Prof. J. T. Spencer, CHE 106 1 Chapter Eight Copyright © James T. Spencer 1995 - 1998 All Rights Reserved Prof. J. T. Spencer, CHE 106 2 Chapter Eight • Ions versus Molecules. • Ionic Bonding - Ions which are held together primarily by electrostatic forces (between oppositely charged ions). • Covalent Bonding - sharing of electrons between atoms in molecules. • Metallic Bonding - sharing of electrons between a very large array of atoms in which the electron are relatively free to move throughout the large array. Prof. J. T. Spencer, CHE 106 3 Lewis Structures • Valence electrons - electrons which participate in bonding (outer shell electrons, usually from an unfilled shell). • Electron Bookkeeping - Lewis Symbols (or more commonly Lewis Structures) keep track of valence electrons involved in bonding (ignoring core electrons). Prof. J. T. Spencer, CHE 106 4 Lewis Structures •Lewis symbols/structures consist of; •Elemental Symbol •a “dot” for each valence electron •form compounds by sharing or exchanging electrons to achieve nearest Noble Gas (group 18) configuration. Since all noble gases have 8 valence electrons, many atoms react to have eight electrons (octet rule). Prof. J. T. Spencer, CHE 106 Lewis Structures 5 . . . . . . . . . . . . . . .. . . Li . Be . B . . C N O . . . . : .F. : Ne .. : Element Li Be B C N O F Ne Elec. Config. Valence Electrons 1s22s1 1 1s22s2 2 1s22s22p1 3 1s22s22p2 4 1s22s22p3 5 1s22s22p4 6 1s22s22p5 7 1s22s22p6 8 Prof. J. T. Spencer, CHE 106 Lewis Structures 6 • Octet Rule - elements tend to surround themselves with eight valence electrons through electron sharing or exchange (to fill shell). Many exceptions, however. Electron Exchange - Ionic Bonding . Na .. . Cl : .. . . -1 . Na+ + . Cl . .: Electron Sharing - Covalent Bonding . .. : .F. .. . .F.: .. .. : .F. : .F.: Prof. J. T. Spencer, CHE 106 Ionic Bonding 7 Na(g) + Cl(g) Na+1(g) + Cl-1(g) E° = I1(Na) - EA(Cl) = +146.8 kJ mol-1 Na r = radius Cl Na+1(g) + Cl-1(g) E n e r g y Na0(g) + Cl0(g) harpooning r = radius Prof. J. T. Spencer, CHE 106 Ionic Bonding 8 • Formation of gaseous ions is Endothermic. Why does NaCl form if ion formation if endothermic? Na + Cl Na+1 + Cl-1 (g) (g) (g) (g) E° = I1(Na) - EA(Cl) = +146.8 kJ mol-1 • Answer: Lattice Energy - Defined - energy required to completely separate a solid ionic compound into its gaseous ions. The reverse (“condensation” to solid is very exothermic) Na+1(g) + Cl-1(g) NaCl(s) E° = Lattice Energy = -788 kJ mol-1 Prof. J. T. Spencer, CHE 106 9 Cmpd LiF LiI NaF NaCl NaI KF KBr KI MgF2 SrCl2 MgO Lattice Energy Lattice Energy (kJ mol-1) 1024 744 911 788 693 815 682 641 2910 2139 3938 • Endothermic written from solid to gas and exothermic from gas to solid. • Higher charged ions have much greater lattice energies • The closer the ions can come together, the larger the lattice energy. E=k k = const. Q1Q2 Q1 = charge ion 1 Q2 = charge ion 2 d d = distance Prof. J. T. Spencer, CHE 106 Ionic Compounds 10 + - + + + - - + - - + + - - + + + - + - + - + - + - + - - - + Unit Cell + - Cell Face Prof. J. T. Spencer, CHE 106 11 Ionic Compounds TiO2 Surface Prof. J. T. Spencer, CHE 106 12 Ionic Compounds LiF Lattice Prof. J. T. Spencer, CHE 106 13 Ionic Compounds • Why do ionic compounds form? • Which ionic compounds are stable? • How do we determine the stoichiometry of the stable forms –i.e., MgCl or MgCl2 or MgCl3 ? Thermodynamics Use Hess’ Law and Born-Haber Cycles Prof. J. T. Spencer, CHE 106 14 Born-Haber Cycle Law - if a reaction is carried out in a series of steps, H for the reaction will be equal to the sum of the H’s for the individual steps. Hess’s Prof. J. T. Spencer, CHE 106 15 Born-Haber Cycle Law - if a reaction is carried out in a series of steps, H for the reaction will be equal to the sum of the H’s for the individual steps. +1 -1 Hess’s E n e r g y Na (g) + Cl (g) Na(g) + Cl(g) Na(s) + 1/2 Cl2(g) HRx NaCl(s) Prof. J. T. Spencer, CHE 106 16 Born-Haber Cycle Law - if a reaction is carried out in a series of steps, H for the reaction will be equal to the sum of the H’s for the individual steps. +1 -1 Hess’s Na (g) + E n e r g y EA I1 Na(g) + Hsub Cl (g) - Lattice Energy (U) Cl(g) U cannot be determined directly BDE Na(s) + 1/2 Cl2(g) HRxn = Hsub + I1 + BDE + EA + U HRx NaCl(s) Prof. J. T. Spencer, CHE 106 17 Born-Haber Cycle Exothermic HRx = Hvap + I1 + BDE + EA + U NaCl Endothermic Prof. J. T. Spencer, CHE 106 18 Born-Haber Cycle Why NaCl and not NaCl2 Na+2(g) + 2 Cl-1(g) Na+1(g) + 2 Cl-1(g) E n e r g y I1 Na(g) + 2 Cl(g) NaCl2(s) HRx Na(s) + Cl2(g) Prof. J. T. Spencer, CHE 106 19 Born-Haber Cycle Why NaCl and not NaCl2 Na+2(g) + 2 Cl-1(g) Very Endothermic I2 - Lattice Energy (U) Na+1(g) + 2 Cl-1(g) E n e r g y I1 EA Na(g) + Hvap Na(s) + 2 Cl(g) BDE NaCl2(s) HRx Cl2(g) HRx = Hvap + I1 + I2 + BDE + EA + U Prof. J. T. Spencer, CHE 106 20 NaCl2 Born-Haber Cycle Very Endothermic HRx = Hvap + I1 + I2 + BDE + EA + U Endothermic Exothermic NaCl HRx = Hvap + I1 + BDE + EA + U Prof. J. T. Spencer, CHE 106 21 Born-Haber Cycle larger U offsets increased endothermicity from I2 and EA2 E n e r g y Hf° for MgO Mg+2(g) + O-2(g) Mg+1(g) + O-1(g) Mg(g) + O(g) HRx Mg(s) + 1/ 2 O2(g) MgO(s) Prof. J. T. Spencer, CHE 106 22 Born-Haber Cycle larger U offsets increased endothermicity from I2 and EA2 Hf° for MgO Mg+2(g) + O-2(g) I2 Mg+1(g) + EA O-1(g) I1 E n e r g y - Lattice Energy (U) EA Mg(g) + O(g) Hvap Mg(s) + BDE 1/ 2 O2(g) HRx MgO(s) HRx = Hvap + I1 + I2 + BDE + EA1 + EA2 + U Prof. J. T. Spencer, CHE 106 23 Electronic Config. of Ions When a positive ion is formed, electrons are always lost first from the subshell with the largest n value. To determinectronic configurations of Cations, remove the highest energy electrons (usually the last one added EXCEPT for transition metal ions - which have NO n(max)s electrons). 25Mn 25Mn+1 4s 3d 3d 3d 3d 3d 4p 4p 4p Prof. J. T. Spencer, CHE 106 24 Ionic Radii • Sizes of ions determines stability and packing arrangements (structures) of ionic compounds. • Size depends upon Zeff and the total number of electrons of the ion. – Cations are smaller than the neutral atoms – Anions are larger than the neutral atoms – size of Ions of the same charge increases down and L. • Isoelectronic Series - ions with the same number of electrons (although different charges): N-3 = O-2 = F-1 = Ne = Na+1 = Mg+2 = Al+3 = 10 elec. Prof. J. T. Spencer, CHE 106 Ionic Radii 25 r a d i u s • Isoelectronic series - ions with the same number of electrons Se-2 Br -2 S ClO-2 F- K+ Na+ Li+ Mg+2 Be+2 Al+3 Ca+2 Sc+3 Ti+4 Rb+ Sr+2 Y+3 +4 Zr smaller w/ incr + charge smaller with smaller max n atomic number Prof. J. T. Spencer, CHE 106 26 Covalent Bonding • Atoms may achieve noble gas configurations by sharing electrons . .. : .F. .. . .F.: . 4 H . + . C. . .. .. : .F. : .F.: H .. .. H H .. C .. H Each atom achieves an outer shell with 8 electrons (at least) Lewis structures show this electron sharing Prof. J. T. Spencer, CHE 106 27 Covalent Bonding • Wave function (wave properties) • Can “constructively” and “destructively add waves (just like ripples on a pond). 1s orbital 2 0 radius H atom 1 Move together to overlap waves 0 radius 1s orbitals 2 H atom 2 0 radius Prof. J. T. Spencer, CHE 106 28 Covalent Bonding Destructive Addition of Waves (out of phase) 2 H atom 2 H atom 1 0 Constructive Addition of Waves (in phase) radius 0 2 H atom 2 H atom 1 0 radius 0 Prof. J. T. Spencer, CHE 106 29 Covalent Bonding s orbital - in-phase addition (bonding) + s orbital - out-of-phase addition (antibonding) + Prof. J. T. Spencer, CHE 106 30 Covalent Bonding p orbital - in-phase addition (bonding) + p orbital - out-of-phase addition (antibonding) + Prof. J. T. Spencer, CHE 106 31 Multiple Bonds • Bond Order (bo) denotes how many electron pairs are shared between two atoms; – single bond (bo = 1) has one shared electron pair. longest – double bond (bo = 2) has two shared electron pair. – triple bond (bo = 3) has three shared electron pair. shortest Single .. .. : .F. : .F.: F = 1s22s22p5 Double .. .. :: O .. O .. O = 1s22s22p4 Triple : N ::: N: N = 1s22s22p3 Prof. J. T. Spencer, CHE 106 32 Polarity and Electronegativity • Electrons shared equally between like atoms is a non-polar covalent bond (i.e., H2, F2, O2, etc...). • Electrons may be unequally shared due to differences in Zeff between unlike nuclei (i.e., HF). Referred to as polar covalent bonding. • Electron density shifts toward the greater Zeff nuclei. • Electronegativity - used to estimate the polarity of a bond - the ability of an atom to attract electrons to itself. Prof. J. T. Spencer, CHE 106 Electronegativity 33 metals metalloids non-metals 1 2 3 1H 4 5 6 7 8 9 10 11 F EN (highest) = 4.0 Cs EN (lowest) = 0.7 12 13 14 15 16 17 Higher EN 18 2 He 3 Li 4 Be 5B 6C 7N 8O 9F 10 Ne 11 Na 12 M g 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 M n 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 M o 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 87 Fr 88 Ra 89 Ac 104 Unq 105 Unp 106 Unh 107 Ns 108 Hs 109 M t 69 Tm 70 Yb 71 Lu Lower EN 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr Prof. J. T. Spencer, CHE 106 Bond Polarity 34 • The larger the difference in EN between the two bonded atoms, the more unequal the electron sharing. • Bonds with unequal electron sharing are a polar bonds with the higher EN atom “holding” the greater electron density. IONIC F F EN = 0 Element F N B Li N F EN = 1.0 B F EN = 2.0 Electronegativity 4.0 3.0 2.0 1.0 Li F EN = 3.0 More Polar Prof. J. T. Spencer, CHE 106 35 Bond Polarity H F EN = 4.0 - 2.1 = 1.9 Polar Covalent Bond Classify the Bonds (give EN): S S non-polar B Cl polar covalent P F polar covalent P Cl polar covalent P O polar covalent Li F ionic N F polar covalent O O non-polar 0 1.0 1.9 0.9 1.4 3.0 1.0 0.0 Element F O N C B Li P Cl Br EN 4.0 3.5 3.0 2.5 2.0 1.0 2.1 3.0 2.8 Prof. J. T. Spencer, CHE 106 36 Drawing Lewis Structures • Lewis Structure describe many properties of molecules. To draw Lewis structures; – Sum the valence electrons for all atoms – Write symbols for the atoms and connect atoms in correct arrangement (with at least a single bond) – Complete the octets of the atoms bonded to the central atom (with some exceptions) – Place any remaining electrons on the central atom – If short electrons to gain central atom’s octet, try using multiple bonds. Prof. J. T. Spencer, CHE 106 Lewis Structures 37 . . . . . . . . . . . . . . .. . . Li . Be . B . . C N O . . . . : .F. : Ne .. : Element Li Be B C N O F Ne Elec. Config. Valence Electrons 1s22s1 1 1s22s2 2 1s22s22p1 3 1s22s22p2 4 1s22s22p3 5 1s22s22p4 6 1s22s22p5 7 1s22s22p6 8 Prof. J. T. Spencer, CHE 106 38 Formal Charge • Bookkeeping of valence electrons – all the unshared electrons are assigned to the atom on which they are found – Half of the bonding electrons are assigned to each atom in a bond – Formal charge - the number of valence electrons on the isolated atom minus the number of electrons assigned to the atom in the Lewis structure. formal ch. = [atomic valence e-] - [assigned e-] • The best Lewis structure (most stable) will be the one with the smallest formal charges and the one in which formal negative charges reside on the most electronegative atoms. Prof. J. T. Spencer, CHE 106 39 Formal Charge B in neutral atom has 3 valence electrons in [BH4]-, B assigned 4 electrons formal charge on B = -1 Atom N O (R) O (L) H -1 H B H H Neutral Assigned Formal Ch. 5 5 0 6 6 0 6 7 -1 -1 O N O Prof. J. T. Spencer, CHE 106 40 Lewis Structures • Best way to learn Lewis structures is PRACTICE! • Example (in class): – PCl3 – PH3 – H2O “Normal” – HClO2 – SeCl2 • [BH4]• [NO2]• [CO3]-2 –SO3 –HCN Charged Multiple Bonding Prof. J. T. Spencer, CHE 106 41 Lewis Structure: PCl3 = electron pair (1) Sum Valence Electrons: P =1x5 =5 Cl = 3 x 7 = 21 Total = 26 electrons = 13 pairs (2) Write symbols and include atom connections: Cl P Cl Cl (3) Complete Octets and place all remaining electrons: Cl P Cl Cl (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 42 Lewis Structure: PH3 = electron pair (1) Sum Valence Electrons: P =1x5 =5 H =3x1 =3 Total = 8 electrons = 4 pairs (2) Write symbols and include atom connections: H P H H (3) Complete Octets and place all remaining electrons: H P H H (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 43 Lewis Structure: H2O = electron pair (1) Sum Valence Electrons: H =2x1 =2 O =1x6 =6 Total = 8 electrons = 4 pairs (2) Write symbols and include atom connections: H O H (3) Complete Octets and place all remaining electrons: H O H (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 44 Lewis Structure: HClO2 (1) Sum Valence Electrons: = electron pair Cl = 1 x 7 =7 O =2x6 = 12 H =1x1 =1 Total = 20 electrons = 10 pairs (2) Write symbols and include atom connections: H Cl O O (3) Complete Octets and place all remaining electrons: H Cl O O (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 45 Lewis Structure: SeCl2 = electron pair (1) Sum Valence Electrons: Se = 1 x 6 =6 Cl = 2 x 7 = 14 Total = 20 electrons = 10 pairs (2) Write symbols and include atom connections: Cl Se Cl (3) Complete Octets and place all remaining electrons: Cl Se Cl (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 46 Lewis Structure: [BH4]- = electron pair (1) Sum Valence Electrons: B =1x3 =3 H =4x1 =4 charge =1 Total = 8 electrons = 4 pairs (2) Write symbols and include atom connections: H -1 H B H H (3) Complete Octets and place all remaining electrons: H -1 H B H H (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 47 Lewis Structure: NO2 - (1) Sum Valence Electrons: = electron pair N =1x5 =5 O =2x6 = 12 charge (-1) =1 Total = 18 electrons = 9 pairs (2) Write symbols and include atom connections: -1 O N O (3) Complete Octets and place all remaining electrons: -1 O N O (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 48 Lewis Structure: CO3 -2 (1) Sum Valence Electrons: = electron pair C =1x4 =4 O =3x6 = 18 charge (-1) =2 Total = 24 electrons = 12 pairs (2) Write symbols and include atom connections: -2 O C O O (3) Complete Octets and place all remaining electrons: -2 O C O O (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 49 Lewis Structure: SO3 -2 (1) Sum Valence Electrons: = electron pair C =1x6 =6 O =3x6 = 18 charge (-1) =2 Total = 26 electrons = 13 pairs (2) Write symbols and include atom connections: -1 O S O O (3) Complete Octets and place all remaining electrons: -1 O S O O (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 50 Lewis Structure: HCN (1) Sum Valence Electrons: = electron pair C =1x4 =4 N =1x5 =5 H =1x1 =1 Total = 10 electrons = 5 pairs (2) Write symbols and include atom connections: H C N (3) Complete Octets and place all remaining electrons: H C N (4) Check (including formal charges). Prof. J. T. Spencer, CHE 106 Resonance Structures 51 • Certain molecules cannot adequately be described using a single Lewis structure – Ozone is perfect example – takes two Lewis structures to describe ozone O O O O O O Prof. J. T. Spencer, CHE 106 Resonance Structures 52 • If the below description of ozone is correct, how is this possible? – Double bond is shorter than single bond so molecule would be lopsided O O O O O O Prof. J. T. Spencer, CHE 106 53 Resonance Structures • When several possible Lewis structures can be written which vary only in electron arrangement (not atom connections), each possible structure is a resonance structure (connected by double headed arrows). • The molecule “exists” as a weighted blend of the possible Lewis structures F F F F B B B F F F F F Prof. J. T. Spencer, CHE 106 54 Resonance Structures • Resonance is of crucial importance in organic chemistry – primary group showing resonance is the aromatics – aromatics are based on benzene Prof. J. T. Spencer, CHE 106 55 Octet Rule Exceptions • Exceptions to the Octet Rule: – Molecules with odd number of electrons (i.e., ClO2, NO2, NO, etc...). – Molecules in which the atom has less than an octet (i.e., BF3, H2, etc...). – Molecules in which an atom has more than an octet (i.e., PCl5, [ICl4]-1, XeOF4, etc...). –only possible for period 3 and beyond elements (requires n = 3 or greater) Prof. J. T. Spencer, CHE 106 56 Octet Rule Exceptions • Exceptions to the Octet Rule: – Molecules with odd number of electrons (i.e., ClO2, NO2, NO, etc...). • Complete pairing of electrons is impossible – Molecules in which the atom has less than an octet (i.e., BF3, H2, etc...). • Most encountered in molecules of Boron and Beryllium • allows the B or Be to bond with another atom using a coordinate covalent bond Prof. J. T. Spencer, CHE 106 Octet Expansions 57 PCl5 P=1x5 =5 Cl = 5 x 7 = 35 40 e- = 20 e- pairs Cl P XeOF4 Xe = 1 x 8 = 8 O=1x6 =6 F = 4 x 7 = 28 21 e- pairs -1 Cl Cl ICl4-1 I=1x7=7 4x7 = 28 charge = 1 18 e- pairs Cl Cl Cl Cl Cl I Cl O F F Xe F F Unshared electron pairs on central atom Prof. J. T. Spencer, CHE 106 58 • Practice: – CHCl3 – NH4+ – XeO4 – HNO2 – C2N2 – SCl2 – ClF3 Lewis Structures 13 prs 4 prs 16 prs 9 prs 9 prs 10 prs 14 prs Prof. J. T. Spencer, CHE 106 59 • Practice: – CHCl3 – NH4+ – XeO4 – HNO2 – C2N2 – SCl2 – ClF3 Lewis Structures 13 prs 4 prs 16 prs 9 prs 9 prs 10 prs 14 prs Cl H C Cl Cl H +1 H N H H Prof. J. T. Spencer, CHE 106 60 • Practice: – CHCl3 – NH4+ – XeO4 – HNO2 – C2N2 – SCl2 – ClF3 Lewis Structures 13 prs 4 prs 16 prs 9 prs 9 prs 10 prs 14 prs Cl H C Cl Cl H +1 H N H H O O H N O O Xe O O Prof. J. T. Spencer, CHE 106 Lewis Structures 61 • Practice: – CHCl3 – NH4+ – XeO4 – HNO2 – C2N2 – SCl2 – ClF3 N C 13 prs 4 prs 16 prs 9 prs 9 prs 10 prs 14 prs C N Cl H C Cl Cl H +1 H N H H O O H N O O Xe O O Cl S Cl F Cl F F Prof. J. T. Spencer, CHE 106 62 Bond Strengths • Stability of molecules is related to the strengths of the bonds within the molecules • Bond strengths can be determined thermochemically – Bond Dissociation Energy (BDE, H) F-F (g) 2 . F (g) H = 155 kJ mol-1 O=O (g) 2. O. (g) H = 495 kJ mol-1 NN (g) (g) H = 941 kJ mol-1 . 2. N . Prof. J. T. Spencer, CHE 106 63 Bond Strengths • BDE for diatomic molecules may be measured directly since only one bond exists within the molecule. • For polyatomic molecules (with many bonds), its difficult to determine the BDE directly so an average BDE is typically used. – Atomization Reactions: H . . .Si. (g) + 4 H (g) H = 1292 kJ mol-1 H Si H (g) . H ave BDE (CH)= 1292 kJ mol-1/ 4= 323 kJ mol-1 •BDEs are endothermic (positive) Prof. J. T. Spencer, CHE 106 64 Bond Dissociation Energies • Larger BDE’s generally lead to lower reactivity. • BDE’s follow bond order (b.o. and BDE 3 > 2 > 1). Bond C-C C=C CC N-N N=N NN C-N C=N CN Ave. BDE (kJ mol-1) 348 614 839 391 418 941 293 615 891 BO 1.0 2.0 3.0 1.0 2.0 3.0 1.0 2.0 3.0 Prof. J. T. Spencer, CHE 106 65 Bond Energies • Bond energies may be used to estimate H for reactions where bonds are broken or formed. – Determine which bonds are broken (endothermic) – Sum the BDE for these bonds – Determine which bonds are formed (exothermic) – Sum the BDE for these bonds – Estimate H by; H = (BDE bonds broken) - (BDE bonds formed) Prof. J. T. Spencer, CHE 106 66 H estimates from BDE’s estimate H for the reaction: H H H - C - O - H + H - Br H - C - Br + H - O - H H H H = (BDE bonds broken) - (BDE bonds formed) H = (BDE C-O + BDE H-Br ) - (BDE C-Br + BDE O-H) H = (358 + 366) - (276 + 463) kJ mol-1 H = (724) - (739) kJ mol-1 H = -15 kJ mol-1 Prof. J. T. Spencer, CHE 106 67 H estimates from BDE’s estimate H for the reaction: H H C=C + H-O-O-H H H H-O-C-C-O-H H H H H H = (BDE bonds broken) - (BDE bonds formed) Prof. J. T. Spencer, CHE 106 68 H estimates from BDE’s estimate H for the reaction: H H C=C + H-O-O-H H H H-O-C-C-O-H H H H H H = (BDE bonds broken) - (BDE bonds formed) H = (BDE C=C + BDE O-O ) - (2(BDE C-O) + BDE C-C) H = (614 + 146) - (2(358) + 348) kJ mol-1 H = (760) - (1064) kJ mol-1 H = -304 kJ mol-1 Prof. J. T. Spencer, CHE 106 69 Bond Lengths • As the number of bonds between two atoms increases (greater bond order), the bond becomes shorter and stronger. Bond C-C C=C CC N-N N=N NN C-N C=N CN Ave. BDE (kJ mol-1) 348 614 839 391 418 941 293 615 891 BO 1.0 2.0 3.0 1.0 2.0 3.0 1.0 2.0 3.0 Bond Length 1.54 1.34 1.20 1.47 1.24 1.10 1.43 1.38 1.16 Prof. J. T. Spencer, CHE 106 70 End Chapter Eight • Ionic Bonding • Octet Rule and Lewis Symbols and Lewis Structures • Hf° for ionic compounds - esp. Lattice energies and Born-Haber cycles. • Ionic Radii • Covalent Bonding • Bond Polarity and electronegativity • Lewis Structures (incl. resonance structures, formal charge, etc...) • Covalent Bond Lengths and Strengths