. 2– b. 1–  c. 0 d. 1 ____ 37. A compound composed of cations and anions is called a(n) _______. a. diatomic molecule b. polar compound c. covalent molecule d. ionic compound ____ 38. Which of the following is true about an ionic compound? a. The chemical formula shows the atoms in a molecule. b. The formula unit gives the number of each type of ions in a crystal. c. It is composed of anions and cations and yet it is electrically neutral. d. The chemical formula shows the ions in a molecule. ____ 39. Ionic compounds are normally in which physical state at room temperature? a. solid b. liquid c. gas d. plasma ____ 40. Which molecule has a single covalent bond? a. CO b. Cl c. CO d. N ____ 41. The chemical formula of an ionic compound shows a. how many atoms of each element a molecule contains. b. the lowest whole-number ratio between ions in the ionic compound. c. which molecules the ionic compound contains. d. how the atoms bond. ____ 42. The side-by-side overlap of p orbitals produces what kind of bond? a. alpha bond b. beta bond c. pi bond d. sigma bond ____ 43. Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen? a. C b. Na c. O d. S ____ 44. Which of the following covalent bonds is the most polar? a. H—F b. H—C c. H—H d. H—N ____ 45. When placed between oppositely charged metal plates, the region of a water molecule attracted to the negative plate is the ____. a. hydrogen region of the molecule  b. geometric center of the molecule c. H—O—H plane of the molecule d. oxygen region of the molecule ____ 46. What type of ions have names ending in -ide? a. only cations b. only anions c. only metal ions d. only gaseous ions ____ 47. Which of the following determines that an element is a metal? a. the magnitude of its charge b. the molecules that it forms c. when it is a Group A element d. whether it loses valence electrons ____ 48. In which of the following are the symbol and name for both of the ions given correctly? a. NH : ammonia; H : hydride b. C H O : acetate; C O : oxalite c. OH : hydroxide; O : oxide d. PO : phosphate; PO : phosphite ____ 49. Which of the following correctly provides the names and formulas of polyatomic ions? a. carbonate: HCO ; bicarbonate: CO b. nitrite: NO ; nitrate: NO c. sulfite: S ; sulfate: SO d. chromate: CrO ; dichromate: Cr O ____ 50. Systematic names are preferred over common names because common names a. do not tell you the actual charge of the ion. b. are derived from the method used to obtain the compound. c. were assigned by the scientist who discovered the compound. d. are not very descriptive. ____ 51. What is the correct formula for potassium sulfite? a. KHSO b. KHSO c. K SO d. K SO ____ 52. Binary molecular compounds are made of two ____. a. metallic elements b. nonmetallic elements c. polyatomic ions d. cations ____ 53. In naming a binary molecular compound, the number of atoms of each element present in the molecule is indicated by ____. a. Roman numerals b. superscripts  c. prefixes d. suffixes ____ 54. Which of the following is a binary molecular compound? a. BeHCO b. PCl c. AgI d. MgS ____ 55. When naming acids, the prefix hydro- is used when the name of the acid anion ends in ____. a. -ide b. -ite c. -ate d. -ic ____ 56. When the name of an anion that is part of an acid ends in -ite, the acid name includes the suffix ____. a. -ous b. -ic c. -ate d. -ite ____ 57. What does an -ite or -ate ending in a polyatomic ion mean? a. Oxygen is in the formula. b. Sulfur is in the formula. c. Nitrogen is in the formula. d. Bromine is in the formula. ____ 58. How many hydrogen atoms are in 5 molecules of isopropyl alcohol, C H O? a. 5 (6.02 10 ) b. 5 c. 35 d. 35 (6.02 10 ) ____ 59. How many molecules are in 2.10 mol CO ? a. 2.53 10 molecules b. 3.79 10 molecules c. 3.49 10 molecules d. 1.26 10 molecules ____ 60. How many atoms are in 3.5 moles of arsenic atoms? a. 5.8 10 atoms b. 7.5 10 atoms c. 2.1 10 atoms d. 1.7 10 atoms ____ 61. The atomic masses of any two elements contain the same number of ____. a. atoms b. grams c. ions d. milliliters ____ 62. The mass of a mole of NaCl is the a. molar mass. b. atomic mass. c. molecular mass. d. compound mass. ____ 63. The volume of one mole of a substance is 22.4 L at STP for all ____. a. gases b. liquids c. solids d. compounds ____ 64. The molar volume of a gas at STP occupies ____. a. 22.4 L b. 0C c. 1 kilopascal d. 12 grams ____ ____ 67. Which of the following gases at STP would have the greatest volume? a. 4.00 mole of He b. 1.00 mole of O c. 0.200 mole of SO d. 5.00 mole of H ____ 68. Which of the following compounds has the highest oxygen content, by weight? a. Na O b. CO c. BaO d. H O ____ 69. Which of the following is an empirical formula? a. C N H b. C H O c. Be (Cr O ) d. Sb4S6 ____ 70. The ratio of carbon atoms to hydrogen atoms to oxygen atoms in a molecule of dicyclohexyl maleate is 4 to 6 to 1. What is its molecular formula if its molar mass is 280 g? a. C H O b. C H O c. C H O d. C H O ____ 71. Chemical reactions a. occur only in living organisms. b. create and destroy atoms. c. only occur outside living organisms. d. produce new substances. ____ 72. What does the symbol above the arrow in a chemical equation mean? a. Heat is supplied to the reaction. b. A catalyst is needed in the reaction. c. Electricity is need in the reaction. d. A precipitate will form during the reaction. ____ 73. This symbol ( ) indicates that ____. a. heat must be applied b. an incomplete combustion reaction has occurred c. a gas is formed by the reaction d. the reaction is reversible ____ 74. Rewrite the following word equation as a balanced chemical equation. What is the coefficient and symbol for fluorine? nitrogen trifluoride nitrogen fluorine a. 6F b. F c. 6F d. 3F ____ 75. What are the coefficients that will balance the skeleton equation below? AlCl + NaOH Al(OH) NaCl a. 1, 3, 1, 3 b. 3, 1, 3, 1 c. 1, 1, 1, 3 d. 1, 3, 3, 1 ____ 76. Which of the following statements is true about what happens in all chemical reactions? a. The ways in which atoms are joined together is not changed. b. New atoms are formed as products. c. The final substances are called reactants. d. Bonds between atoms are broken and new bonds are formed. ____ 77. When the equation KClO (s) KCl(s) + O (g) is balanced, the coefficient of KClO3 is a. 1. b. 2. c. 3. d. 4. ____ 78. Which of the following statements is correct? a. All combustion reactions are also combination reactions. b. All chemical reactions can be classified as one of five general types. c. Incomplete combustion of a hydrocarbon may produce carbon monoxide and water. d. A single reactant is the identifying characteristic of a single replacement reaction. ____ 79. When the equation for the complete combustion of one mole of C H OH is balanced, the coefficient for oxygen is ____. a. 13 2  b. 11 2 c.7 2 d. 9 2 ____ 80. Which of the following statements is true about the decomposition of a simple binary compound? a. The products are unpredictable. b. The reactants are the constituent elements. c. The reactant is a single substance. d. The product could be an ionic or a molecular compound. ____ 81. Use the activity series of metals to complete a balanced chemical equation for the following single replacement reaction. Ag(s) KNO (aq) a. AgNO K b. AgK NO c. AgKNO d. No reaction takes place because silver is less reactive than potassium. ____ 82. Which of the following is the correctly balanced equation for the incomplete combustion of heptene, C H ? a. C H 14O 7CO 7H O b. C H 7O 7CO 7H O c. 2C H 21O 14CO 14H O d. C H O C O 7H ____ 83. What is the molarity of a solution containing 7.0 moles of solute in 569 mL of solution? a. 81M b. 0.081M c. 12M d. 4.0M ____ 84. What measures are used to calculate the percent by volume of a solution? a. mass of solute and volume of solvent b. mass of solute and volume of solution c. volume of solute and volume of solvent d. volume of solute and volume of solution ____ ____ 87. In which of the following is concentration expressed in percent by volume? a. 100% b. 100% c. 100%  d. 100% _ ____ 89. What is the mole fraction of ethanol in a solution of 3.00 moles of ethanol and 5.00 moles of water? a. 0.375 b. 0.6 c. 1.67 d. 15 Short Answer 90. Chlorine has two naturally occurring isotopes, Cl-35 and Cl-37. The atomic mass of chlorine is 35.45. Which of these two isotopes of chlorine is more abundant? Cl-35 because it is closer to the actual atomic mass. 91. Give the electron configuration for a neutral atom of chlorine. 1s²2s²2p63s²3p5 92. Which group of elements in the periodic table is known as the alkali metals? Group 1 93. Which group in the periodic table is known as the noble gases? Group 8 Or (18) 94. What is the electron configuration of oxygen? 1s2 2s2 2p4 95. Find the number of moles of argon in 607 g of argon. N = m/m.m = 607 / 40 = 96. Find the mass, in grams, of 1.40 1023 molecules of N. 1. n = # of molecules/ Avog. Number 1.4x1023/ 6.02 x 1023= 0.233 mole 2. m = n x m.m = 0.233 x (2 x14) = 6.51 grams  97. What is the number of moles of solute in 650 mL of a 0.40M solution? N = M x V(L) = 0.4 x 650/1000 = 98. How many liters of a 1.5M solution are required to yield 5.0 grams of solute? (molar mass of solute = 30.0 g) M = n / V so V = n / M n = m /m.m = 5.0 / 30 = 0.167 mole so V = n / M = 0.167 / 1.5 = 0.11 L 99. Calculate the molality of a solution prepared by dissolving 175 g of KNO in 750 g of water. n = m /m.m = 175 / ( 39 + 14 + 3x16 ) = 175 / 101 = 1.73 mole V = 750 / 1000 = 0.75 L so M = n / V = 1.73 / 0.75 = 2.31 molar Numeric Response 100. How many neutrons are present in an atom of the isotope U? Neutrons = 235 – 92 = 143 101. How many electrons are present in the d sublevel of a neutral atom of nickel? 8 102. How many electrons does the ion Ca contain? 20 – 2 = 18 electrons 103. How many valence electrons are in bromine? 7 104. Write balanced chemical equations, and indicate the type of the reaction: 1. Zinc + Silver nitrate Zn (s) + 2AgNO3 (aq) → Zn(NO3)2 (aq) + 2Ag(s) 2. Aluminum + Hydrogen chloride 2Al(s) + 6HCl(aq ) → 3H2(g) + 2AlCl3(aq) 3. Magnesium oxalate + Ammonium carbonate MgC O + (NH ) CO → MgCO + C H N O 2 4 4 2 3 3 2 8 2 4 4. Calcium + Aluminum nitrate  2Al(NO ) + 3Ca → 3Ca(NO ) + 2Al 3 3 3 2 5. Potassium flouride + Lead (II) Nitrate  Pb(NO3)2 + 2KF → PbF2 + 2KNO3 6. Calcium bromide + Silver nitrate 7. CaBr2 (aq) + 2 AgNO3 (aq) → 2 AgBr (s) + Ca(NO ) 3 2 (aq) 8. Sodium chloride + Potassium K(s) + NaCl(aq) = Na(s) + KCl(aq) 9. Magnesium nitrate + ammonium chloride MgCl (aq) + 2NH NO (aq) → Mg(NO )2(aq) + 2NH Cl(aq) 2 4 3 3 4 10. Iron (III) chlorate + calcium 2Fe(ClO3)3(aq) + 3 Ca(s) 3 ca (ClO3)2(aq) + 2 Fe(s). 105. A sample with a molar mass of 34.00 g/mol is found to consist of 0.44g H and 6.92g O. Find its molecular formula. Divide each mass by respective molar mass Moles of H = 0.44 / 1 = 0.44 mole Moles of O = 6.92 / 16 = 0.43 ( round to 0.44 mole Ratio H : O = 1 : 1 Empirical formula = HO Formula mass of emp.formula = 1+16 = 17 g / formula Multiplication factor = 34 g / 17 g = 2 Therefore molecular formula = (HO)2 = H2O2 106. Determine the empirical formula of a compound consisting of : 55.3%K, 14.6%P, and 30.1% O. Assume 100 g of compound. We have 55.3 g of K, 14.6 g P, and 30.1 g O. Look up the molar masses of K, P, and O and convert these to moles. 107. If 4.04g of N combine with 11.46g O to produce a compound with a molar mass of 108.0 g/mol, what is the molecular formula of this compound? 108.A 1 grams of potassium fluoride is dissolved to make 0.10 L of solution. Calculate its molarity KF M = # moles / volume = # moles = mass/ m.m = 1/( 39 + 17 ) = 1/56 = 0.018 mole M.018 / 0.1 = 0.18 M 109.How many molecules are in 0.400 moles of N2O5? # of molecules = # moles x Avog. Number = 0.4 x 6.02 x1023= 110. Calculate the percent composition of Fe2O3 ? % Fe = 2x56 /(2x56 + 3x16 ) x100 = %O = 3x16 / ( 2x56 = 3x16) x100 = 111..A compound composed of hydrogen and oxygen is found to contain 0.59 g of hydrogen and 9.40 g of oxygen. The molar mass of this compound is 34.0 g/mol. Find the empirical and molecular formulas. Moles of H = 0.59 / 1 = 0.59 mole Moles of O = 9.4 / 16 = 0.588 mole Mole ratio 1 : 1 Emp.formula HO 112. A chemist determines that 1.26 g of iron reacts with 0.54 g of oxygen to form rust. What is the percent composition of each element in the new compound? %Fe = m/total mass x 100 = 1.26/(1.26+0.54) = % O = m/total mass x 100 = 0.54/(1.26 +.54 ) x100 = 113.Write the electronic configuration of the following:. Cupper atom Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Calcium Ion Ca+2 1s2 2s2 2p6 3s2 3p6  Chloride Ion Cl-1 1s2 2s 2 2p6 3s2 3p6 114.Arrange the following elements according to their electronegativity. Ca , Cl , O , Na , Si ________________________________________________________________ element period group O > Cl > Si > Na > Ca Ca 4 2 Cl 3 7 O 2 6 Na 3 1 Si 3 4 Name: Chemical formula 1 Calcium phosphate Ca3(PO4)2 2 Ammonium Carbonate [NH 4] 2CO3 3 Aluminum Chloride AlCl3 115. How do ionic bonds differ from metallic bonds?. Ionic bonds are held together by the electrostatic attraction between contrarily charged ions. Metallic bonds are held together by the sharing of electrons between metal atoms 116. State four of the properties of metallic compounds. Very High melting and boiling points Very Good Conductors of heat and electricity Malleable (can be made into different shapes without breaking) Ductile (can be molded into wiring)  Metallic luster (shiny) Sometimes magnetic 117. Explain how atoms (ions) are held together in an ionic bond. Give an example of an ionic compound. In ionic bond the atoms are held together by electrostatic force of attraction. In ionic bond the anions and cations are present in the ratio where the total charge of the compound becomes zero. For example, Let us consider NaCl compound 118. Can some atoms exceed the limits of the octet rule in bonding? If so, give an example. Yes, Examples include sulfur hexafluoride (SF6) and phosphorus pentachloride (PCl5). If all the phosphorus-chlorine bonds in a PCl5 molecule are covalent, it would imply that the phosphorus molecule is violating the octet rule by holding a total of 10 valence electrons