OC 1.1-2.PDF

Chapt*r 1: Chemical Bontling -esson 1.1 Covalent Bonds Introduction Organic chemistry is broadly defined as the chemistry of carbon-containing compounds. Because the sludy of organic chemistry builds on the foundations covered in general chemistry coursework, Chapters ' and 2 serve as a...

Chapt*r 1: Chemical Bontling -esson 1.1 Covalent Bonds Introduction Organic chemistry is broadly defined as the chemistry of carbon-containing compounds. Because the sludy of organic chemistry builds on the foundations covered in general chemistry coursework, Chapters ' and 2 serve as a platform to bridge selected concepts from general chemistry to applications in organic :remistry. 1.'1.01 Valence Electrons and the Octet Rule dividual atoms will either transfer or share valence electrons to give each atom access to eight ,alence electrons, a feature commonly known as the octet rule. By fulfilling an octet, individual atoms :ecome rsoeiectronic with the noble gases and gain increased stability (Table 1.1). Table 1.1 The octet rule for elements in the first (n = 1) and second (n = 2) rows of the periodic table. n=1 n=2 Valence Needed Valence Needed Group Element. electrons for octeta Element electrons for octet 1A H 1 1 Li 1 7 2A Be 2 6 3A B 3 5 4A c 4 4 5A N 5 3 6A o 6 2 7A F 7 4 8A He Ne I 0 t First-row elements strive for a total of 2 valence electrons l€ments in the first row of the periodic table (ie, H, He) require only two valence electrons (the maximum -lmber allowed by a lone 1s orbital). Elements in the third row of the periodic table (or greater) have access to d and f subshells that may allow more than eight valence electrons (an expanded octet). -1owever, within organic chemistry most of the relevant elements appear in the first and second row. 1 i.A2 Formation of Chemical Bonds -here are two primary ways that atoms gain valence electrons, complete their octet, and form a chemical bond. r an ionic bond, one atom fullytransfers one or more electrons to another, and ions are generated. -he atom that gives up valence electrons becomes positively charged (a cation), while the atom that acrepts valence electrons becomes negatively charged (an anion). Salts form when the oppositely :.rarged ions become electrostatically attracted. Most ionic bonds are formed between a metal and a  ChaPt*r "l: Chemical Bonding they nonmetal, as shown in Figure 1.1. Consequently, ionic bonds are relevant to organic chemistry, but appear less frequently than covalent bonds. +..O :ClI --------) @:'.o Na" +.Ct, -----* Na@ Na :Cl: 1 Valen*e *leclr*n 7 Valence electrons Sodium cation crrnrioe anion Sodium.iiorio" (a salt) Fulln=2shell lsoelectronic lsoelectronic with Ne with Ar Figure 1.1 Formation of an ionic bond. 1'2' A covalent bond is formed when atoms share electrons to achieve a full octet, as shown in Figure Covalent bond t9r' +.Gli ---------> Chlorine atom Chlorine atom 7 Valence electrons 7 Valence electrons Figure 1.2 Formation of a covalent bond. 1.1.03 Lewis Structures each A Lewis (dot) structure represents chemical bonding between atoms in a molecule by depicting a pair of dots or, valence etection as a single dot. Shared bonding electrons are represented by either more commonly in org"nL chemistry, a single line between the atoms (a single bond). Nonbonding ef ectrons (ie, lone pairs) and are represented by dots around (but not between) atoms. Lone pairs are very importanf in determining the reactivity of organic compounds' full octet' A Sometimes atoms must share more than one pair of electrons for all atoms to achieve a and a triple bond forms when three pait double bond forms when two pairs (four electrons) are shared, (six electrons) are shared between atoms. Double and triple bonds are examples of multiple bonds, which are almost always represented by lines rather than dots (Figure 1.3). A multiplebond Nonbonding electrons (double bond) (lone Pairs) H :O: H :O: :: I lt H: C: C:9: H = H-C-C-O-H t" H H Allvalenceelectrons Onlynonbondingelectrons represented by a dot represented by a dot Figure 1.3 Multiple bonds and nonbonding electrons in Lewis structures.  *hnpi*r 1: Chemical i$*nrllr:g ' 1.04 Formal Charge =crmal charge quantitatively indicates the electron density of an atom in a molecule relative to the n?::'on density of the atom in isolation according to the equation: Formal charge : Atomic valence electrons - Nonbonding electrons - I lnonaing electrons) ::-al charges often influence the reactivity and properties of a substance, and the sum of the formal :::'jes of all atoms in a molecule is always equal to the overall charge of the molecule or ion (the *olecular net charge). A general chemistry text can provide a review of the process for generating -=r, s structures and assigning formal charges for i"rFlyillilnrirr lrl{.}i[)()tiir];c. ffimnnept ffi{tws:${'*.1 ,', hat is the formal charge for each nonhydrogen atom in the following Lewis structure? H H:O: llll.. H-N-C-C-O: tl HH So luticn : ',:ie: The appendix contains the answer. :,---:tgh the formal charge equation is useful, a recognition of how common bonding patterns affect ":-. al charge is a great time-saver. Often, the number of valence electrons of an element indicates the -::: ^umber of covalent bonds that the element prefers to form within molecules. However, there may be - -: rle different patterns of single, double, and triple bonds that result in the same formal charge. ::' example, the element carbon has four valence electrons and prefers to form a total of four covalent :':-:s with other elements to fulfill its octet and have a formal charge of 0. Although there are several :,:-: combinations (Figure 1.4)that meet this criterion, these variations create carbons with very different ,'ity profiles. Deviations from these patterns result in a nonzero formal charge. =::: I lt -c- I -c- Four single bonds One double bond Two double bonds One triple bond Two single bonds One single bond =€;re 1.4 Bonding scenarios for a carbon atom with a total of four covalent bonds. *-;.nost common bonding patterns for the second-row elements relevant to organic chemistry are :*: :ied in Figure 1.5. Committing these patterns to memory and understanding the relationship :€:/.een variants are likely to be very helpful throughout organic chemistry.  Chapter 1: Chemical Bondtng c N o X Element Carbon Nitrogen Oxygen Halogens Valence electrons 4 5 6 7 +1 Formal charge (cation) Total number of bonds 3 4 3 2 Number of lone pairs 0 0 1 2 @ lo..(o _1@ Example -c- I -N: I -oY-I No formal charge (neutral) Total number of bonds 4 3 2 1 Number of lone pairs 0. I 2 3 Example -i-I I -trj- I -i)- -X: -1 Formal charge (anion) Total number of bonds 3 2 1 0 Nurnber of lone Pairs 1 2 3 4 aO Example I -Nv -o: Figure 1.5 Relationship between valence electrons, covalent bonding, and formal charge. The distance between nuclei connected by a covalent bond (bond length) represents the optimal, lowes energy distance with the least repulsion and greatest constructive interaction (Figure 1.6). lf the nuclei are too close together, their electron clouds will strongly repel one another (ie, they will have high potential energyy, and if the nuclei are too far apart, their valence electrons will be unable to interact'  Chapter 1: Chemical Bonding Dissociation energy Internuclear distance Point D Point C Point B Point A @ W#W ry-1--:-'-@ Tm close Optimaldistance Farther apart. Separated Compressed bond Bond formed Met constiuctive. No interactions ',er repulsion Forces cancel interaction =gr,re 1.6 Potential energy diagram for the distance between nuclei in a covalent bond, rL:'--s ,'vith valence electrons in higher shells farther from the nucleus have longer bond lengths. For :!:-3 e. the bond iengths of the hydrohalic acids (H-X) increase from fluorine to iodine. 3:n,: engths also vary due to the multiplicity (bond order) of the covalent bond. Multiple bonds (eg, r,-c € and triple bonds) feature additional orbital interactions that bring the nuclei closer together and :=.:se the bond length (see Concept 1.4.04). For example, a carbon-carbon single bond has a larger lrr-'l sngth than a carbon-carbon double bond, which is larger than the bond length of a carbon-carbon *. - ^^^/ --e bond dissociation energy is the energy required to break a covalent bond and evenly divide the :e-*;:ls in the covalent bond (Figure 1.7). nergy A:B !* A.+.B =g-re 1.7 Bond dissociation energy. l":r-se:uently, the bond dissociation energy for a multiple bond is larger than that of a comparable single rnr,: -s-*rre more electrons must be disrupted (Figure 1.8). C-C C=C C=C Increasing bond strength, increasing bond dissociation energy a5Lre 1.8 Comparison of bond dissociation energies by bond order.  Chapter 1: f;hemical $onding Hybridization &nic orbitals are represented by mathematical expressions, which describe the behavior of electrons n r ;::in. The principal quantum number n (electron shell) describes the distance of an atomic orbital rEr :-. ^ucleus. The angular momentum quantum number /describes an atomic orbital's shape. -.t-e-:-- numbers are discussed in greater detail in General Chemistry Lesson 1.3. This lesson gives a ''F,'r"rrr :' s and p atomic orbitals and examines different types of associated hybrid orbitals. n :r::- : chemistry,the s and p atomic orbitals are moql prevalent. An s orbital (/= 0) has a spherical ;*€cE electron density distributed evenly around the nucleus; a p orbital (t = 1) has a dumbbell ,'r'1h i*a€ ,','ih two lobes that have opposite phases (ie, opposite signs [+ or -]within the mathematical xt:.s:n) and are connected by a nad* (ie, a point where the mathematicalexpression changes phase rc :-€'3 rs zero probability of electron density) (Figure 1.9). s orbitals n ='1 n=2 n.3 l=0 t=0 t=0 p orbitals z A r, \t, l, \----.,.rY \.f.r"'ol t' -'o' -tt \ - \x '/, \X '' i n*2 n=2 n=2 l*1 l=1 l=1 :r$rE 1.9 Shapes of s and p atomic orbitals. tms :.;anic molecules contain bonds made from hybrid orbitals, which are generated by the linear rrFc'..3: cn of s and p atomic orbitals on the same atom, yielding sp3, sp2, and sp hybrid orbitals. $er= - of the conservation of orbitals, the number of hybrid orbitals generated must be equal to the llulT-c* oi atomic orbitals that were hybridized. Each hybrid orbital has a distinct electron domain IE-€:".. that minimizes orbital repulsion and dictates the bond angles in the molecule (see Lesson 2.1). ' 2 :2 sp3 Orbitals -re ::'-cination of one s orbital and three p orbitals (four total atomic orbitals) results in four sp3 hybrid :rotds Figure '1.10), each having 25o/o s character and75% p character. 11  Chapter 1: Chemical Bon 4 atomic orbitals I C- l'] 4 hybrid orbitals tr ffi-ffi,ffi LW], ffiffff ffi 1;iri*5':35"' ffiffi.\l.w,*w, *#ffi y.y.y.y )tsv"pcharacter 2s ZPx ' Zpy # 2pz sp3 orbitals Figure 1.10 Formation of sp3 hybrid orbitals' counting the by examininglhe Lewis structure and The hybridization of an atom can be determined around the ato number of electron domains (ie, bonding ;g;tt'";J pairs 6t nonbonding electrons) domains must ec ete"ctron domain' The number of electron Each double or triple oonO.ount, as one hybrid orbital' the sum of the s and p superscripts in"the Atomsthataresp3hybridizedhavefourelectrondomains.Forexamp|e,thecarboninmethano|(CH more o*vfen (Figure 1.11) (see concept 2'1'02 for has three bonds to hydrogen and one oonolo information). i ol'l ,l *lI.rr.:it r.rtl ijlil ri!iilrit l,uH ''i H'"tH ,, sp3 hybridized Methanol I sp3 hybridization' Figure 1.11 Electron domains and two p atomic orbitars (three totar atomic orb*ars) form tl The combination of one s atomic orbitar and "E;h p ,no*n in Figure 1. 1 t. ;rbital has 33% s character and 670/o charact sp2 hybrid orbitats p"tp"noi.ur"r to the plane of the three sf orbitals; unhyb ", unhybridized p orbital remains oriented p orbitals can be uu."ii,'coniain a lone puii:Siputticipate in a pi bond (see Concept 1'4'03)' 3 atomic orbitals f r,'\ 3 hYbrid orbitals I sl orbitals have -W-W lu] ilff-U Hybridizatiofi ,* I sSZ" s character and ur* p character. v ) #* \-. ZPx zPv sdorbitals Figure 1'12 Formation of sPz hYbrid orbitals' 12  Chapter 1: Chemical Bonding r€rs that participate in double bonds are a common example of sl hybridization. The carbon atoms in s€r€ ;CHzCHz) each have two single bonds to hydrogen and one double bond to the other carbon; f^e=,':re. each carbon atom has three electron domains and is sp2 hybridized (Figure 1.13). HH ts*^' ,:: i.t. a.:: liii"'rl sp2 hybridized HH Ethene qgllrr.€ 1.13 Electron domains and sp2 hybridization. ::r-e:cn orbitals occurs by mixing one s atomic orbital and one p atomic orbital, of two sp hybrid GBr'r-'J :wo unhybridized p orbitals (Figure 1.14). Because only one p orbital is used to form the two sp mcrc :rbitals, the hybrid orbitals must be oriented along only the axis of the contributing p orbital. As s.r sr l-rybrid orbitals have a linear relationship. Each sp orbital has 50% s character and 50% p 1a 3_\= - Each blue lobe is an sp lobe i (the yellow lobes are omitted). : i- I I- sp orbitals have 50% s character and p character. x )so"t" 2s \J2p, sp orlcitals trrr 1.14 Formation of sp hybrid orbitals. rcrs :aticipating in one triple bond or two double bonds are sp hybridized. The carbon atoms in $"-€ 3:Hz) each have one bond to hydrogen and a triple bond to the other carbon; therefore, these :mcrs 'ave two electron domains (Figure 1.15). H*C=C*H sp hybridized Ethyne trc 1,15 Electron domains and sp hybridization. 13  Chapter 1: *h*mieal ffione;ling 'il.? ffi fi'nurceptu" Sl^rr*nil' ldentify the hybridization of the labeled atoms in the molecule in the following structure. I t' ryH (t\ - {r,,*t-i-"*;- t--/"'H l\==.-F ii -C -,^- N /'4'--/ " Ho ') cY cI Ato ilI ^ ' ll "+- 3 "$*lutielr"t Note: The appendix contains the answer..