1. Does a cation gain protons to form a positive charge or does it lose electrons? 2. Iron(III) sulfate [Fe 2 (SO 4 ) 3 ] is composed of Fe 3 + and SO 4 2 − SO 4 2 − ions. Explain why a sample of iron(III) sulfate is uncharged. 3. Which of the following atoms would be expected to form negative ions in binary ionic compounds and which would be expected to form positive ions: P, I, Mg, Cl, In, Cs, O, Pb, Co? 4. Which of the following atoms would be expected to form negative ions in binary ionic compounds and which would be expected to form positive ions: Br, Ca, Na, N, F, Al, Sn, S, Cd? 5. Predict the charge on the monatomic ions formed from the following atoms in binary ionic compounds: (a) P (b) Mg (c) Al (d) O (e) Cl (f) Cs 6. Predict the charge on the monatomic ions formed from the following atoms in binary ionic compounds: (a) I (b) Sr (c) K (d) N (e) S (f) In 7. Write the electron configuration for each of the following ions: (a) As 3 – (b) I – (c) Be 2 + (d) Cd 2 + (e) O 2– (f) Ga 3 + (g) Li + (h) N 3– (i) Sn 2 + (j) Co 2 + (k) Fe 2 + (l) As 3 + 8. Write the electron configuration for the monatomic ions formed from the following elements (which form the greatest concentration of monatomic ions in seawater): (a) Cl (b) Na (c) Mg (d) Ca (e) K (f) Br (g) Sr (h) F 9. Write out the full electron configuration for each of the following atoms and for the monatomic ion found in binary ionic compounds containing the element: (a) Al (b) Br (c) Sr (d) Li (e) As (f) S 10. From the labels of several commercial products, prepare a list of six ionic compounds in the products. For each compound, write the formula. (You may need to look up some formulas in a suitable reference.) 11. Why is it incorrect to speak of a molecule of solid NaCl? 12. What information can you use to predict whether a bond between two atoms is covalent or ionic? 13. Predict which of the following compounds are ionic and which are covalent, based on the location of their constituent atoms in the periodic table: (a) Cl 2 CO (b) MnO (c) NCl 3 (d) CoBr 2 (e) K 2 S (f) CO (g) CaF 2 (h) HI (i) CaO (j) IBr (k) CO 2 14. Explain the difference between a nonpolar covalent bond, a polar covalent bond, and an ionic bond. 15. From its position in the periodic table, determine which atom in each pair is more electronegative: (a) Br or Cl (b) N or O (c) S or O (d) P or S (e) Si or N (f) Ba or P (g) N or K 16. From its position in the periodic table, determine which atom in each pair is more electronegative: (a) N or P (b) N or Ge (c) S or F (d) Cl or S (e) H or C (f) Se or P (g) C or Si 17. From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity: (a) C, F, H, N, O (b) Br, Cl, F, H, I (c) F, H, O, P, S (d) Al, H, Na, O, P (e) Ba, H, N, O, As 18. From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity: (a) As, H, N, P, Sb (b) Cl, H, P, S, Si (c) Br, Cl, Ge, H, Sr (d) Ca, H, K, N, Si (e) Cl, Cs, Ge, H, Sr 19. Which atoms can bond to sulfur so as to produce a positive partial charge on the sulfur atom? 20. Which is the most polar bond? (a) C–C (b) C–H (c) N–H (d) O–H (e) Se–H 21. Identify the more polar bond in each of the following pairs of bonds: (a) HF or HCl (b) NO or CO (c) SH or OH (d) PCl or SCl (e) CH or NH (f) SO or PO (g) CN or NN 22. Which of the following molecules or ions contain polar bonds? (a) O 3 (b) S 8 (c) O 2 2− O 2 2− (d) NO 3 − NO 3 − (e) CO 2 (f) H 2 S (g) BH 4 − BH 4 − 23. Name the following compounds: (a) CsCl (b) BaO (c) K 2 S (d) BeCl 2 (e) HBr (f) AlF 3 24. Name the following compounds: (a) NaF (b) Rb 2 O (c) BCl 3 (d) H 2 Se (e) P 4 O 6 (f) ICl 3 25. Write the formulas of the following compounds: (a) rubidium bromide (b) magnesium selenide (c) sodium oxide (d) calcium chloride (e) hydrogen fluoride (f) gallium phosphide (g) aluminum bromide (h) ammonium sulfate 26. Write the formulas of the following compounds: (a) lithium carbonate (b) sodium perchlorate (c) barium hydroxide (d) ammonium carbonate (e) sulfuric acid (f) calcium acetate (g) magnesium phosphate (h) sodium sulfite 27. Write the formulas of the following compounds: (a) chlorine dioxide (b) dinitrogen tetraoxide (c) potassium phosphide (d) silver(I) sulfide (e) aluminum fluoride trihydrate (f) silicon dioxide 28. Write the formulas of the following compounds: (a) barium chloride (b) magnesium nitride (c) sulfur dioxide (d) nitrogen trichloride (e) dinitrogen trioxide (f) tin(IV) chloride 29. Each of the following compounds contains a metal that can exhibit more than one ionic charge. Name these compounds: (a) Cr 2 O 3 (b) FeCl 2 (c) CrO 3 (d) TiCl 4 (e) CoCl 2 ∙6H 2 O (f) MoS 2 30. Each of the following compounds contains a metal that can exhibit more than one ionic charge. Name these compounds: (a) NiCO 3 (b) MoO 3 (c) Co(NO 3 ) 2 (d) V 2 O 5 (e) MnO 2 (f) Fe 2 O 3 31. The following ionic compounds are found in common household products. Write the formulas for each compound: (a) potassium phosphate (b) copper(II) sulfate (c) calcium chloride (d) titanium(IV) oxide (e) ammonium nitrate (f) sodium bisulfate (the common name for sodium hydrogen sulfate) 32. The following ionic compounds are found in common household products. Name each of the compounds: (a) Ca(H 2 PO 4 ) 2 (b) FeSO 4 (c) CaCO 3 (d) MgO (e) NaNO 2 (f) KI 33. What are the IUPAC names of the following compounds? (a) manganese dioxide (b) mercurous chloride (Hg 2 Cl 2 ) (c) ferric nitrate [Fe(NO 3 ) 3 ] (d) titanium tetrachloride (e) cupric bromide (CuBr 2 ) 34. Write the Lewis symbols for each of the following ions: (a) As 3 – (b) I – (c) Be 2 + (d) O 2– (e) Ga 3+ (f) Li + (g) N 3– 35. Many monatomic ions are found in seawater, including the ions formed from the following list of elements. Write the Lewis symbols for the monatomic ions formed from the following elements: (a) Cl (b) Na (c) Mg (d) Ca (e) K (f) Br (g) Sr (h) F 36. Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed: (a) MgS (b) Al 2 O 3 (c) GaCl 3 (d) K 2 O (e) Li 3 N (f) KF 37. In the Lewis structures listed here, M and X represent various elements in the third period of the periodic table. Write the formula of each compound using the chemical symbols of each element: (a) (b) (c) (d) 38. Write the Lewis structure for the diatomic molecule P 2 , an unstable form of phosphorus found in high-temperature phosphorus vapor. 39. Write Lewis structures for the following: (a) H 2 (b) HBr (c) PCl 3 (d) SF 2 (e) H 2 CCH 2 (f) HNNH (g) H 2 CNH (h) NO – (i) N 2 (j) CO (k) CN – 40. Write Lewis structures for the following: (a) O 2 (b) H 2 CO (c) AsF 3 (d) ClNO (e) SiCl 4 (f) H 3 O + (g) NH 4 + NH 4 + (h) BF 4 − BF 4 − (i) HCCH (j) ClCN (k) C 2 2+ C 2 2+ 41. Write Lewis structures for the following: (a) ClF 3 (b) PCl 5 (c) BF 3 (d) PF 6 − PF 6 − 42. Write Lewis structures for the following: (a) SeF 6 (b) XeF 4 (c) SeCl 3 + SeCl 3 + (d) Cl 2 BBCl 2 (contains a B–B bond) 43. Write Lewis structures for: (a) PO 4 3− PO 4 3− (b) ICl 4 − ICl 4 − (c) SO 3 2− SO 3 2− (d) HONO 44. Correct the following statement: “The bonds in solid PbCl 2 are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl 2 are located on the Cl – ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms.” 45. Write Lewis structures for the following molecules or ions: (a) SbH 3 (b) XeF 2 (c) Se 8 (a cyclic molecule with a ring of eight Se atoms) 46. Methanol, H 3 COH, is used as the fuel in some race cars. Ethanol, C 2 H 5 OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO 2 and H 2 O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas. 47. Many planets in our solar system contain organic chemicals including methane (CH 4 ) and traces of ethylene (C 2 H 4 ), ethane (C 2 H 6 ), propyne (H 3 CCCH), and diacetylene (HCCCCH). Write the Lewis structures for each of these molecules. 48. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl 2 CO. Write the Lewis structures for carbon tetrachloride and phosgene. 49. Identify the atoms that correspond to each of the following electron configurations. Then, write the Lewis symbol for the common ion formed from each atom: (a) 1 s 2 2 s 2 2 p 5 (b) 1 s 2 2 s 2 2 p 6 3 s 2 (c) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 (d) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 4 (e) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 1 50. The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms. (a) the amino acid serine: (b) urea: (c) pyruvic acid: (d) uracil: (e) carbonic acid: 51. A compound with a molar mass of about 28 g/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound. 52. A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound. 53. Two arrangements of atoms are possible for a compound with a molar mass of about 45 g/mol that contains 52.2% C, 13.1% H, and 34.7% O by mass. Write the Lewis structures for the two molecules. 54. How are single, double, and triple bonds similar? How do they differ? 55. Write resonance forms that describe the distribution of electrons in each of these molecules or ions. (a) selenium dioxide, OSeO (b) nitrate ion, NO 3 − NO 3 − (c) nitric acid, HNO 3 (N is bonded to an OH group and two O atoms) (d) benzene, C 6 H 6 : (e) the formate ion: 56. Write resonance forms that describe the distribution of electrons in each of these molecules or ions. (a) sulfur dioxide, SO 2 (b) carbonate ion, CO 3 2− CO 3 2− (c) hydrogen carbonate ion, HCO 3 − HCO 3 − (C is bonded to an OH group and two O atoms) (d) pyridine: (e) the allyl ion: 57. Write the resonance forms of ozone, O 3 , the component of the upper atmosphere that protects the Earth from ultraviolet radiation. 58. Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, NO 2 – . NO 2 – . 59. In terms of the bonds present, explain why acetic acid, CH 3 CO 2 H, contains two distinct types of carbon-oxygen bonds, whereas the acetate ion, formed by loss of a hydrogen ion from acetic acid, only contains one type of carbon-oxygen bond. The skeleton structures of these species are shown: 60. Write the Lewis structures for the following, and include resonance structures where appropriate. Indicate which has the strongest carbon-oxygen bond. (a) CO 2 (b) CO 61. Toothpastes containing sodium hydrogen carbonate (sodium bicarbonate) and hydrogen peroxide are widely used. Write Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule, with resonance forms where appropriate. 62. Determine the formal charge of each element in the following: (a) HCl (b) CF 4 (c) PCl 3 (d) PF 5 63. Determine the formal charge of each element in the following: (a) H 3 O + (b) SO 4 2− SO 4 2− (c) NH 3 (d) O 2 2− O 2 2− (e) H 2 O 2 64. Calculate the formal charge of chlorine in the molecules Cl 2 , BeCl 2 , and ClF 5 . 65. Calculate the formal charge of each element in the following compounds and ions: (a) F 2 CO (b) NO – (c) BF 4 − BF 4 − (d) SnCl 3 − SnCl 3 − (e) H 2 CCH 2 (f) ClF 3 (g) SeF 6 (h) PO 4 3− PO 4 3− 66. Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures: (a) O 3 (b) SO 2 (c) NO 2 − NO 2 − (d) NO 3 − NO 3 − 67. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in nitrosyl chloride: ClNO or ClON? 68. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH? 69. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in sulfur dioxide: OSO or SOO? 70. Draw the structure of hydroxylamine, H 3 NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges? 71. Iodine forms a series of fluorides (listed here). Write Lewis structures for each of the four compounds and determine the formal charge of the iodine atom in each molecule: (a) IF (b) IF 3 (c) IF 5 (d) IF 7 72. Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound. 73. Which of the following structures would we expect for nitrous acid? Determine the formal charges: 74. Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H 2 SO 4 , which has two oxygen atoms and two OH groups bonded to the sulfur. 75. Explain why the HOH molecule is bent, whereas the HBeH molecule is linear. 76. What feature of a Lewis structure can be used to tell if a molecule’s (or ion’s) electron-pair geometry and molecular structure will be identical? 77. Explain the difference between electron-pair geometry and molecular structure. 78. Why is the H–N–H angle in NH 3 smaller than the H–C–H bond angle in CH 4 ? Why is the H–N–H angle in NH 4 + NH 4 + identical to the H–C–H bond angle in CH 4 ? 79. Explain how a molecule that contains polar bonds can be nonpolar. 80. As a general rule, MX n molecules (where M represents a central atom and X represents terminal atoms; n = 2 – 5) are polar if there is one or more lone pairs of electrons on M. NH 3 (M = N, X = H, n = 3) is an example. There are two molecular structures with lone pairs that are exceptions to this rule. What are they? 81. Predict the electron pair geometry and the molecular structure of each of the following molecules or ions: (a) SF 6 (b) PCl 5 (c) BeH 2 (d) CH 3 + CH 3 + 82. Identify the electron pair geometry and the molecular structure of each of the following molecules or ions: (a) IF 6 + IF 6 + (b) CF 4 (c) BF 3 (d) SiF 5 − SiF 5 − (e) BeCl 2 83. What are the electron-pair geometry and the molecular structure of each of the following molecules or ions? (a) ClF 5 (b) ClO 2 − ClO 2 − (c) TeCl 4 2− TeCl 4 2− (d) PCl 3 (e) SeF 4 (f) PH 2 − PH 2 − 84. Predict the electron pair geometry and the molecular structure of each of the following ions: (a) H 3 O + (b) PCl 4 − PCl 4 − (c) SnCl 3 + SnCl 3 + (d) BrCl 4 − BrCl 4 − (e) ICl 3 (f) XeF 4 (g) SF 2 85. Identify the electron pair geometry and the molecular structure of each of the following molecules: (a) ClNO (N is the central atom) (b) CS 2 (c) Cl 2 CO (C is the central atom) (d) Cl 2 SO (S is the central atom) (e) SO 2 F 2 (S is the central atom) (f) XeO 2 F 2 (Xe is the central atom) (g) ClOF 2 + ClOF 2 + (Cl is the central atom) 86. Predict the electron pair geometry and the molecular structure of each of the following: (a) IOF 5 (I is the central atom) (b) POCl 3 (P is the central atom) (c) Cl 2 SeO (Se is the central atom) (d) ClSO + (S is the central atom) (e) F 2 SO (S is the central atom) (f) NO 2 − NO 2 − (g) SiO 4 4− SiO 4 4− 87. Which of the following molecules and ions contain polar bonds? Which of these molecules and ions have dipole moments? (a) ClF 5 (b) ClO 2 − ClO 2 − (c) TeCl 4 2− TeCl 4 2− (d) PCl 3 (e) SeF 4 (f) PH 2 − PH 2 − (g) XeF 2 88. Which of these molecules and ions contain polar bonds? Which of these molecules and ions have dipole moments? (a) H 3 O + (b) PCl 4 − PCl 4 − (c) SnCl 3 − SnCl 3 − (d) BrCl 4 − BrCl 4 − (e) ICl 3 (f) XeF 4 (g) SF 2 89. Which of the following molecules have dipole moments? (a) CS 2 (b) SeS 2 (c) CCl 2 F 2 (d) PCl 3 (P is the central atom) (e) ClNO (N is the central atom) 90. Identify the molecules with a dipole moment: (a) SF 4 (b) CF 4 (c) Cl 2 CCBr 2 (d) CH 3 Cl (e) H 2 CO 91. The molecule XF 3 has a dipole moment. Is X boron or phosphorus? 92. The molecule XCl 2 has a dipole moment. Is X beryllium or sulfur? 93. Is the Cl 2 BBCl 2 molecule polar or nonpolar? 94. There are three possible structures for PCl 2 F 3 with phosphorus as the central atom. Draw them and discuss how measurements of dipole moments could help distinguish among them. 95. Describe the molecular structure around the indicated atom or atoms: (a) the sulfur atom in sulfuric acid, H 2 SO 4 [(HO) 2 SO 2 ] (b) the chlorine atom in chloric acid, HClO 3 [HOClO 2 ] (c) the oxygen atom in hydrogen peroxide, HOOH (d) the nitrogen atom in nitric acid, HNO 3 [HONO 2 ] (e) the oxygen atom in the OH group in nitric acid, HNO 3 [HONO 2 ] (f) the central oxygen atom in the ozone molecule, O 3 (g) each of the carbon atoms in propyne, CH 3 CCH (h) the carbon atom in Freon, CCl 2 F 2 (i) each of the carbon atoms in allene, H 2 CCCH 2 96. Draw the Lewis structures and predict the shape of each compound or ion: (a) CO 2 (b) NO 2 − NO 2 − (c) SO 3 (d) SO 3 2− SO 3 2− 97. A molecule with the formula AB 2 , in which A and B represent different atoms, could have one of three different shapes. Sketch and name the three different shapes that this molecule might have. Give an example of a molecule or ion for each shape. 98. A molecule with the formula AB 3 , in which A and B represent different atoms, could have one of three different shapes. Sketch and name the three different shapes that this molecule might have. Give an example of a molecule or ion that has each shape. 99. Draw the Lewis electron dot structures for these molecules, including resonance structures where appropriate: (a) CS 3 2− CS 3 2− (b) CS 2 (c) CS (d) predict the molecular shapes for CS 3 2− CS 3 2− and CS 2 and explain how you arrived at your predictions 100. What is the molecular structure of the stable form of FNO 2 ? (N is the central atom.) 101. A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen. What is its molecular structure? 102. Use the simulation to perform the following exercises for a two-atom molecule: (a) Adjust the electronegativity value so the bond dipole is pointing toward B. Then determine what the electronegativity values must be to switch the dipole so that it points toward A. (b) With a partial positive charge on A, turn on the electric field and describe what happens. (c) With a small partial negative charge on A, turn on the electric field and describe what happens. (d) Reset all, and then with a large partial negative charge on A, turn on the electric field and describe what happens. 103. Use the simulation to perform the following exercises for a real molecule. You may need to rotate the molecules in three dimensions to see certain dipoles. (a) Sketch the bond dipoles and molecular dipole (if any) for O 3. Explain your observations. (b) Look at the bond dipoles for NH 3 . Use these dipoles to predict whether N or H is more electronegative. (c) Predict whether there should be a molecular dipole for NH 3 and, if so, in which direction it will point. Check the molecular dipole box to test your hypothesis. 104. Use the Molecule Shape simulator to build a molecule. Starting with the central atom, click on the double bond to add one double bond. Then add one single bond and one lone pair. Rotate the molecule to observe the complete geometry. Name the electron group geometry and molecular structure and predict the bond angle. Then click the check boxes at the bottom and right of the simulator to check your answers. 105. Use the Molecule Shape simulator to explore real molecules. On the Real Molecules tab, select H 2 O. Switch between the “real” and “model” modes. Explain the difference observed. 106. Use the Molecule Shape simulator to explore real molecules. On the Real Molecules tab, select “model” mode and S 2 O. What is the model bond angle? Explain whether the “real” bond angle should be larger or smaller than the ideal model angle.