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AP Chem Problem Set # 6

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Lewis Symbols and Ionic Bonding

1. (a) What are valence electrons? (b) How many valence electrons does a nitrogen atom posses? (c) An atom has the electron configuration 1s22s22p63s23p2. How many valence electrons does the atom have?

2. Write the Lewis symbol for atoms of each of the following elements: (a) Ca; (b) P; (c) Ne; (d) B.

3. Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance MgO.

4. Predict the chemical formula of the ionic compound formed between the following pairs of elements: (a) Al and F; (b) K and S; (c) Y and O; (d) Mg and N.

5. Write the electron configuration for each of the following ions, and determine which one possesses noble-gas configurations: (a) Sr2+; (b) Ti2+; (c) Se2-; (d) Ni2+; (e) Br-; (f) Mn3+.

6. (a) Does the lattice energy on an ionic solid increase or decrease (i) as the charges of the ions increase; (ii) as the sizes of the ions increase? (b) Using a periodic table, arrange the following substances according to their expected lattice energies, listing them from lowest lattice energy to the highest: LiCl, KCl, KBr, CaO. Compare your list with the data in Table 8.2.

7. Explain the following trends in lattice energy: (a) MgO > MgS; (b) LiF > CsBr; (c) CaO > KF.

8. Energy is required to remove two electrons from Ca to form Ca2+ and is also required to add two electrons to O to form O2-. Why, then, is CaO stable relative to the free elements?

Sizes of Ions

9. (a) Why are monoatomic cations smaller than their corresponding neutral atoms? (b) Why are monoatomic anions larger than their corresponding neutral atoms? (c) Why does the size of ions increase as one proceeds down a column in the periodic table?

10. (a) Why do the radii of isoelectronic ions decrease with increasing nuclear charge? (b) Which experiences the greater effective nuclear charge, a 2p electron in F-, a 2p electron in Ne, or a 2p electron in Na+?

11. Arrange the atoms and ions in each of the following sets in order of increasing size: (a) Li+, Rb+, K+; (b) Br-, Na+, Mg2+; (c) Ar, Cl-, S2-, K+; (d) Cl, Cl-, Ar.

Covalent Bonding, Electronegativity, and Bond Polarity

12. (a) Construct a Lewis structure for O2 in which each atom achieves an octet of electrons. (b) Explain why it is necessary to form a double bond in the Lewis structure. (c) The bond in O2 is shorter than the O-O bond in compounds that contain an O-O single bond. Explain this observation.

13. Using only the periodic table, select the most electronegative atom in each of the following sets: (a) P, S, As, Se; (b) Be, B, C, Si; (c) Zn, Ga, Ge, As; (d) Na, Mg, K, Ca.

14. Which of the following bonds are polar: (a) P-O; (b) S-F; (c) Br-Br; (d) O-Cl? Which is the more electronegative atom in each polar bond?

Lewis Structures; Resonance Structures

15. Draw Lewis structures for the following: (a) SiH4; (b) CO; (c) SF2; (d) H2SO4 (H is bonded to O); (e) ClO2-; (f) NH2OH (N and O are bonded to one another).

16. Write Lewis structures that obey the octet rule for each of the following, and assign formal charges to each atom: (a) NO+; (b) POCl3 (P is bonded to three Cls and to the O); (c) ClO4-; (d) HClO3 (H is bonded to O).

17. Draw the resonance structures that obey the octet rule for the following: (a) N2O (a nitrogen atom is the central atom); (b) CO32-; (c) HCO2- (H and both O atoms bonded to C).

18. Predict the order of the C-O bond lengths in CO, CO2, and CO32-.

19. Based on Lewis structures, predict the ordering of N-O bond lengths in NO+, NO2-, and NO3-.

20. (a) How can the concept of resonance be used to explain that all six C-C bonds in benzene are equal in length? (b) The C-C bond lengths in benzene are shorter than C-C single bonds but longer than C=C double bonds. Explain how this observation is consistent with the notion of resonance.

Exceptions to the Octet Rule

21. Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule and explain why they do not. (a) CO32-; (b) BH3; (c) I3-; (d) GeF4; (e) AsF6-.

Additional Exercises

22. Three Lewis structures can be drawn for N2O:    

   :NN-O: :N-NO: :N=N=O:


(a) Using formal charges, which of these three resonance forms is likely to be the most important? (b) The N-N bond length in N2O is 1.12 , slightly shorter than a typical N=O bond. (See Table 8.4) Rationalize these observations in terms of the resonance structures shown above and your conclusion for (a).

Molecular Shapes; the VSEPR Model

23. Describe the characteristic electron-domain geometry of each of the following numbers of electron domains about a central atom: (a) 3; (b) 4; (c) 5; (d) 6.

24. What is the difference between the electron-domain geometry and the molecular geometry of a molecule? Use the ammonia molecule as an example in your discussion.

25. Give the electron-domain and molecular geometries of molecules that have the following electron domains on their central atoms: (a) four bonding domains and no nonbonding domains; (b) three bonding domains and two nonbonding domains; (c) five bonding domains and one nonbonding domain.

26. Give the electron-domain and molecular geometries for: (a) PBr3; (b) CH3+; (c) BH4-; (d) SO3; (e) AsCl5; (f) BrF5.

27. The three species NH2-, NH3, and NH4+ have H-N-H bond angles of 105 , 107 , and 109 , respectively. Explain this variation in bond angles.

Polarity of Polyatomic Molecules

28. Consider the AF3 molecules in Figure 9.43. Which of these will have a nonzero dipole moment? Explain.

29. Will the following molecules be polar or nonpolar: (a) CCl4; (b) CS2; (c) SO3; (d) SF4; (e) NF3; (f) PF5?

Orbital Overlap; Hybrid Orbitals

30. Indicate the hybridization and bond angles associated with each of the following electron-domain geometries: (a) linear; (b) tetrahedral; (c) trigonal planar; (d) octahedral; (e) trigonal bipyramidal.

31. Indicate the hybrid orbital set used by the central atom in each of the following molecules and ions: (a) SiH4; (b) CH3+; (c) ICl2-; (d) MgCl2; (e) PF6-.

Multiple Bonds

32. (a) Sketch a bond that is constructed from p orbitals. (b) Sketch a bond that is constructed from p orbitals. (c) Which is generally the stronger, a bond and a bond? Explain.

33. (a) Draw Lewis structures for methane, CH4, and formaldehyde, H2CO. (b) What is the hybridization at the carbon atom in CH4 and H2CO? (c) The carbon atom in CH4 can not participate in multiple bonding whereas that in H2CO can. Explain this observation by using the hybridization at the carbon atom.

Molecular Orbitals

34. (a) What are the similarities and differences between atomic orbitals and molecular orbitals? (b) Why is the bonding molecular orbital of H2 at lower energy than the electron in a hydrogen atom? (c) How many electrons can be placed into each MO of a molecule?

35. According to molecular-orbital theory, would Be2 or Be2+ be expected to exist? Explain.

36. (a) What is meant by the term paramagnetism? (b) How can one determine experimentally whether a substance is paramagnetic? (c) Which of the following ions would you expect to be paramagnetic: O2+; N22-; Li2+; O22-?

Additional Exercises

37. Predict the molecular geometry of (a) PCl4+; (b) SO2; (c) HCN; (d) H2Te; (e) Br3-.

38. From their Lewis structures, determine the number of and bonds in each of the following molecules or ions: (a) CO2; (b) thiocyanate ion, NCS-; (c) formaldehyde, H2CO; (d) formic acid, HCO(OH), which has one H and two O atoms attached to C.