Describe the hybridization of each carbon atom in each of the following structures. Base

Draw a Lewis structure for each of the following molecules and assign charges where appropriate. The order in which the atoms are connected is given in parentheses. O (a) ClF (b) BrCN (c) SOCl2(ClSCl) (d) CH3NH2 (e) CH3OCH3 (f) N2H2 (HNNH) (g) CH2CO (h) HN3 (HNNN) (i) N2O (NNO)

Using electronegativity values from Table 1-2 (in Section 1-3), identify polar covalent bonds in several of the structures in Problem 25 and label the atoms d1 and d2, as appropriate.

Draw a Lewis structure for each of the following species. Again, assign charges where appropriate. (a) H2 (b) CH3 2 (c) CH3 1 (d) CH3 (e) CH3NH3 1 (f) CH3O2 (g) CH2 (h) HC2 2 (HCC) (i) H2O2 (HOOH)

For each of the following species, add charges wherever required to give a complete, correct Lewis structure. All bonds and nonbonded valence electrons are shown. (a) H O H H H H C (b) C O H H H (c) H H H C (d) H H H N O H (e) O H H O O B (f) N O H H

(a) The structure of the bicarbonate (hydrogen carbonate) ion, HCO3 2, is best described as a hybrid of several contributing resonance forms, two of which are shown here. O O   O O  HO   HO (i) Draw at least one additional resonance form. (ii) Using curved electron-pushing arrows, show how these Lewis structures may be interconverted by movement of electron pairs. (iii) Determine which form or forms will be the major contributor(s) to the real structure of bicarbonate, explaining your answer on the basis of the criteria in Section 1-5. (b) Draw two resonance forms for formaldehyde oxime, H2CNOH. As in parts (ii) and (iii) of (a), use curved arrows to interconvert the resonance forms and determine which form is the major contributor. (c) Repeat the exercises in (b) for formaldehyde oximate ion, [H2CNO]2.

Several of the compounds in Problems 25 and 28 can have resonance forms. Identify these molecules and write an additional resonance Lewis structure for each. Use electron-pushing arrows to illustrate how the resonance forms for each species are derived from one another, and in each case indicate the major contributor to the resonance hybrid.

Draw two or three resonance forms for each of the following species. Indicate the major contributor or contributors to the hybrid in each case. O (a) OCN2 (b) CH2CHNH2 (c) HCONH2 (HCNH2) (d) O3 (OOO) (e) CH2CHCH2 2 (f ) ClO2 2 (OClO) (g) HOCHNH2 1 (h) CH3CNO

Compare and contrast the Lewis structures of nitromethane, CH3NO2, and methyl nitrite, CH3ONO. Write at least two resonance forms for each molecule. Based on your examination of the resonance forms, what can you say about the polarity and bond order of the two NO bonds in each substance?

Write a Lewis structure for each substance. Within each group, compare (i) number of electrons, (ii) charges on atoms, if any, (iii) nature of all bonds, and (iv) geometry. (a) chlorine atom, Cl, and chloride ion, Cl2 (b) borane, BH3, and phosphine, PH3 (c) CF4 and BrF4 2 (C and Br are in the middle) (d) nitrogen dioxide, NO2, and nitrite ion, NO2 2 (nitrogen is in the middle) (e) NO2, SO2, and ClO2 (N, S, and Cl are in the middle)

Use a molecular-orbital analysis to predict which species in each of the following pairs has the stronger bonding between atoms. (Hint: Refer to Figure 1-12.) (a) H2 or H2 1 (b) He2 or He2 1 (c) O2 or O2 1 (d) N2 or N2 1

For each molecule below, predict the approximate geometry about each indicated atom. Give the hybridization that explains each geometry. (a) Br Br H2C CH2 (b) C O H3C CH3 (c) H3C O CH CH2 (d) H3C NH2 (e) HCqCOCH2OOH (f) H2C NH2

For each molecule in Problem 35, describe the orbitals that are used to form every bond to each of the indicated atoms (atomic s, p, hybrid sp, sp2, or sp3).

Draw and show the overlap of the orbitals involved in the bonds discussed in Problem 36.

Describe the hybridization of each carbon atom in each of the following structures. Base your answer on the geometry about the carbon atom. P r o b l e m s (a) CH3Cl (b) CH3OH (c) CH3CH2CH3 (d) CH2CH2 (trigonal carbons) (e) HCCH (linear structure) (f) C H3C O H (g) C O C H2C H2C O H H

Depict the following condensed formulas in Kekul (straight-line) notation. (See also Problem 42.) (a) CH3CN (b) (CH3)2CHCHCOH H2N O (c) CH3CHCH2CH3 OH (d) CH2BrCHBr2 (e) O CH3CCH2COCH3 O (f ) HOCH2CH2OCH2CH2OH

Convert the following bond-line formulas into Kekul (straight-line) structures. (a) OH (b) O N (c) Br Br (d) (e) O CN (f) S

Convert the following hashed-wedged line formulas into condensed formulas. (a) H2N NH2 C C H H H H } C i i [ & O (b)  C x N ~ i i [ & C C O H H H H H H H i }C & O O  (c) H ~ i & Br Br Br

Depict the following Kekul (straight-line) formulas in their condensed forms. (a) C N C H H H H H H H (b) C N C H O H H H H H H C H C H (c)  C H H H H C H H C H H C O S H H D (d) C C  H H H O F F F (e) H C H H H H H H H H H C C C C (f ) C H C C H H H H H C O

Redraw the structures depicted in Problems 39 and 42 using bond-line formulas.

Convert the following condensed formulas into hashed-wedged line structures. (a) C N CH3 C HOCH3 (b) CHCl3 (c) (CH3)2NH (d) SH CH3CHCH2CH3

Construct as many constitutional isomers of each molecular formula as you can for (a) C5H12; (b) C3H8O. Draw both condensed and bond-line formulas for each isomer.

Draw condensed formulas showing the multiple bonds, charges, and lone electron pairs (if any) for each molecule in the following pairs of constitutional isomers. (Hint: First make sure that you can draw a proper Lewis structure for each molecule.) Do any of these pairs consist of resonance forms? O (a) HCCCH3 and H2CCCH2 (b) CH3CN and CH3NC (c) CH3CH and H2CCHOH

Two resonance forms can be written for a bond between trivalent boron and an atom with a lone pair of electrons. (a) Formulate them for (i) (CH3)2BN(CH3)2; (ii) (CH3)2BOCH3; (iii) (CH3)2BF. (b) Using the guidelines in Section 1-5, determine which form in each pair of resonance forms is more important. (c) How do the electronegativity differences between N, O, and F affect the relative importance of the resonance forms in each case? (d) Predict the hybridization of N in (i) and O in (ii).

The unusual molecule [2.2.2]propellane is pictured in the margin. On the basis of the given structural parameters, what hybridization scheme best describes the carbons marked by asterisks? (Make a model to help you visualize its shape.) What types of orbitals are used in the bond between them? Would you expect this bond to be stronger or weaker than an ordinary carbon carbon single bond (which is usually 1.54 long)?

(a) On the basis of the information in Problem 38, give the likely hybridization of the orbital that contains the unshared pair of electrons (responsible for the negative charge) in each of the following species: CH3CH2 2; CH2CH2; HCC2. (b) Electrons in sp, sp2, and sp3 orbitals do not have identical energies. Because the 2s orbital is lower in energy than a 2p, the more s character a hybrid orbital has, the lower its energy will be. Therefore the sp3 (14 s and 34 p in character) is highest in energy, and the sp (12 s, 12 p) lowest. Use this information to determine the relative abilities of the three anions in (a) to accommodate the negative charge. (c) The strength of an acid HA is related to the ability of its conjugate base A2 to accommodate negative charge. In other words, the ionization HA H1 1 A2 is favored for a more stable A2. Although CH3CH3, CH2CH2, and HCCH are all weak acids, they are not equally so. On the basis of your answer to (b), rank them in order of acid strength.

A number of substances containing positively polarized carbon atoms have been labeled as cancer suspect agents (i.e., suspected carcinogens or cancer-inducing compounds). It has been suggested that the presence of such carbon atoms is responsible for the carcinogenic properties of these molecules. Assuming that the extent of polarization is proportional to carcinogenic potential, how would you rank the following compounds with regard to cancercausing potency? (a) CH3Cl (b) (CH3)4Si (c) ClCH2OCH2Cl (d) CH3OCH2Cl (e) (CH3)3C1 (Note: Polarization is only one of the many factors known to be related to carcinogenicity. Moreover, none of them shows the type of straightforward correlation implied in this question.)

Certain compounds, such as the one pictured below, show strong biological activity against cell types characteristic of prostate cancer. In this structure, locate an example of each of the following types of atoms or bonds: (a) a highly polarized covalent single bond; (b) a highly polarized covalent double bond; (c) a nearly nonpolar covalent bond; (d) an sp-hybridized carbon atom; (e) an sp2-hybridized carbon atom; (f ) an sp3-hybridized carbon atom; (g) a bond between atoms of different hybridization; (h) the longest bond in the molecule; (i) the shortest bond in the molecule (excluding bonds to hydrogen). C O N O H3 CH3 C Cl H H 3 Team Problems Team problems are meant to encourage discussion and collaborative learning. Try to solve team problems with a partner or small study group. Notice that the problems are divided into parts. Rather than tackling each part individually, discuss each section of the problem together. Try out the vocabulary that you learned in the chapter to question and convince yourselves that you are on the right track before you move to the next part. In general, the more you use the terms and apply the concepts presented in the text, the better you will become at correlating molecular structure and reactivity and thus visualizing bond breaking and bond making. You will begin to see the elegant patterns of organic chemistry and will not be a slave to memorization. The collaborative process used in partner or group study will force you to articulate your ideas. Talking out a solution with an audience instead of to yourself builds in checks and balances. Your teammates will not let you get away with, Well, you know what I mean, because they probably do not. You become responsible to others as well as yourself. By learning from and teaching others, you solidify your own understanding.

Consider the following reaction: CH3CH2CH2 CH3 O CH3CH2CH2CCH3 HCN C HO CN A B P r o b l e m s (a) Draw these condensed formulas as Lewis dot structures. Label the geometry and hybridization of the bold carbons in compounds A and B. Did the hybridization change in the course of the reaction? (b) Draw the condensed formulas as bond-line structures. (c) Examine the components of the reaction in light of bond polarity. Using the notation for partial charge separation, d1 and d2, indicate, on the bond-line structures, any polar bonds. (d) This reaction is actually a two-step process: cyanide attack followed by protonation. Depict these processes by using the same electron pair pushing technique that we employed for resonance structures in Section 1-5, but now show the fl ow of electrons for the two steps. Clearly position the beginning (an electron pair) and the end (a positively polarized or charged nucleus) of your arrows. Preprofessional Problems Preprofessional problems are included to give you practice solving the type of problems found on exams required for entry into professional schools, such as the MCAT, DAT, chemistry GRE, and ACS exams, as well as on many undergraduate tests. Do these multiple-choice questions as you go through this course and then return to them before you take a professional school exam. These questions are to be answered closed book that is, no periodic table, calculator, or the like.

A certain organic compound was found on combustion analysis to contain 84% carbon and 16% hydrogen (C 5 12.0, H 5 1.00). A molecular formula for the compound could be (a) CH4O (b) C6H14O2 (c) C7H16 (d) C6H10 (e) C14H22

The compound A A O O O Br A Br A CH3 CH3 Br Al N CH2CH3 has a formal charge of (a) 21 on N (b) 12 on N (c) 21 on Al (d) 11 on Br (e) none of the above

The arrow in the structure in the margin points to a bond that is formed by (a) overlap of an s orbital on H and an sp2 orbital on C (b) overlap of an s orbital on H and an sp orbital on C (c) overlap of an s orbital on H and an sp3 orbital on C (d) none of the above

Which compound has bond angles nearest to 1208? (a) OCS (b) CHI3 (c) H2CO (d) HOCCOH (e) CH4

The pair of structures that are resonance hybrids is (a) P HO CHCH3 and  O HO CHCH3    (b) CH2 and A CH CHN K CH2 (c) CH3 H A C and OB CH2 H H C O K (d) CH3CH2 and  CH2CH3