CHAPTER 7

PROBLEMS

    1.     Give the empirical formulas and names of the compounds formed from the following pairs of ions: (a) Rb+ and I-, (b) Cs+ and SO42-, (c) Sr2+ and N3-, (d) Al3+ and S2-.

    2.     List the following bonds in order of increasing ionic character: the lithium-to-fluorine bond in LiF, the potassium-to-oxygen bond in K2O, the nitrogen-to-nitrogen bond in N2, the sulfur-to-oxygen bond in SO2, the chlorine-to-fluorine bond in ClF3.

    3.     Classify the following bonds as ionic, polar covalent, or covalent, and give your reasons:  (a) the C-C bond in H3CCH3, (b) the K-I bond in KI, (c) the N-B bond in H3NBCl3, (d) the C-H bond in CH4.

    4.     Draw Lewis formulas for the following molecules:  (a) ICl, (b) PH3, (c) H2S, (d) N2H4, (e) HClO3, (f) COBr2 (C is bonded to O and Br atoms).  

    5.     The following Lewis formulas are incorrect.  Explain what is wrong with each one and give a correct Lewis structure for the molecule.  Note:  The relative positions of atoms are shown correctly.

              (a)

              (b)

              (c)

              (d)

              (e)

              (f)

              (g)

    6.     Draw Lewis formulas for HCO2-, including all resonance forms, and show formal charges.  The relative positions of the atoms are as follows:

             


    7.     Draw three reasonable resonance formulas for the OCN- ion. Show formal charges.

    8.     Of the noble gases, only Kr, Xe, and Rn are known to form a few compounds with O and/or F.  Draw Lewis formulas for the following molecules: (a) XeF2, (b) XeF4, (c) XeF6, (d) XeOF4, (e) XeO2F2.  In each case, Xe is the central atom.

    9.     Use Lewis formulas to show the transfer of electrons between the following atoms to form cations and anions:  (a) Na and F, (b) K and S, (c) Ba and O, (d) Al and N.

10.     Four atoms are arbitrarily labeled D, E, F, and G.  Their electronegativities are as follows:  D = 3.8, E = 3.3, F = 2.8, and G = 1.3.  If the atoms of these elements form the molecules DE, DG, EG, and DF, how would you arrange these molecules in order of increasing covalent bond character?

11.     Classify the following substances as ionic compounds or covalent compounds containing discrete molecules:  (a) CH4, (b) KF, (c) CO, (d) SiCl4, (e) BaCl2.

12.     Draw Lewis formulas for SeF4 and SeF6.  Is the octet rule satisfied for Se?

13.     Draw Lewis formulas for BrF3, ClF5, and IF7.  Identify those in which the octet rule is not obeyed.

14.     Draw Lewis formulas for the following four isoelectronic species:  (a) CO, (b) NO+, (c) CN-, (d) N2.  Show formal charges.

15.     Draw three resonance structures for the cyanate ion (NCO-) and rank the resonance structures in order of increasing importance.

16.     The amide group plays an important role in determining the structure of proteins:

              Draw another resonance structure for this group.  Show the formal charges.

17.     A rule for drawing plausible Lewis formulas is that the central atom is invariably less electronegative than the surrounding atoms.  Explain why this is so.

18.     Nitrous oxide (laughing gas) has the formula N2O.  If the relative position of the atoms is N  N  O, draw two Lewis formulas and explain which is preferred.

19.     Hydrogen cyanide can be drawn with two different Lewis formulas (HCN or HNC).  Explain which formula is preferred and explain why.

20.     Nitric acid (HNO3) can be thought of as being made from H+ and NO3-.  Explain the structure and bonding in nitric acid.

21.     An ionic bond is formed between a cation A+ and an anion B-.  How would the energy of the ionic bond be affected by the following changes?

              a.   doubling the radius of A+

              b.   tripling the charge on A+

              c.   doubling the charges on A+ and B-

              d.   decreasing the radii of A+ and B- to half their original values


22.     Predict the structure of each of the following molecules or ions (using VSEPR):

              a.   SiH4

              b.   TeF6

              c.   NH2Cl

              d.   SbF5

              e.   O3

              f.    CO32-

              g.   H3O+

              h.   PO43-

              i.     AlCl3

              j.     SiF62-

              k.   TeF4

              l.     BrF3

              m.  ICl4-

              n.   ICl2-

              o.   XeOF4

23.     Predict which of the structures in #1 would be polar.

24.     List the following molecules in order of increasing dipole moment:  H2O, CBr4, H2S, HF, NH3, CO2.

25.     Which of the following molecules has a higher dipole moment?

             

26.     Show the distribution of valence electrons in the orbitals of the central atom in each of the following molecules or ions just prior to bonding to the other atoms in the compound, and show the distribution after bonding takes place:

              a.   GeCl4

              b.   AlCl3

              c.   BrF3

              d.   BrF5

27.     The allene molecule H2C=C=CH2 is linear (the three C atoms lie on a straight line).  What are the hybridization states of the carbon atoms?  Draw diagrams to show the formation of sigma bonds and pi bonds in allene.


28.     How many sigma bonds and pi bonds are there in each of the following molecules?

             

29.     Indicate the type of hybridization for each central atom in #1.

30.     Determine the bond orders of the homonuclear diatomic molecules formed by the first 10 elements of the Periodic Table.

31.     Explain why the bond order of N2 is greater than that of N2+, but the bond order of O2 is less than that of O2+.

32.     Consider the diatomic ions X22+, where X is each one of the first ten elements of the Periodic Table.  Write the molecular orbital electronic configurations of these ions.  Determine the bond order of each and determine if the structure would be stable.

33.     Explain why the symbol on the right is a better representation of benzene molecules than that on the left.

                    

34.     Predict whether each of the following atoms, ions, or molecules should be paramagnetic or diamagnetic.

              a.   He

              b.   O

              c.   F2

              d.   NO

              e.   B

              f.    O2

              g.   Co

              h.   H2+

              i.     B2

              j.     F

              k.   CO

              l.     F2+

35.     Identify the homonuclear diatomic molecules or ions that have the following electronic configurations:

              a.   X2+         KK(s2s)1

              b.   X2:          KK(s2s)2(s2s*)2(p2py, p2pz)4

              c.   X2-:         KK(s2s)2(s2s*)2(p2py, p2pz)3

              d.   X2+:        KK(s2s)2(s2s*)2(s2px)2(p2py, p2pz)4(p2py *, p2pyz*)3

36.     The bond angle of SO2 is very close to 120°, even though there is a lone pair on S.  Explain.


37.     Draw a molecular orbital energy level diagram for HHe and calculate the bond order.

38.     Use molecular orbital theory to compare the relative stabilities of F2 and F2+.

39.     Predict the geometry of sulfur dichloride (SCl2) and predict the hybridization of the sulfur atom.

40.     Draw the Lewis formula for the BeCl42- ion.  Predict its geometry and describe the hybridization state of the Be atom.

41.     Cyclopropane (C3H6) has the shape of a triangle in which a C atom is bonded to two H atoms and two other C atoms at each corner.  Cubane (C8H8) has the shape of a cube in which a C atom is bonded to one H atom and three other C atoms at each corner.  (a) Draw the Lewis formulas of these molecules. (b) Compare the C-C-C bond angles in these molecules with those predicted for an sp3 hybridized C atom.  (c) Would you expect these molecules to be easy to make?  (Explain your answer.)