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) Se^2-; (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 O^2-. Why, then, is CaO stable relative to the free elements?

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^-, S^2-, K⁺; (d) Cl, Cl^-, Ar.

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?

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) CO₃^2-; (c) HCO₂^- (H and both O atoms bonded to C).

18. Predict the order of the C-O bond lengths in CO, CO₂, and CO₃^2-.

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.

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) CO₃^2-; (b) BH₃; (c) I₃^-; (d) GeF₄; (e) AsF₆^-.

22. Three Lewis structures can be drawn for N₂O:    

   :NÉN-O: ¢ :N-NÉO: ¢ :N=N=O:

[NOT ALL UNSHARED ELECTRONS SHOWN] 

(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).

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.

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?

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^-.

32. (a) Sketch a ¦ bond that is constructed from p orbitals. (b) Sketch a ¹ bond that is constructed from p orbitals. (c) Which is generally the stronger, a ¦ bond and a ¹ 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.

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⁺; N₂^2-; Li₂⁺; O₂^2-?

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 ¹ 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.