Intro Inorganic Chem
Intro Inorganic Chem CHM 218
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Date Created: 11/01/15
Ionization and Dissociation of Acids and Bases Section 71 72 Arrhenius Acid H20 HCI g 9 H aq C aq Arrhenius Base H20 BaOH2s e Ba2aq 20H aq Brmsted Lowry Acid HN03I H20 e H30aq N03 aq Brmsted Lowry Base NH3aq H20 a NH4aq OH aq Autoionization of Water Section 72 Amphiprotic H200 H200 H30aCI OHaCI IonProduct Constant for Water KW AcidBase Equilibrium Constants Section 73 AcidIonization Constant Ka BaseIonization Constant Kb Kw KaKb pKw pKa pr 1400 at 25 C Values of K3 and Kb for Common Weak Acids and Bases Table 1 71 Table 1 2 AcidIonization Constants at 25 312 I Basealonization Constants at 25 C Substance Formula K Substance Formula Kb Acetic acid HC2H302 17 X 10 5 Ammonia 7 NH3 13 x 10 5 Benzoic acid HC7H502 63 x 10395 Aniline c HSNHZ 42 x 1010 i r 10 Bone and H330 59 X 10 Dimelhylamine CH3an 51 x 10 4 Carbonic acid 112013 43 x 10 Ethylamine CBHSNH2 457 x 10 4 I 303 48 X 0 Hydrazinc N2H4 17 x 10 6 SW 1ij SOC 3 5 X 04 Hydroxylamine NHQDH 11 x 10 3 C acid H3302 13 m Mcthylamine CH3NH2 44 x 10 4 39 39 d1no c H N 14 x 10 9 Hydro uoric acid HF 53 x 111 4 1 141701 1 5 x 114 Hydrogen sulfate ion HSO4 11 x 10 2 2 Hydrogen sul de H38 89 X 10 3 Hg I 1 x I 531 Table 72 Base ionization constants of various Inorganic bases Hypocl orous acid HCIO 35 X 10 8 Base A HA Kb at 25 C pr Nitrous 45 X 10 quot Phosphate ion P043 HP042 X 10 2 oxalic acid H2904 5 5 x quotr Cyanide ion CN HCN 16 x 10 5 479 HC204 539 x 390 Ammonia NHQ NH4 18 x 105 474 Phosphor acld 234 X 1345 Hydrazinc N2H4 NszJr 85 X 104 63907 2 4 X Hpof 48 x 10quot3 phosplmmnS acid 11213110 If X 101 Table 71 Acid ionization constants of various inorganic acids HPHD3 7 x 10 7 Acid HA 39A Ka at 25 C pKa 7 W 5 Er x g39ggm 3233533 13 Pcrchloric acid HCIO4 310 1010 10 S erumus acid H 5303 3 I 1 Dan Hydrochloric acid HCI Cli 39 102 2 Hz 6 X 103 Hydro uoric39acid HF F 35 X 10 4 345 3 quot Ammonium ion NH4 NH3 55 x 10 10 926 The ionization constanls for polyprotic acids are for successive ionizations For example for H4304 the equilibrium is PERL H20 113039quot HZPO4 For HQPOK lhc aquiiibrium is HIPOK H20 H304 HPOfi i l his vain is in doubt Somc cvidcncc suggcsts that i is about ll W Scc RJ Myers J Chem Educ 613 137 1936 BranstedLowry Acids Section 74 Binary Acids HXaQ H200 9 H30aQ X39aQ Acid pKa Bond energy kJ mol l HFaq L 3 565 HCKaq 7 428 HBqu 9 362 HIM 10 295 Oxyacids 0HnX0m BranstedLowry Acids Continued Polyprotic Acids Diprotic Acid 3 H303CI HSO43CI H23043CI H200 HSO4 aq H20 lt2 H30aq SO4 aq Triprotic Acid H3PO4aq H200 H30aq H2PO aCI H2PO4 aq H20 lt2 30781 HPO4 aq Hpo42aq H20 ltgt H30aq P0431801 BrmstedLowry Bases Section 75 NH3aq H20 zgt NH4aq OH aq Conjugate Bases of Weak Acids POE ac H20 lt2 HPO42 aq OH aq SZaq H20 ltgt HS aq OH aq F aq H20 HFaq OH aq HS aq H20 a H28aq OH aq HP04239aq H200 lt gt H2P0439aq OH aq H2PO4 aq H20 lt quot H3PO4aq OH aq Trends in AcidBase Behavior Section 76 Aciditiy of Metal Cations Table 74 Acidity of some metal ions If 1362 Slightly acidic Weakly acidic Na M g2 A13 Neutral W eakly acidic Acidic K 332 SC3 Ti4 Neutral Slightly acidic Acidic Very acidic Na MgOH2 Mg2 A13 gt A1OH3 gtlt AMOH gt Trends in AcidBase Behavior Continued Basicity of Nonmetal Anions Basicity of Oxyanions Table 75 Basicity of some nonmetal ions N3 Very basic P3 Very basic A337 Very basic 02 Very basic 82 Basic 8e2 Basic quotfez Weakly basic F Weakly basic Cl Neutral Br Neutral 1 Neutral Table 76 Basicity of some common X0 oxyanions Classification Type Examples Neutral 04 ClO4 M1104 X03 NOs C103 Weakly basic XOf N02 C102 Moderately basic XO C10 Trends in AcidBase Behavior Continued Basicity of Oxyanions Cont Table 77 Basicity of some common XOJ39 oxyanions Ciassi cation Type Exampfes Neutral VC4w 1041 N11104 Weakly basic XOf 8042 CrOfquot M004 Moderately basic X043quot P043 V043 Strongly basic X044quot SiOf39 Table 78 Basicity of some common X0quot oxyanions Classi cation Examples Neutral NOg 3 C103 Moderately basic X032quot C032quot 803239 Hso43 r SO42quot H3P04 H2P04 m HPO4Z w P043 gt Y 3104 Network Covalent Substances Section 311 Network Covalent Bonding Diamond Quartz Intermolecular Forces Section 312 Two tvpes of interactions Intramolecular Attractions between atoms within molecules these are the chemical bonds that determine the chemical properties of the molecule Intermolecular Attractions between molecules these are intermolecular forces that determine the physical properties of the molecules Types of Intermolecular Forces IMF 1 London Dispersion Forces 2 DipoleDipole 3 Hydrogen Bonding 4 IonDipole Dispersion London Forces Section 313 London Dispersion Forces Induced Dipole s Na Na Na Na A B C Polarizability Factors 1 Number of Electrons 2 Molecular Shape Electronegativity Section 314 Electronegativity Molecular Geometry and Polarity Dipole Moment 1 a measure of the charge separation in molecules containing atoms of different electronegativities Overall dlpule momeni 0 Oman momenl Why does NH has a larger dipole moment p 14 D than NF p 02 D DipoleDipole Forces Section 315 DipoleDipole Attractions v 1 LI 39quot I x 1 N2 and CO HCI and HBr Hydrogen Bonding Section 316 Hydrogen Bonding 5 6 H quot 8 8 CI 3 5 TH3 H E H 2Hydrogen must be directly bonded to 0 N or F for hydrogen bonding 1 to occur 7 ymmm aH O H N H F H a f J 0 yo H quotH0VIH Summary of Intermolecular Forces Strengths of IMF39 London Dispersion lt DipoleDipole lt HBond lt IonDipole xcmow Wm cmumlmm 10mg mm mammum gt quot4 He me smug K H c H X H A mam m 5 Duane saw nMuduced a w e DlvnMnduoed mums D spels un Fundamental Properties of Group 17 Elements Fluorine Chlorine Bromine Iodine Astatine Symbol F C1 Br I At Atomic number 9 17 35 53 85 Natural isotopes 19F 100 35C1 7553 791315054 1271100 210A A abundances 37C1 2447 Br4946 Total no of isotopes 6 9 17 23 24 Atomic weight 1900 3545 7990 1269 210 Vaience electrons 2322125 3323515 4s24p5 55255 6s26p5 mp bp C 220 188 101 34 73 588 114184 302337 Density 6 181 gL 321 g L 312 gcm3 494 g cm3 Covalent radius A 072 099 114 133 Ionic radius a ShannonPrewitt A CN 1196 1676 1826 2066 Pauling EN 40 30 28 25 22 Charge density 084 060 055 049 chargeionic radius unit charge A E b v 287 136 107 054 03 Oxidation states ml 1 to 7 a 1 to 7 1 to 7 1 1 3 5 Ionization energy k1 mol 1680 1251 1143 1009 916 Electron af nity kJm01 333 348 324 295 Discovered by date Moisson Scheele Balard Courtois 6015011 1886 1774 1826 1811 MacKenzie Segre 1940 rpwC O2 02F2 None None None None Acidbase character Acid Acid Acid Acid Acid of oxide rpw N2 None None None None None rpw halogens See Section 184 on Neutral and Ionic Interhaiogens rpw hydrogen HF BC 1181 HI HAt Crystal structure 39 Ortho Ortho rhomb rhomb Group Trends Section 171 toCommon Features Halogens 00 Melting and Boiling Points Table 171 Melting and boiling points of the Group 17 elen Melting Boiiing Number of Element point C point 0 electrons F2 219 188 A 18 C12 101 34 34 Br 7 60 70 12 114 185 106 ozoOxidation States Anomalous Nature of Fluorine Section 172 otoUniqueness Principle ozoWeakness of F F bond otoBonding Limitations otoHigh Electronegativity ozolonic Character of Fluorides ozoStabilization of High Oxidation States in Metals ozoDifferences in Fluoride Solubilities Fluorine Section 173 t Most Reactive Element Hg Clltggt Fzg 183 k Clzg T105 k 333 k 349 k Fg V Clwg v ozoCovalent vs Ionic Fluorides 460 k 348 k 730 k J C1 aql F011V Table 174 The lattiCe energies of sodium halides Lattice energy Halide kJ39morl NaF quot 915 NaCi 78 1 NaBr 743 Nail 699 Chlorine Section 174 Free energy V39 Inol e 10 8 6 4 2 10 Acidic solution C10 C10 4 C10 3 Basic solution C102 C1 I I I I I 7 5 3 1 0 1 Oxidation state Sections 175 176 oonydrogen Fluoride and Hydrofluoric Acid Hydrochloric Acid Halides Section 17 7 ozolonic Halides ozoCovalent Halides Chlorine Oxides and Oxyacids Sections 178 179 t Chlorine Monoxide t Chlorine Dioxide O ozoChlorine Oxyacids amp Oxyanions H W C C1 0C1 OH OH OH Chlorine Bromine amp Iodine Oxyacids The Known Chlorine Bromine and Iodine Oxoacids Oxidation state 41 3 5 7 Hypohalous Halons Hallo Perhalic acid acid acid acid Chlorine HOCI HOCIO H0002 HOCIO3 K 34 gtlt 10quot8 11 x 102 55 x 102 1x 108 5 acid 163 Vquot 164 v3 147 Vb 142 Vb Equot anion 089 vquot 078 V 063 V c 056 V Bromine HOBr HOBIOQ HOBr02 HOBrO3 Ka 2 x 109 10 Large 5 acid 159 Va 152 Vquot 159 vb 5 anion 076 Vc 061 V 069 v6 Iodine H01 HOIO2 11110103d Ka 2 x 10 16 x 101 Large Equot acid 145 Va 120 vb 134 vquot Equot anion 049 v 026 V0 039 V aFor acid to halogen in acid solution bFor oxoanion to halogen in acid solution Cifor oxoanion to halide in basic solution dAlso occurs as H5106 Ka 5 x 10 K02 w 5 x 109 K83 2 x 1039 Interhalogen Compounds and Polyhalide Ions Section 1710 otolnterhalogens XYn ozoPolyhalide Ions Interhalogen Compounds and Polyhalide Ions The Neutral Binary Interhalogens Gasphase structure Color and phase at STP Preparation CIF Linear Colorless gas Direct and C12 4 C1173 3ng Distorted T shaped Colorless gas Direct 3ng Square pyramidal Colorless gas F2 CW3 BrCl Linear Red brown gas Direct BrF Linear Palebrown gas Direct and 132 33ng 13ng Distorted T shaped Yellow liquid Direct BIFS Square pyramidal Colorless liquid BrF3 F2 IBr Linear Black solid Direct ICl Linear Red solid Direct IF Linear Brown solid Direct and 12 IF3 IC13 Planar dimer Orange solid Direct IF3 Triagonal planar Yellow solid Direct IF5 Square pyramidal Colorless liquid Direct IF7 Pentagonal bipyramldal Colorless gas F2 IF Binary Interhalogen and Polyatomic Monohalogen Ions Central halogen Terminal halogen Chlorine Bromine Iodine Fluorine CIFS CIFJ Ber Ber IF2 IF Cl 2F CIR Ber BrF IF IF CIR BrFff BrFf 1E IF CIFg39 IFgquot Chlorine CE Cl Brle BrCl 2 1C1 ICI Brzcr 120 Cl 12C 1C1 Bromine Br Br 13 IBrzquot Brr 12 Br 1281 w Brs Iodine I g 13 13 Is Ts 17 13 15 13 1 Element Reaction Flowchart Section 17 12 2 BTO FeF3 8116 Fe 510 1310 8amp4 F2 2 HF 4um 5 aw w w izo 396 OH F 117 C1133 F HFZ F2 HCIO I e 12 103 12 gt 1 gt I OH H u cg N33 10 38 0400 NH oa v 1c1 Fed3 Fe af l HCI Fe FeClg P 2 OH39A 1 P015 c103 m e gtc1o4 H 103 Isotopes of Hydrogen Section 101 Protium Hydrogen H2 Deuterium D2 Tritium T2 Table 101 Physical properties of the isotopes of hydrogen Isotope Molar mass g mol l Boiling point K Bond energy kJ mol l Hz 202 206 436 D2 403 239 443 T2 603 252 447 Properties of Hydrogen Section 103 2Hg 0g F otoProperties 436 k H2g 0g 928 k 248 k H2ggt 02ggt otoPreparation 244 k H20g J The Hydrogen Economy Metallic hydrides Storage for Hydrogen production facility 1 Oil or 0031 syngas 2 Photochemical 3 Thermal cycles H 2 g mm woe Truck and or train transport peak demands Metallic hydrides Storage facilities Electrici Fuel ty cells H20 f etroleum Fuels re ning W Organic chemicals for drugs plastic etc Hydmgam W Partially saturated a on solids fats facilities LUnsaturated oils Metallic Metals ore re ning LOres Coal W Fuels gagi cgtmn M Organic chemicals an for drugs piaslic etc liquefaction LCoal Haber NH3 g nitrates process facility Lung Hydrides Section 104 mm BCNOF Fe Co Ni Ru RhAg Cd In Sn Sb Te I Os Ir Pt Au Hg TI PbEmAtm EISaline El Metallic 1 Intermediate Molecular Unknoan 1 Saline Ionic Hydrides 2 Metallic Hydrides Hydrides Cont 3 Molecular Covalent Hydrides 4 Extended Polymeric Covalent Hydrides Boiling Points of Hydrides 120 100 80 3 E 393 2 m lt100 20 IZG 40 gt140 60 gt160 HES 20 40 60 80 00 I20 20 40 60 80 100 120 Molecular Weight Molecular weight A B Bonding in Diborane BZHG H 1 H 3 B H H H Figure 106 Electron pair arrangement in diborane BZHB HB H H H H H Figure 107 Geometry of the diborane molecule Bridging H Terminal H atom Fig 426 The structure of B2H6 determined by electron di raction CH3 Hydrogen bonds HCCCO HNH N N Adenine C H C N H N c O C H C C H N N Water and Hydrogen Bonding Section 105 Second strand Clatherates Section 106 Clatherates substance in which moleculese or atoms are trapped within the crystalline framework of other molecules Element Reaction Flowchart Section 108 NaH Na NaBH4 LiA1H4 2 CU NaH LiH NH3 Ca LaH2r87 Na Li UH3 La catalyst catalyst CH HO gt H2 lt HOCO 4 2 co co2 2 CuO catalyst N F8304 Cu0 NH3 W0 Fe0 Fundamental Properties of Group 18 Elements Helium Neon Argon Krypton Xenon Radon Symbol He Ne Ar Kr Xe Rn Atomic number 2 10 18 36 54 86 Natural isotopes He 100 20Ila9092 36Ar0337 78Kr 035 12 Xe0096 222 Rna A abundances 3Hes300013 2 Ne A1257 38Air0063 8017227 126Xe0090 22Ive882 Ar9960 82Kr1156 128Xe192 83191155 129Xe2644 8 Kr5690 13 Xe408 Kr1737 131Xe2118 132 Xe 2689 134 Xe 1044 136 Xe 887 Total no of isotopes 5 8 8 21 25 20 Atomic weight 4003 2018 3995 8380 1313 222 Valence electrons 152 2322p6 3533p6 432416 5325126 636 mp bp C 272 269 249 246 189 486 157 152 112 107 71 62 Density gliter 0177 0900 13784 3733 5887 973 Covalent radius 103 39 39 39 39 39 115 126 Van der Waals radius A 140 154 188 202 216 Oxidation states 0 0 0 0 2 0 2 4 6 0 2 ionization energy 2373 2080 1521 1241 1167 kJ mol Electron af nity 2 1 29 35 39 40 kJmol estimates Isolated by date Ramsay Travers Rayleigh Travers Travers Dom 1895 amp Ramsay amp Ramsay amp Ramsay amp Ramsay i900 1898 1894 1898 r 1898 rpwb halogens Ker Xan n 2 4 6 Crystal structure hep fcc fcc fcc fcc fcc aLongest lived isotope 212 382 days Group Trends Sections 181 183 ozoCommon Features and Uses Noble Gases t Melting and Boiling Points Table 181 Melting and boiling points of the noble gases Melting Boiling Number of Noble gas point C point C electrons He 269 2 Ne 249 245 10 Ar 189 186 18 Kr 157 152 36 Xe l 12 109 54 Rn 71 62 86 Xenon Fluorides and Oxides Sections 185 186 a Xer 6 X603 3 X604 g Xe02 3 b XeF4 Xng h f X606 1 Xeom O 900 I o 3 o O X F4 H XEFZ Decay step Element Reaction Flowchart Section 188 1 H10 3 2 XeF6 X60134 Xe03 i HXeO4 igt Xe064 E BaZXeO E91 er4 Halflife 2 U gt 233 gHe 23m gt2 fPa fie Zara 2331 fie 23311 32 3 233m gt 2 Ra gHe 233m gt 23an 3H6 2an gt l iPo gHe zgipo 925131 He 2ng 623151 fie 28131392ng 16 281po 92ng 3H6 2ng 53131 36 Z gBi z gPo 31c 2 2P0 gt 23ng g e 45 X 109 years 24 days 12 min 25 X 105 years 80 X 104 years 16 X 103 years 38 days 31 min 27 min 20 min 16 x 104 s 22 years 50 days 138 days 151 Chapter 18 The Group 18 Elements The Noble Gases Element mp C bp 0C He W 7269 All the elements in Group 18 are colorless odorless monatomic gases at Ne 7249 7245 room temperature They neither bum nor support combustion in fact Ar 189 186 they make up the least reactive group in the Periodic Table with only xenon forming a variety of compounds The very low melting and boiling Kr 7157 7152 points of the noble gases indicate that the dispersion forces holding the Xe 7112 7109 atoms together in the solid and liquid phases are very weak The increase in the melting and boiling points with increasing atomic weight is R 71 62 due to the increase in polarizability To date chemical compounds have been isolated at room temperature only for the three heaviest members of the group krypton xenon and radon Few compounds of krypton are known and tend to be very unstable at room temperature whereas xenon has an extensive chemistry The study of radon chemistry is very dif cult because all the isotopes of radon are highly radioactive the longest lived isotope has a halflife of less than 4 days Unique Features of Helium When helium is cooled to as close to absolute zero as possible it is still a liquid At 10 K a pressure of about 25 MPa is required to cause it to solidify However liquid helium is an amazing substance At ambient pressure the gas condenses at 42 K to form an ordinary liquid helium I but when cooled below 22 K the properties of the liquid now helium II are dramatically diiTerent For example helium II is an incredibly good thermal conductor 106 times greater than helium I and much better than even silver the best metallic conductor at room temperature Even more amazing its viscosity drops to close to 0 When helium II is placed in an open container it literally quotclimbs the wallsquot and runs out over the edges These and many other bizarre phenomena exhibited by helium II are best interpreted in terms of quantum behavior in the lowest possible energy states of the element A lll discussion is in the realm of quantum physics The Chemistry of the Noble Gases In 1962 Neil Bartlett and DH Lohmann at the University of British Columbia discovered that Pth was a su iciently potent oxidizing agent to oxidize dioxygen to the dioxygenyl cation 02 02 1th a 021 Pthli Since the ionization energy of 02 1177 ldmol is similar to that of Xe 1170 ldmol Bartlett reasoned that Pth might also oxidize Xe The direct reaction of Xe with Pth yielded a redorange crystalline product which was insoluble in CCL and originally thought to be Xe PtF 639 Later work suggested that the product was likely XeF Pt2F1139 or a mixture of XeF P12F1139 and XeF PtF539 Xenon Fluorides The first binary compounds of Xe to be prepared were uorides and since uorine is the only element that reacts directly with Xe the uorides also serve as starting materials for the synthesis of other Xe compounds All of the uorides are white solids and have negative free energies of formation at 25 C that is they are stable to dissociation at room temperature Xenon di uoride is prepared by the reaction of excess xenon with uorine at high pressure aided by heating or electromagnetic radiation Xeg F2 g a Xer S The preparation of xenontetra uoride is carried out by heating a 15 xenon uorine mixture to about 400 C under a pressure of about 6 atmospheres for several hours Xeg 2F2g XeF4 S These conditions are also employed in the preparation of XeFG exceptthat a large excess of F2 is used and pressures of gt50 atm are required for quantitative conversion Xeg 2F2g XeF4 S The structures of XeF2 and XeF4 are predictable on the basis of the VSEPR model With 8 valence shell electrons from the xenon atom and 2 additional electrons from the two uorine atoms there are 10 electrons surrounding the xenon atom in XeFZ Thus XeF2 is a linear molecule 8quot 39n X n The structure of XeF4 is derived from the fact that there are 12 electrons surrounding the Xe atom in the molecule 8 from the Xe atom and 1 l lr39F from each uorine atom Therefore the 12 electrons reside in six F7xe F orbitals pointing toward the comers of an octahedron and the molecular F structure is square planar The structure of XeF6 is not a regular octahedron because the xenon atom has 14 valence electrons The IF7 molecule is similar in this regard and it I has a pentagonal bipyrarnid structure However with one unshared pair of F electrons there is some question as to where the unshared pair resides Xltuf Also the molecule is not rigid and it is described as an capped 1 F octahedron 1n condensed phases monomme b m K 1 g 5 3 5 g slmple as those oerF2 andXeF W 111 11 Lummh m m 1mm WW1 thmm 1n gma al1nfzennolecular forces are R m E A E w 1 mtaesnng that me memng p01an oerFz Ken and Ken are 129 117 and 50 C respemvely but me sohds readdly subhme Reacu39nns nr Xennn Flunrides but slowly m and sdudms and a 015 M sduam ean be prepared at 0 C 1n base sduam me hydrolysxs occurs rapxdly Xers ZOH39 aq a Xegg 201g 2F39 aq Hzo 1 The hydrolysxs of XeF 15 Exh39Ernely vxgorous andxt undegoes a daspropomonanon Eamon that 15 sxmlarto that ofthe halogens SXeF rlZHZO e2 lteos 4Xe 30224HF whmh represents Xeav producmg XeOI and Xe0 The omdeXeO1s a vary dangerously explosxve compound Hydrolyss of Xer also produces ths ends by the reaction Ken 3 H20 X20 SHF although the reaction appears to take p1aee m two seeps as shown by me dudmng equaadns 2XeF H20 e2 lteo12m XeOF 2HZO X2034HF l 5 4 The uorides are strong uorinating agents For example xenon di uoride can be used to uorinate double bonds in organic compounds It is a very clean uorinating agent in that the inert xenon gas can be easily separated from the desired product Xer s HzcCH2g a HzFCCFH2gXeg A use ll derivative of uracil is 5 uorouracil which is used inthe treatment of some types of skin cancer One preparation of the compound is by the uorination of uracil with XeFZ O H H F N N l k XeFZ gt I x Xe HF N o N o The highest oxidation state of many uorides can be produced by using xenon uorides as reagents For example xenon tetra uoride will oxidize sulfur tetra uoride to sulfur hexa uoride XeF4 s 2 SF4 g a 2 SF6 g Xeg Previously several reactions of interhalogens were shown in which cations were produced by a reaction with a strong uoride ion acceptor A reaction of this type is 01F3 SbF5 C1F2SbF539 The xenon uorides undergo similar reactions with uoride acceptors Such as SbF5 Ast TaF5 and Pth The reaction of F XeF2 with SbF5 can produce XeF SbZFH39 XeFSbF539 or X6 X6 XezFf SbFG39 with F39 bridges between Xe centers in the cation F V F Thus Xe2F3 represents a uoride ion bridging between two 1510 XeF ions as shown In species that contain the XeF ion the Xe is associated with the ME anions by attaching to one ofthe uoride ions F39 quot 39 Muquotquot39F in the complex F r I FXef F The XeF3 cation is generated when XeF4 reacts with BiF 5 as shown by the equation XeF4 BiF5 a XeFfBiFG39 Solid XeF6 contains XeF ions that have uoride ions bridging them The cation is also generated when XeF6 reacts with penta uorides such as RuF5 Xer RuF5 H XerlRuFG39 A cation having the composition Xe2F11 that has the arrangement F5Xe 39 39 39F39 39 39 XeF has also been identi ed In some cases Xer itselfforms anions such as XeFf and XeFSZ39 by virtue of its Lewis acidity A general reaction can be written as Xer MF MXeFf When heated some compounds of this type undergo a reaction to produce M2XeF8 and XeFG Xenon Oxy uorides and Oxides Xenon oxides are prepared from the uorides As we have already mentioned hydrolysis of XeF4 and Xer lead to the formation of XeO3 by reactions that can be shown as 6XeF4l2H20 a 2XeO3 4 Xe 24HF 302 XeF6 3 H20 4 XeO3 6 HF The latter reaction probably involves the steps XeF6 H20 4 XeOF4 2 HF XeOF4 2 H20 4 XeO3 4 HF These reactions are similar to the hydrolysis of PC15 which can lead to OPC13 when a limited amount of water is present Solid XeO3 is a violently explosive white solid that has a very high positive heat of formation of approximately 400 kJmol Since the four atoms in XeO3 have a total of 26 valence shell electrons the predicted structure has 0 I three bonds and an unshared pair of electrons surrounding the xenon atom quot I 39 However the resulting 3 formal charge on the xenon atom indicates that there should be contributions to the actual structure from resonance structures having multiple bonds between Xe and O In the presence of OH the reaction of XeO3 leads to the formation of a hydrogen xenate HXeO439 ion XeO3 OH39 H HXeO439 Disproportionation of this unstable species occurs according to the equation 2HXe0439 20H39 a XeO6 39 Xe022H20 Solid perxenate Xe05439 salts can be obtained that contain cations of Group IA and IIA metals The Xe05439 ion has a very weak conjugate acid so the hydrolysis reactions Xe05439 H20 HXeOf OH39 HXeOf H20 HZXeOGZ39 OH39 are extensive and the solutions are basic as well as very strong oxidizing agents In reactions that are very similar to that between PC15 and P4010 6 PCl5 P4010 10 OPCI3 xenon uorides react with xenon oxides to produce oxy uorides XeF6 2 XeO3 a 3 Xe02F2 2 XeF6 XeO3 a 3 XeOF4 The Chemistry of Krypton and Radon Krypton di uoride is obtained when an electric discharge is passed through Kr and F2 at 183 C or when the gases are irradiated with highenergy electrons or protons It is a volatile white solid that decomposes slowly at room temperature It is a highly reactive uorinating agent The linear KrF2 molecule is thermodynamically unstable whereas XeF2 is stable as the following free energies of formation show Kr g mg 1er s AGf 95 1d Xe g F2 g a Xer 5 Ag 90 1d 157 No other molecular uoride of Kr has been isolated The cationic species KrF and Kr2F3 are formed in reactions of Ker with strong uoride acceptors such as AsF5 and SbF5 and compounds have formulas such as KrFSb2F1139 KrFSbF539 and Kr2F3AsF539 The KrF salts can be used to uorinate NF3 to NFf There is some eVidence for KrFeCO5 in gammairradiated matrices at low temperatures Since radon has only a short halflife study is dif cult but tracer studies allow some properties to be deduced for example the formation of Ran RnFlTaFG39 and possibly RnO3 Oxidation of Rn by ClF 3 and study on a uorinated ionexchange material N a on suggests that Rn can displace Na or Kl Group Trends Section 91 otoGroups of elements tend to have characteristic properties Descending a group in the periodic table there are often smooth trends in these properties 1 Alkali Metals Group 1 Table 91 Melting and boiling points for the alkali metals Element Melting point C Boiling point 0 Lithium 180 I 1330 Sodium 98 892 Potassium 64 759 Rubidium 39 700 Cesium 29 690 Table 9239 Some physical properties of the alkali metals Ionization energy Hydration enthalpy 1 Element kJmol39l kJmol Electrode potential V Lithium 526 519 304 Sodium 502 406 271 Potassium 425 322 294 Rubidium 403 301 292 Cesium 376 276 303 Group Trends Cont 2 Halogens Group 17 Table 93 Melting and boiling points for the halogens Element Melting point C Boiling point C Fluorine quot21 9 188 Chlorine 101 W 34 Bromine 7 60 Iodine 1 14 185 Table 94 Some physical properties of the halogens Bond energy Electron affinity Element kJ39mol39l kJmol l Electrode potential V Fluorine 155 328 305 Chlorine 240 349 136 Bromine 190 331 109 Iodine 149 301 054 Group Trends Cont 3 Group 15 Pnicogens Table 95 Melting and boiling points for the Group 15 elements Element Melting point 0 Boiling point C Nitrogen i 210 196 Phosphorus 44 281 Arsenic 61 5sub Antimony 63 1 1387 Bismuth 27 1 1 5 64 Periodic Trends in Bonding Section 92 t Crossing a period we observe systematic patterns in chemical formulas of the compounds formed by the elements In addition there are partial trends in physical and chemical properties of the elements 1 Bonding Trends in the 2quotdl Period Elements Table 96 Melting points of the Period 2 elements Element Li Be B C N2 02 F2 Ne Meltingpoint C 180 1287 2180 4100sub 210 229 219 249 Periodic Trends in Bonding Cont 2 Bonding Trends in the 3rd Period Elements Table 97 Melting points of the Period 3 elements Element Na Mg vAl Si P4 38 CI2 Ar Meltingpoin C 98 649 660 1420 44 119 101 189 H2 He c N2 02 F2 Ne P4 Sg C12 Ar Seg Brz Kr 12 Xe Atz Rn Figure 92 The common forms of some nonmetallic elements Periodic Trends in Bonding Cont 3 Bonding Trends in the Highest Fluorides of the 2ndl and 3lrdl Periods Table 98 Formulas bonding types and phases at room temperature of the highest fluorides of the Periods 2 and 3 elements Compound LiF BeF2 BF3 CF4 NF3 OF2 Bonding Ionic Network Covalent Covalent Covalent Covalent type phase solid covalent gas gas gas gas solid Compound NaF Mng AlF3 SiF4 PFS SF6 ClF Bonding Ionic Ionic Network Covalent Covalent Covalent Covalent type phase solid solid covalent gas gas gas gas solid Periodic Trends in Bonding Cont 4 Bonding Trends in the Highest Oxides of the 2ndl and 3lrdl Periods Table 99 Formulas bonding types and phases at room temperature of the highest oxides of the Periods 2 and 3 elements Compound Li20 BeO B203 002 N205 F20 Bonding Ionic Ionic Network Covalent Covalent Covalent type phase solid solid covalent gas gas gas solid compound N320 A1203 P4010 C1207 Bonding Ionic Ionic Ionic Network Covalent Covalent Covalent type phase solid solid solid covalent solid solid liquid solid Periodic Trends in Bonding Cont 5 Bonding Trends in the Hydrides of the 2nd and 3rdl Periods Table 911 Formulas bonding types and phases at room temperature of the hydrides of the Periods 2 and 3 elements Compound LiH BeHzX 82H6 CH4 NH3 H20 HF Bonding Ionic Network Covalent Covalent Covalent Covalent Covalent type phase solid covalent gas gas gas liquid liquid solid Compound NaH MgHz AlH3x SiH4 PH3 H28 HCl Bonding Ionic Ionic Network Covalent Covalent Covalent Covalent type phase solid solid covalent gas gas gas gas solid Trends in AcidBase Properties Section 94 1 AcidBase Patterns in the Highest Oxides of the 2nd and 3rd Periods Table 913 Acid base properties of the highest oxides of the Period 3 elements Compound N820 MgO A203 Si02 P4010 8033 C207 Acid base Basic Basic Amphoteric Acidic Acidic Acidic Acidic behavior Trends in AcidBase Properties Cont 2 AcidBase Patterns in the Highest Oxides of Group 15 Table 914 Acid base properties of the highest oxides of the Group 15 elements N205 Acidic P4010 Acidic A5203 Acidic Sb203 Amphoteric Bi203 Basic 3 AcidBase Patterns in the Covalent Hydrides of the 2quotd and 3rd Periods Table 915 Acidbase patterns in the covalent hydrides of the Periods 2 and 3 elements Compound BZH6 CH4 NH3 H20 HF Acid base behavior Neutral Neutral Basic Weakly acidic Compound SiH4 PH3 H28 HCl Acid base behavior Neutral Neutral Very weakly acidic Strongly acidic The n Group and n 10 Group Similarities Section 95 There are similarities in chemical formulas and structures of the hihest oxidation state of some members of the n Group elements and of the members of the correspondin n 10 Group elements 313 414 515 616 717 122 The n Group and n 10 Group Similarities Cont 1 Aluminum Group 13 and Scandium Group 3 Table 916 A comparison of some properties of Group 3 and Group 13 elements Group 3 elements Group 13 elements Element mp C E V ement mp C E V A1 660 166 Sc 1540 188 Ga 30 053 Y 1500 237 In 160 034 La 920 252 T1 300 072 Ac 1050 26 The n Group and n 10 Group Similarities Cont 2 TinV Group 14 and TitaniumV Group 4 3 PhosphorusV Group 15 and VanadiumV Group 5 4 SulfurVl Group 16 and ChromiumVl Group 6 Table 917 Similarities between chromiumVl and sulfurV species Group 6 Group 16 Formula Systematic name Formula Systematic name Cr03 ChromiumltVI oxide 803 Sulfur trioxidc CrOZClz Chromyl chloride SOZClz Sulfuryl chloride CrOf Chromate ion SO42 Sulfate ion Cr2072 Dichromate ion 820727 Pyrosulfate ion 5 ChlorineV Group 17 and ManganeseVl Group 7 6 XenonV Group 18 and OsmiumVll Group 8 The n Group and n 10 Group Similarities Cont 7 Alkali Metals Group 1 and the Coinage Metals Group 11 Table 918 Contrast of the alkali metals and the coinage metals Property Alkali metals Coinage metals Common oxidation numbers Chemical reactivity Density Melting points Aqueous redox chemistry Solubilities of common salts Always 1 Very high increasing down the group Very low increasing down the group 05 to 19 gcm s Very low decreasing down the group I 181 C to 29 C None All soluble Silver 1 but copper and gold rarely 1 Very low decreasing down the group High increasing down the group 9 to 19 gcm g High all about 1000 C Yes eg Cu2nq gtCunq 1 Oxidation state compounds insoluble 8 Magnesium Group 2 and Zinc Group 12 Isomorphism in Ionic Compounds Section 96 Alums have the general formula NFMOH263SO42392 6H20 Isomorphous Substitution analogous crystal structures 1 2 Table 920 A comparison of cation radii in alums Monopositive ions Tripositive ions K 152 pm 1A1 68pm Rb 166 pm Crer 75 pm NH4 151 pm Fe3 78 pm Table 921 Some comparative ion sizes ionic radii 1 Charge 2 Charge 3 Charge 4 Charge 5 Charge Small Be 1 A1 Fe Si P CrSi Medium Li Mg Fe Ti W5 Large NaJr Ca2 L213 Very large 1C NHL B212 Diagonal Relationships Section 97 ZThe similarity in chemical properties between an element and the element to the lower right of it This relationship is found for element in the upperleft corner of the Periodic Table The quotKnight39s Move Relationship Section 98 otoThe similarity between an element and the element of Group n and Period m with the element in Group n 2 and Period m 1 in the same oxidation state This relationship is found among elements in the lowerright portion of the Periodic Table Cu Zn Ga Ag Cd In Sn Sb T1 Pb Bi The Lanthanoid Relationships Section 910 oonhe lanthanoids have very similar properties across the period The Combo Elements Section 911 otrThe combination of an n X group element with an n X group element to form compounds that parallel those of the n group element Boron nitride Figure 99 Comparative layer structures of boron nitride and graphite Borazine B N O 0 Figure 910 Repeating layer structure of boron nitride The Combo Elements Cont Combo elements and Semiconductors Zintl Principle The PseudoElements Section 912 03A polyatomic ion whose behavior in many ways mimics that of an ion of an element or a group of elements 1 NH4 2 CN39 Review of Thermodynamics Enthalpy Chapter 61 First Law of Thermodynamics Internal Energy U AUqw 1 Constant Volume AU qv Enthalpy H H u pV 2 Constant Pressure AH qp Enthalpy Change AH Endothermic Exothermic Entropy and Gibbs Free Energy Section 61 Second Law of Thermodynamics Entropy S Entropy Change AS Gibbs Free Energy G AG AH TAS AG as Criterion for Spontaneity Section 61 AG AH TAS Standard State 298 K and 100 kPa 1 bar AG AH TAS Table 1 53 Effect of Temperature on the Spontaneity 0 Reactions AH A5 A6 Descriplinn39 Example Spontaneous at all T C6H206s gt ZCZHSOHU 2C02g Nonspomaneous at all T 302g gt 203g or Spontaneous at low T 2N39H3g C02g gt NH2C0NH2aq H100 nonsponlaneous at high T or 7 Nonsponlaneous at low T Ba 0H2 stom ZNH4NO3 gt spontaneous at high T BaNO2aq 2NH3g IOHZOU Enthalpy of Formation Section 61 Standard Enthalpy of Formation AH f P4S 5029 9 P4O1OS AH f P4S AH fP4O1OS AH fOZg Write the equation for the formation of NH3g from its elements AHf for an element in its standard state reference form is zero Enthalpy of Reaction Section 61 Standard Enthalpy of Reaction AH the enthalpy change for a reaction in which the reactants in their standard states yield products in their standard states Two different methods to calculate AH 1 Direct Method Reactants 9 Products AH AH fproducts AH freactants 2 Indirect Method Hess s Law A B 9 C AH 1 100 kJ C B 9 D AH 2 350 kJ A 2 B 9 D AH 3 450 kJ Standard Enthalpy of Reaction AH Direct Method aA b8 9 CC dD AH cAH fC dAH fD aAH fA bAH fB Calculate AH for the following reaction using values ofAH f 2 HgSg 3 029 9 2 H20I 2 3029 AH AH fHZSg 2o2 kJmol AH fHZOI AH f 3029 2858 kJmol 2961 kJmol Standard Entropy of Reaction AS Direct Method aA b8 9 CC dD AS cS C dS D aS A bS B Nas 1639 29 9 NaCls AS 8 Nas 51 Jmol K 8 NaCls 72 Jmol K 8 C2g 223 Jmol K Standard Enthalpy of Reaction AH Hess 5 Law Hess s Law For a chemical equation that can be written as the sum of two or more steps the enthalpy change AH for the overall equation equals the sum of the enthalpy changes for the individual steps Given the following data AH kJ H2g 9 2 Hg 4364 Br2g 9 2 Brg 1925 724 H2g Br2g 9 2 HBrg Calculate AH kJ for the following reaction Hg Brg 9 HBrg Standard Enthalpy of Reaction AH Hess 5 Law Given the following data AH le COCI2g H20 9 CH2CI2I 029 475 HCg 9 12 H2g 12 C2g 2300 COCI2g 2 H20I a CH202I H2g 32 029 4025 Calculate AH kJ for the following reaction H200 C2g 9 2HCKQJ 1202g Bond Energies Section 61 Bond energy BE the enthalpy change required to break a particular bond in 1 mole of a gaseous molecule F2g 9 2 Fg F F 9 2 F Cl Cl 9 2 Cl Br Br 9 2 Br I I 9 2 I What is the 0 H bond energy in CH4g CH49 9 09 4Hg AH 158 BEF F 158 kJmol BECI CI 242 kJmol BEBr Br 193 kJmol BEI I 151 kJmol AH 1664 kJmol Bond Energies Cont Table 95 Bond Energies in ldmol single Bends H C N O S F Cl Br I O TlUIQZO p u 0 m u m N S 3 N 208 175 149 N 3 N L l to o l N l on Multiple Bonds 615 CO 745 799 in cog 887 CEO 1072 607 50 in 02 532 494 SO in 50 469 u 1 Hu 9 new ml m r Ch Mum New York HarperCollins 1993 pp A217A34 H l zfno 2200 is a on w u 2min 0022 o 4 N 0 Using Bond Energies BE to Calculate AH AH BEbonds broken BEbonds formed Using bond energies estimate the enthalpy of reaction AH xn for the following 2CH49 9 H3CCH39 H29 Bond BEkJmol H H 432 C H 411 C C 346 CC 602 C C 835 Lattice Energies Section 61 Lattice Energy or Lattice Enthalpy U M g nX g 9 MXns BornLand Equation NA Avogadro s number 6023 x 1023 mol391 A Madelung Constant 2 relative cation charge 239 relative anion charge e electron charge 1602 x 103919 C en permittivity of free space 8854 x 10 12 Czm391J391 r0 sum of the ionic radii n average Born exponent BornLand Equation Madelung Constant A 12 8 Each Na is surrounded by 6 Cl39 ions at r0 by 12 Na ions at Era by 8 more Cl ions at xgro and by 6 more Na ions at 2ro A6 3 Table 64 Madelung constants for common lattice types Type of lattice Madelung constant A Sphalcritc ZnS Wurtzitc ZnS Sodium chloride NaCl Cesium chloride CsCl Rutilc TiOg Fluorite CaFZ 1638 1641 1748 1763 2408 2519 BornLand Equation Born Exponent n Table 65 Values for the Born exponent Electronic configuration of ion Born exponent n Examples He 5 Li NC 7 Na Mg2 02 F Ar 9 K Ca 82 31 Cu Kr 10 RH Br Ag X6 12 Cs I Au To calculate the lattice energy of NaCl the value of the Born exponent n is the average of 7 Na and 9 Cl39 so n 8 Lattice Energy of NaCI NA 6023 x 1023 mol391 NAAzze2 1 A1748 n z 1 e 1602 x 103919 0 e0 8854 x 103912 CZm391J391 r0 116 pm 167 pm 283 pm n 8 Lattice Energy Kapustinski Equation 5 U1202gtlt10 vz z 1345kJm011 7 0 7 0 V number of ions in the empirical formula r0 is in units of pm Formation of Ionic Compounds Section 62 Nas 1m 29 9 NaCIs AHoforma on 411 kJmol 1 Nas e Nag 2 1127291 9 C9 3 Nag 9 Nag equot 4 Clg equot 9 C39g 5 Nag C39g 9 NaCIs AH atomization AH bond 121 kJmol AHoionization AHwelectron affinity 349 AH lattice energy 793 The BornHaber Cycle Section 63 Nag e C1g l 349 k 502 k Nag Clg V Nag C1g Naltggt Clzltggt 1211ltJ Nas C12g 108 k 793 k 4111lt NaCl39s L Figure 63 Born Haber cycle for the formation of sodium chloride AcidsBase Reactions of Oxides Section 77 Acidic 0029 2 NaOHaq 9 Na2003aq H20l Basic 908 2 HN03aCI 9 M9N032aCI H20 oz Free Energy AG of Reaction Correlates Acid or Base Strength Lewis Theory Section 78 Lewis Acid Lewis Base Ag 2NH3 9 H3NAgNH3 Pearson HardSoft AcidBase HSAB Concepts Section 79 Pearson s Principle Hard Acids Most metal ions Soft Acids Lower right portion of the metallic elements Borderline Acids border between soft and hard acids Hard acids Borderline acids Soft acids H Li Na K 1362 Mg2g Ca2 Srz BF BC13 BOR3 BCH33 BH3 T1 T1CH33 A1 A1CH33 A1C13 AlH3 Cr Mn Fe C03 Fe Co Ni Cu Zn2 Cu Ag Au Cd Hg22 Rh Ir Ru Os2 Hg CH3Hg C0CN53 Pd2 139 Pt4 Ions with oxidation states of Br2 12 4 or higher Metals with zero oxidation state HX hydrogen bonding 7t acceptors trinitrobenzene molecules quinones tetracyanoethylene etc SOURCE Adapted from R G Pearson J Chem Edna 1968 45 581 Pearson HardSoft AcidBase HSAB Hard Bases Soft Bases Borderline Bases Hard bases Borderline bases Soft bases H Fquot Cl Br 1 H20 OH 02 HZS HS 8 ROH RO R20 CH3COO39 RSH RS R28 NO3 CIOAf N02 N3 SCN39 CN RNC CO CO32 SO42 PO4339 SO32 8203 NH3 RNHz NZH4 C6H5NH2 CSHSN N2 R3P RO3P R3As C2H4 C6H6 SOURCE Adapted from R G Pearson J Chem Educ 1968 45 581 Applications of the HSAB Concept Section 710 oz Prediction of Chemical Reactions HgF2g Be2g 9 BeF2g Hg2g AgBrS 39g 9 Aglg Br39g OHquot CHngso ltgt CH3HgOH 032quot CH3HgF H803quot ltgt CH3H980339 HF CH3HgOH H803quot CH3HgSO3 HOH Applications of the HSAB Concept Section 710 oz Qualitative Analysis of Cations Qualitative analysis separation Group 1 Group 2 Group 3 Group 4 Group 5 HSAB Acids Soft Borderline and Soft Borderline Hard Hard Reagent HCl HZS acidic HZS basic NH42CO3 Soluble Precipitated AgCl HgS MnS CaCO3 Na PbCl2 CdS FeS SrCO3 K HgZCI2 CuS COS BaCO3 NH4 SnS NiS AS253 ZnS SbZS3 AlOH3 BiZS3 CrOH3 Precipitates All groups Solutions HZS acidic Group 1 H28 basic Group 2 NH42CO3 Group 3 Group 4 Soluble Group 5 Na m Mg Biological Aspects Section 711 A1 K Ca Sc Ti Cr Mn Fe Co Ni Cu Zn Rb Sr Zr Nb Mo TC Ru Rh Pd Ag Cs Ba Lu Hf Ta W Re Os Ir Pt Au Fr Ra Lr Rf Db 3g Bh Hs Mt Uun Uuu Uub Ga Sb Bi Po La Ce Pr Nd Pm Sm Eu Gd Tb Dy H0 Er Tm Yb Ac Th Np Pu Am Cm Bk Cf Es Fm Md Fundamental Properties of Group 141Elements Carbon Silicon Germanium Tm Lead Symbol C Si Ge Sn Pb Atomic number 6 14 32 513 82 Natural isotopes 12C9889 28819221 76Ola2052 Sn1430 2041M 1 48 A abundances g11 1 39314711 7zGe2743 7511 761 2 Pb236 Si309 Gs776 31112403 2m1132213 MGoa 1654 98118521 web523 7661a 776 08113285 132511 472 124811594 Total no of isotopes 7 8 141 31 21 Atomic weight 1201 2809 7259 1187 2072 Valence electrons 2522192 3323132 4124112 5135122 6536172 mpbp C 3570sublimes 14142355 9372830 23252270 3281750 Density gcm3 225c 233 532 730quot 1 135 Atomicradius 139 158 175 metallic 21 Covalent radius 9 077 117 122 141 147 ShannonPrawn A Ionic radius 02594 06104 31376 4 08336 4 03796 ShannonPrewitt A 2 1228 2 l336d ON Pauling EN 25 18 18 18 19 Charge density 14 10 60 4 48 4 51 chargeionic radius 2 16 2 15 unit charge A 13 V 021 m 091 00 011 0679 MO2 a M acid soln Oxidationstates 41014 4 2 4 4 24 4 2 4 2 4 Ionization energy 1886 786 760 7119 716 ldmol Electron af nity 123 120 118 121 1111 ldmol Discovered by date Antiquity Be rzeiius Winkler Antiquity Antiquity 1824 1886 39 1pr 02 C0 C02 8102 G802 51102 PbO Acid base character Acid Acid Ampho Ampho Ampho of oxide rpw N2 None 8131514 None 8113114 None rpw halogens CX4 SiX4 GCX4 SnX4 PbXZ rpw hydrogen CH 4 None None None None Crystal structure Diamond Diamond Diamond Tetragonal fcc graphite Group Trends Section 141 U nCommon Features Table 141 Melting and boiling points of the Group 14 e Element Melting point C Boiling point C C Sublimes at 4100 Si 1420 3280 G6 945 2850 Sn 232 2623 Pb 327 1751 ot Melting and Boiling Points Free energy Vmol e toOxidation States l I l 1 4 2 0 m2 4 Oxidation state Carbon Section 142 Allotropes of Carbon otoDiamond otoGraphite E in Q9060 4 ltiw lt quot is A 1 1 n L 391 at r r J x f mg a t t 3 4139 I otoNanotubes I L 1 I tube The Extensive Chemistry of Carbon Section 144 ozoCatenation otoMultiple Bonding Carbon Monoxide Carbon Dioxide Carbon Disulfide Sections 146 147 1410 ozoCarbon Monoxide Carbon Dioxide ozoCarbon Disulfide Hydrogen Carbonate Carbonate and Cyanide Sections 149 and 1415 oonydrogen Carbonates ozoCarbonates ozoCyanides Carbon Tetrahalides Chlorofluorocarbons and Methane Sections 1412 1414 ozoCarbon Tetrahalides ozoChlorofluorocarbons ozoMethane Silicon Section 14 16 otoPreparation and Properties Heating coil Pure Si Molten zone Very pure Si Figure 1418 Zone refining method for the purification of silicon otoApplications ozosnica SiOz topndn Bonding Silicon Dioxide Section 1417 H3C Figure 14207 Overlap of a full oxygen p orbital with an empty silicon d orbital 39 T WSiH3 CI3 31V Silicates Section 1419 Om oxygen on silicon j I 0635 AVA a b C FIGURE 1513 Silicate structures a ortho SiOlf b pyro 20675 c cyclic 5139302 d cyclic Si5018 e single chain Si0 f doublechain 403 and g sheet Sigo k The dashed lines indicate the repeating units 12 Zeolites and Glass Sections 1420 8 1418 ozoZeolites Supercaga Cubic cage Sodalite cage otoGlass Germanium Tin and Lead Section 1423 0 3 PbOZ LIA C 2 I 0 1 53 8114 2 amp 5 o Pb SnPb E a 71 Sn2 lt1 4 2 0 2 4 Oxidation state Figure 1431 Frost diagram for tin and lead Germanium Tin and Lead Oxides and Halides Sections 1424 1425 ozoOxides oonalides Biological Aspects Section 1427 ozoCarbon Cycle ozoSilicon otoTin Toxicity ot Lead Poisoning Carbon Reaction Flowchart Section 1428 Nazcz H20 c1Z HCOOH sz qsz cc14 COCIZ H2804 C 12 02 V 02 5 X502 0 HO 39 O NiCO4 N CO 2 C02 2 CH4 2 A14C3 HZ s CaOH2 HCl NH3 CH3OH cos CaCO3 HCN COZ A OH H CaHCO32 CNi Silicon Reaction Flowchart NaZSiO3 NaOH C HCI SiHC13 1 Si 3102 SC CH3C1 HF IE ZCQ H 120727 8113627 An Overview of the Periodic Table Chapter 2 Triads Law of Octaves Mendeleev and Meyer Flaws in Mendeleev s Periodic Table Moseley Organization of the Modern Periodic Table Section 21 2Elements placed in order of increasing atomic number of protons Mainnmup Elemems r A I Farina m I 2 IIA ymLmI Alomic wighl Almmc number s Malnr mup Eiemems z IVA DEM TmnsiuonMcmls n 9 3 Na Mg 3 4 5 a 7 a vuIE m n 71W W ma VB vs vna vua A m 1 2 1 u u 14 5 m 1 m u 4 x Ca Sr n v Cr m r to Ni Cu wnw mm mm 41 mm m uwm um 51mm mm M w u u a w s n s v y m n m m m M cd mm m WW my uzuwx m Imm Vamm sz mm um mm V7 72 74 7 M 1 m w m m o m L1 m w m m n n M u n 117117 my um mm m 1mm m mm mam MW 1mm mun 7 i m w m m w W nu m m m 7 Fr 11 A m m s m u rm um mm m mm mm mm mm mm mm mm m mm mm m m MLhui su m M m a a u 1 m m m 74 n yr m pm s L m n u m u m n L mm MA mm mm mm 71 WW may wwnz um mung mm luau Mcumaxd w 4 m as 1 a M Mr H m nu 1 n u N1 n An Cm an c a n mu m Lr 11mm guns mm 37 mu m mm any my my 2m 11 5 onmuux Organization of the Modern Periodic Table Terminolgy Period Group or Family Main Group Elements Alkali Metals Alkaline Earth Metals Pnicogens Chalcogens Halogens Noble Gases Transition Metals Lanthanoids and Actinoids Existence of the Elements Section 22 Big Bang Theory Exothermic Nuclear Reactions Endothermic Nuclear Reactions Stability of the Elements and Their Isotopes Section 23 20nly 81 stable elements 130 Alpha emlsmon y Why 100 Bela emission Band of slahility Electron capture and positron emission Number of neutrons N 01 e 0 IO 20 30 40 50 60 70 80 90 100 Number of protons Z Quantum or Shell Model of the Nucleus 2Quantum or Shell Model of the Nucleus Magic Number For Electrons 2 10 18 36 54 86 For Protons 2 8 20 28 50 82 114 For Neutrons 2 8 20 28 50 82 126 Examples a rm 39 aims Classification of the Elements Section 24 Standard Temperature and Pressure STP Standard Ambient Temperature and Pressure SATP Criteria for Classification of Metals and Nonmetals 1 Luster 2 Density 3 Hardness 4 Malleability or Ductility 5 Thermal Conductivity 6 Three Dimensional Electrical Conductivity Semimetals B Si Ge As and Te Weak metals Be Al Zn Ga Sn Sb Bi Po Periodic Properties Sections 25 27 1 Effective Nuclear Charge zeff 2 Atomic Radius 3 Ionization Energy 4 Electron Af nity zeff Ionization Energy Atomic Radius Electron Affinity INCREASES INCREASES Shielding or Screening and 237 Section 25 2 Innercore electrons act as a shield for outercore electrons farther out from the nucleus Effective Nuclear Charge Zen ncreases going left to right across a period Decreases going down a group INCREASE Zen T immense mums mnesz INCREASE Zen Slater s Rules Section 25 zeff Z 6 Rules for determining 039 1 The electronic structure of the atom is written in groupings as follows 13 23 2p 3s 3p 3d 4s 4p 4d 42 5s 5p etc 2 Electrons in higher groups to the right in the list above do not shield those in lower groups they contribute zero 3 For ns or np valence electrons a Other electrons in the same ns np group contribute 035 b Electrons in the n 1 group contribute 085 c Electrons in the n 2 group contribute 100 4 For nd or nf valence electrons a Other electrons in the same nd or nf group contribute 035 b Electrons in all the other groups to the left contribute 100 Calculate Ze for the outermost electron in oxygen Calculate Ze for the 3d and 4s electrons of nickel Periodic Properties Atomic Radius Section 25 What is Atomic Size 1 Covalent radius rm 2 van der Waals radius rvdw 3 metallic radius M 1 2 Covalent Radius Covalent Radius F 64 pm Li 134 pm Cl 99 pm Be 106 pm Br 114 pm B 88 pm I 133 pm C 77 pm N 70 pm 0 66 pm Atomic Radius and Relativistic Effects ozoVariations in Atomic Radius Aluminum rCOV 126 pm and Galium rCOV 126 pm ozoRelativisitic Effects in Heavier Elements Period 6 When Z 1 velocity of electrons are 73 x 10393 When Z 80 velocity of electrons are 058 Therefore electrons are 20 heavier when Z 80 1 Relativistic Contraction of s orbitals 2 Relativistic Expansion of dand f orbitals Overall Atomic Radius Decreases Cr 129pm M0 140 pm W 141 pm Periodic Properties Ionization Energy Section 26 Ionization Enery IE First Ionization Energies IE1 energyE1 Li 9 Li equot IE1 520 kJmol energyE1 Ne 9 Ne equot IE1 2080 kJmol Factors 1 IE increases going left to right across a period because 2 IE decreases going down a group because INCREASE IE1 I INCREASE IE1 in ame T ease Increase Periodic Properties Ionization Energy Cont ozoVariations in IE1 Li and Be B 0 Ga Ianizalicn energy kJmal 0 w mere V 3A M 65 Em Io 6A I Periodic Properties 1quot 2quotquot 3m Ionization Energies IE1 lt IE2 lt IE3 lt Table 83 Successive Ionization Energies of the First Ten Elements kJmol39 Element First Second Third Fourth Fi 39h Sixth Seventh H l 3 1 2 He 2372 5250 11815 47276 Periodic Properties Electron Affinity Section 27 Electron Affinity EA ozoLarger values of EA easier to add an electron F e 9 F EA 328 kJmol ozoSmaller values of EA more difficult to add an electron Na e 9 Na EA 53 kJmol otoFormation of Naquot is energetically preferred to forming Nal Periodic Properties Electron Affinity Cont ozoVariations in IE1 Li Be B C N O F Ne EA 2 1 4 7 141 28 kJmol 60 ve 6 5 3 ve Periodicity Examples Which pair of elements is most similar in their chemical properties CaampBa SampP AgampRb CsampBa Which one of the following should have the lowest 1st ionization energy As K Cl Na Be Which element has the greatest atomic radius AI Si P S CI Which of the following atoms will have the greatest electron affinity S P Ga Li Br Biochemistry of the Elements Section 28 Medicinal Chemistry p 26 in text Antacids Tellurium Lithium Boron Bismuth Platinum Gold Bioinorganic Chemistry Essential Elements H Na K Mg Ca Fe Cu Zn C N O P 8 Cl Ultratrace Elements V C Mo Mn Co Ni B Si Sn Se F Bertrand s Rule s and p Orbital Mixing t Orbitals with similar but not equal energies can interact if they have similar symmetry t The 03923 and 62p orbitals both have 039 symmetry and are symmetric to inversion gerade They interact or mix The 620 MO is raised and the 03925 39 G M0 IS lowered 1 o u I I 117gquot 111 l n gt1 nggtllt 8 l I I 39 I I I I I I 1 I l 2 2 P FP 217 2p I z x z 439 I I 6 x x g I I l 7T 75 TCquot TC I l 6g gtllt gtxlt 6 cu 2 2s I 5 2s ZS 6 g z G g h No mixmg Mixing of 5g orbitals M O Diagrams of Heteronuclear Diatomic Molecules l 39 I 239 r 2 I I I x x x x x I lquot 1 l x x 139 x 1 v x X quot l x x S l 39 X x x I x z I x x X x x fquot quotquot 39 x v u A A A A 39 A A A B B A A39 B B Equal energws Unequal energles Very unequal energies 0 10 H 9 3 k 20 Es 21 0 72 E 30 4 5 40 Atomic orbitals MO Diagram of CO Molecular orbitals Atomic orbitals Atomic orbitals MO Diagram of HG Molecular orbitals 1L Nonbonding 1L Atomic orbitals Cl i Epy 3pz Fundamental Properties of Alkaline Earth Metals Beryllium Magnesium Calcium Strontium Barium Radium Symbol Be Mg Ca Sr Ba Ra Atomic number 4 12 20 38 56 88 Natural isotopes 9100 247839 409695 84056 130 0101 226 100 A abundances 25 1000 42 0646 86986 1320097 261101 430135 87702 134242 44208 888256 135659 460186 136781 48018 1371132 1387166 Total no of isotopes 6 8 14 16 39 2O 16 Atomic weight 9012 2431 4008 8762 1373 226 Valence electrons 252 352 452 Ss 2 632 752 mpbp C 12832970 6501120 8451420 7701380 7251640 7001140 Density gcm3 185 174 155 260 351 5 Atomic radius A 112 160 197 215 224 logic radius Shannon1 rewitt A CN 0414 0714 l146 1326 1496 Pauling EN 15 12 10 10 09 09 Charge density 49 28 1 7 1 5 13 charge ionico radius uni charge A 39 E a V 185 237 287 h 289 290 292 Oxidation states 2 2 2 2 2 2 Ionization energy k mol 899 738 590 549 503 Estimated electron 241 230 154 120 52 af nity k1 mol Discovered bydate W6h1erBussy Davy Davy Davy Davy Curie 1828 1808 1808 1808 1808 191 1 rpwb 02 B120 MgO CaO SrO SrOz 131102 R510 Acid base character Ampho Weak base Base Base Base Base of oxide rpw N2 None Mg3N3 21ng2 Sr3N2 BagNg Ra3N2 rpw halogens 86X 2 MgX 2 CaX z SrX 2 Bax l RaX 2 rpw hydrogen None MgH 2 Call 3 SrH 2 BaH 2 Crystal structure hop hep cop cop bcc Group Trends and Common Features of Alkaline Earth Metal Compounds Sections 121 122 ozoCommon Features 75 pm 116 pm 152 pm 166 pm m m vChemIcal ReactIVIty w 86 pm 114m 132 pm 49 pm otolonic Character Table 123 Usual hydration number of common Element MCI Mme MSO Mg 6 v 7 I o 4 7 M vlon Hydration Sr 6 4 0 Ba 2 o o Solubility of Alkaline Earth Metal Salts Section 123 2Many Alkaline Earth Metals Salts are Insoluble AG AH TAS Table 126 Calculated nee energy changes luv the solution process fol magneslum chloride and sodium chloride Enlnalpy change Entrupy change Free energy cnange Compound kJ39mol39l Jmolquot mo 1 MgCl2 7133 734 799 Nlel 4 T 1 3 7 1 l Beryllium Section 124 ozoPreparation and Properties ozoApplications Magnesium Section 125 ozoPreparation and Properties ozoApplications Calcium and Barium Section 126 ozoPreparation and Properties ozoApplications Oxides and Hydroxides Sections 127 128 Z0xides Z Peroxides Z Hydroxides Table 127 Solubilities of the alkaline earth metal hydroxides Hydroxide SolubeLy g39Lquot Mg 00001 Ca 12 Sr 1 0 Ba 47 Calcium Carbonate and Cement Sections 129 1210 Calcium Chloride and Calcium Sulfate Sections 1211 1212 Calcium Carbide Section 1213 irlti3w gure 123 Crysta structure of ca cium dicarbideae which mosey resembles the sodium ch onde crysta structure Biological Aspects Section 1214 otoMagnesium ozoCalcium Magnesium Reaction Flowchart Section 1215 quot12353 H 39 39 H L 3033 39CEHEMgEr I39vng h MEL3 Calcium Reaction Flowchart CEiENE H 21 V H 3ng 4 a C MOHM ii cag DH39 CDT 39quot CD CL I j 0 up 3 JUL G39CLC L L CE quot 4 139 m 1 339 1130 1 L 5 139 Hgm Barium Reaction Flowchart BEHNE N v El 1 139 H r 504 Bach Ba quot39 BE WOHM i 3213 8304 Of 1 Covalent Bonding Chapter 3 Foundation for Modern Chemical Bonding Theory ozoLewis Theory ozoValence Bond VB Theory ozoMolecular Orbital MO Theory A Brief Review of Lewis Theory Section 36 Rules 1 Count total number of valence electrons present polyatomic anions add the of negative charges polyatomic cations subtract the of positive charges 2 Write the skeletal structure of the compound least electronegative atom occupies center position H and F usually occupy terminal end positions 3 Draw single covalent bond between all bonded atoms 4 Complete the octets of all atoms bonded to central atom with the remaining valence electrons lonepair electrons H holds only 2 valence electrons 5 If the octet rule is not satisfied for the central atom try adding double or triple bonds A dditional Rules for Writing Lewis Structures Number of valence electrons an atom has is equal to its Periodic Table Group Number IA IIA IIIA IVA VA VIA VIIA VIIIA Therefore Li has 1 val e Be has 2 val e B has 3 val e Al has 3 val e C has 4 val e N has 5 val e O has 6 val e F has 7 val e Ne has 8 val e Moreover H has 1 val e H atoms only form 1 bond with other atoms H atoms hold a maximum of2 electrons in its valence shell Most other elements can hold a maximum of 8 electrons in its valence shell Octet Rule Halogens F Cl Br I tend to form only one bond with other atoms except in XOn or XYn structures where X and Y are halogens Oxygen atoms tend to form only 1 or 2 bonds a single bond two single bonds or one double bond with other atoms One exception is ll30 Triple bonds are possible Carbon atoms tend to form four bonds with other atoms In a molecule with the formula OXn where X is any element all the O atoms are bonded to X Ony C N O and S ordinarily form multiple bonds double or triple bonds Remember to place brackets around Lewis structures that have an overall positive or negative charge If a formula is written this way NPO then the Lewis structure will tend to be drawn in the same order P is the central atom with O bonded to P and N bonded P Procedure to calculate of bonds in a Lewis structure using example of 002 Won t work every time 1 Assume each atom has an octet usually 8 electrons but H can only hold 2 electrons and add up total electrons possible C holds 8 electrons 0 holds 8 electrons total electrons possible 8 28 24 total e 2 Calculate of valence electrons 426 16 valence e 3 of bonds 12tota e39 valence e39 For 002 of bonds 1224 total e 16 valence e 128 4 bonds in 002 Lewis structure has 2 double bonds 4 total bonds OCO Writing Lewis Structures Examples NF3 COCIZ C02 302 Nze Exceptions to the Octet Rule 1 Lewis structures containing atoms that do not have 8 valence electrons 1 Incomplete octetH He Be B AI AlClg H quot BeHz gi Al CI 9 2 Expanded octet elements with n 3 only PCI5 239 Partial Bond Order Section 37 Bond Order Resonance 0032 N03 Formal Charge Section 38 Calculating Formal Charge of an Atom using Lewis Structures Formal Of valqnce of nonbonding 12 of bonding electrons In free Charge atom lone paIr electrons electrons Example C032quot Formal Charge Cont Rules For Determining Most Plausible Lewis Structure 1 Smaller formal charges either or are preferred over larger formal charges 2 Like charges on adjacent atoms are not preferred 3 Negative formal charge prefer to reside on more electronegative atoms Example Using formal charges determine the most preferred Lewis structure for the cyanate ion NCO Valence Shell Electron Pair Repulsion Theory Section 39 How to predict the shape of molecules 3D arrangement of atoms in a molecule ls H20 linear or bent ls NH3 planar or nonplanar ValenceShell ElectronPair Repulsion Theory VSEPR otoAssumption electron pairs in the valence shell of an atom repel each other Bondingpair bp electrons electrons involved in bonding Lonepair Ip electrons electrons not involved in bonding otoMolecule adopts a geometry around the central atom which minimizes the repulsive forces between bondingpair and lonepair electrons Maximize the distance between these bp and lp electrons Arrangement of Electron Pairs Around a Central Atom E Pe Geometry 2 Linear T genal Planar 4 Tetrahedral Arrangement of Electron Pairs Around a Central Atom EPs Geometry T gonal Bipyramidal 6 Octahedral Applying VSEPR 1 Draw the Lewis structure of the molecule 2 Count the number of electron pairs around the central atom Multiple bonds double and triple bonds count as one electron pair 3 Determine the molecular geometry based on the number of lonepair lp and bondingpair bp electrons otoln molecules with both lp and bp electrons the repulsions decrease in the order less favorable Ip Ip gt Ip bp gt bp bp more favorable ie minimize the number of close lp Ip interactions Applying VSEPR Five Electron Pair Example 391 F t 5 regions at electron density around the 3 31 central etem Cl I It i Trigonal Bipyramidal Arrangement Elli quotif of Electron Pairs There are three peeeible structures In Iquot F I we F I fn39t39F I ee F 391 cr F Cr 7 F 7 1 K 43 F F A E C 90 lp lp 0 1 0 90 lp bp 6 3 4 Tshaped Geometry Applying VSEPR Six Electron Pair Example KEF E regime uf electrun de neity around the 4 g fig central aturn Cl Iii tir 1 Octahedral Arrangement of 3 H 139 Electron Pairs There are two pessible structures Iii ill Fax Wall 13 39RVF F 1 F I A E Geometry Square Planar Summary of ElectronPair Arrangements and Geometries Total of of of Electron Bonding Lone ElectronPair Molecular Pairs Pairs Pairs Arrangement Geometry Example 2 2 0 linear linear 3 3 0 trigonal planar trigonal planar 3 2 1 trigonal planar bent 4 4 0 tetrahedral tetrahedral 4 3 1 tetrahedral trigonal pyramidal 4 2 2 tetrahedral bent 5 5 0 trigonal bipyramidal trigonal bipyramidal 5 4 1 trigonal bipyramidal see saw 5 3 2 trigonal bipyramidal Tshaped 5 2 3 trigonal bipyramidal linear 6 6 0 octahedral octahedral 6 5 1 octahedral square pyramidal 6 4 2 octahedral square planar Summary of Selected Molecular Geometries A bent I I axe393 a l X X 1quot A a K 339 lt7 X trigclnal pg an dal 939 H l 39 12 A H T shaped H 6 9n39 39 K I K 12 45 EVA seesaw K K lt K J ff A k x XH 90 l 9039 H H 439 K33 square pmn al square planar 9 lt galKt Effect of Lane Pairs on Bond Angles N02 N02 and N02 CH4 NH3 and H20 SF4 IF4 Bonding in Greater Than Six Directions MX7 species Pentagonal bipyramidal Capped trigonal prism Capped octahedron MXs species Square Antiprismatic Valence Bond Theory Section 310 Valence Bond Theory assumes that covalent bonds are formed when atomic orbitals on different atoms overlap and the pair of electrons is shared Valence Bond Theory Cont How do we explain the bonding of carbon in methane Hybridization 2p Energy 1s C atom ground state N Is C atom promoted Hybridization Hybrid Orbitals In methane CH4 the one 25 orbital and the three 2p orbitals hybridize into four sp3 hybrid orbitals Kinds of Hybrid Orbitals You can describe the bonding around the central atom in a molecule if you first determine the geometry of the electron pairs around the central atom Table 1 02 Kinds of Hybrid Orbitals Hybrid Geometric Number of Orbitals Arrangement Orbitals Example sp Linear 2 Be in BeF2 sp2 Trigonal planar 3 B in BF3 sp3 Tetrahedral 4 C in CH4 sp3d Trigonal bipyramidal 5 P in PC15 sp3d2 Octahedral 6 s in SF6 Spatial Arrangements of Hybrid Orbitals Linear arrangement Trigonal planar arrangement 39I etrahedral arrangement sp hybrid orbitals rpz hybrid orbitals Sp3 hybrid orbitals Trigonal bipyxainidal arrangement Octahedml arrangement sp3d hybrid orbitals sp3d2 hybrid orbitals 1 BHe2822p1 FHe2822p5 Hybridization of NF 3 Hybrid Orbital Formation in BF3 2p 23 2p 2p sp2 23 2P sp2 sigma and pi Bonding sigma bond pi bond Hybridization of 002 c He 232 2p2 0 He 2s2 2p4 Hybrid Orbital Formation in C02 2P 2P 2P 2P sp sp 23 23 Today Monday Jan 13 quot and FYI ozolntroductions Student Information ozoCourse Syllabus Start Lecture on Chapter 1 What is Inorganic Chemistry Schrodinger Wave Equation Section 11 Quantum Numbers Section 12 Shapes of the Atomic Orbitals Section 12 oz For Lecture on Wednesday Jan 15th Downoad and print Chapter 1 Lecture Notes Read pages xx xxvii and Applications of Inorganic Chemistry colored slides between pages 260 and 261 Read Chapter1 in textbook Sections 11 15 pages 1 15 Do Assigned Homework Problems in Chapter 1 1 7 Finish Lecture on Chapter 1 I will announce Problem Set 1 which is due this Friday Jan 17th Writing Electron Configurations More Examples What is the electron configuration for Pb using the noblegas core Pb Xe 632 4f14 5d1O 6p2 What is the electron configuration for Au using the noblegas core Au Xe 632 4f14 5d9 9 Au Xe 631 4f14 5d1O What is the electron configuration for Cu using the noblegas core Cu Ar 4s2 3d9 a Cu Ar 4s1 3d10 What is the electron configuration for W using the noblegas core W Xe 632 42 14 5d4 not an exception What is the electron configuration for Fe using the noblegas core Fe Ar 4s2 see Mo Nizquot Ion Electron Configurations More Examples Clquot Ne 332 3p5 Clquot Ne 332 3p6 or Ar Pb Xe 632 4f14 5d1O 6p2 Pb Xe 632 4f14 5d1O 6p Fe Ar 4s2 3d6 Fe2 Ar 4303616 Fe3 Ar 430 3amp5 Mo Kr 5s1 4d5 an exception Mo Kr 5304cO or Kr Fe Ar 4s2 3d8 Fe Ar 4s0 3d8 Fundamental Properties of Group 16 Elements Oxygen Sulfur Selenium Tellurium Polonium Symbol 0 S Se Te Po Atomic number 8 16 34 52 84 Natural isotopes 1609976 3239511 quotSe087 120Teo089 210mm AW abundances 1700037 338076 Se902 122Tezms 1800204 3434212 77Se758 123quotre087 3630014 78Se2352 12quotquot1 e4n 1 8 8e4982 SR699 82Se919 126Teism 138 Te 3 1 79 Total no of Isotopes 8 10 17 21 27 Atomic weight V 1600 3207 7896 1276 210 Valence electrons 23324 35231 4524p4 5525p 6s26p4 mp bp C 218 183 112444 217685 450990 254962 Density g cm3 a l43g1iter 207b 479 624 932 Covalent radius A 073 102 116 136 Ionic radius 0 39 39 1436 0566 6 1706 1 6 08116 Shannon Prewitt A CN 41116 41086 Pauling EN 35 25 24 21 20 Charge density charge o 14 11 686 674 ionic radius unit charge A 43 60 103170 E0 V 123 014 011 069 10 Oxidationstates w1 2 2to 6 2t0 6 2to 6 Zto 6 Ionization energy kJmoi 1314 1000 941 870 814 Electron af nity kJ mol m 142 200 e 195 190 NA Discovered by date Priestley Antiquity Berzelius Miiller Curie 1774 1817 1782 1898 1 de 02 02 802 T602 P002 Acid Base character of oxide Acidic Acidic Ampho rpw N2 N0 NO2 None None None None rpw halogens 02 F2 51743135 Sex 4 Tex4 FOX 4 8262312 Ser TeF6 PoC12PoBr2 82 BIZ Se 20233213133 rpw hydrogen H 20 H ZS 11286 None None Crystal structure Cubic Orthoi Hexag Hexag Cubic Group Trends Section 161 o UnCommon Features Chalcogens toOxidation States Table 161 Melting and boiling points of the Group 16 e toMelting and Boiling Points Element Melting point C Boiling point C 02 219 183 88 119 445 ch v 221 685 TC 452 987 P0 254 962 Anomalous Nature of Oxygen Section 162 oonniqueness Principle oonigh Stability of Multiple Bonds ozoLack of Catenated Compounds Oxygen Section 163 otoDioxygen ozoOzone Bonding in Oxygen Compounds Section 164 ozoCommon Bonding 3 Orbital Mixing otoBent s Rule More electronegative substituents prefer hybrid orbitals with less 3 character and more electropositive substituents prefer hybrid orbitals with more 3 character Trends in Oxides Section 165 otoTransition in Bonding Type Ionic to Covalent Table 99 Formulas bonding types and phases at room temperature of the highest oxides of the Periods 2 and 3 elements Compound Li20 BeO B203 002 N205 F20 Bonding Ionic Ionic Network Covalent Covalent 39 Covalent type phase solid solid covalent gas gas gas solid compound N320 A1203 P40 10 C1207 Bonding Ionic Ionic Ionic Network Covalent Covalent Covalent type phase solid solid solid covalent solid solid liquid solid otoAcidBase Behavior H20 H202 and OH39 Sections 167 169 ozoWater oonydrogen Peroxide oz Hyd roxides Sulfur and Its Allotropes Sections 161 1 1612 Sulfur otoAllotropes of Sulfur Free energy V39mol e sof Acidic solution A Sl H28 Basic solution SO32 d 2 so4 I I I I I 6 4 2 0 Z Oxidation state Sulfides and Hydrogen Sulfides Sections 1614 1615 ozoSulfides 00 Hydrogen Sulfide Sulfur Oxides Section 1 61 6 t Sulfur Dioxide ozoSulfur Trioxide Sulfuric Acid Section 161 7 1618 ozoProperties otoAcid ozoDehydrating Agent ozoOxidizing Agent ozoSulfonating Agent oz Base Sulfites and Sulfates Sections 1619 1623 ozoSulfites ozoSulfates oonyd rogen Sulfates oonhiosulfates and Peroxydisulfates Summary of the OX0 Acids of Sulfur Name Formula Structure l Acids Containing One Sulfur Atom Sulfurous H2503 SO in sul tes 3 Sulfuric H2504 O OH OH Acids Containing Two Sulfur Atoms O I Thiosulfuric H25203 O SH OH I F Dithionous H28204 HOS S OH 0 1 Disulfurousc H28205 HO S S OH O O I Dithionic H28206 HO yhlS OH O 0 Cl 39Disulfuric H28207 HO39S O ISOH O O Acids Containing Three or More Sulfur Atoms CF I Polythionic H28 206 HO ISHSn IS OH O O Peroxo Acids I Peroxomonosulfuric H2805 HOD f OH O I O Peroxodisulfuric H28208 Sulfur Halides Sections 1624 1625 ozoSulfur Hexafluoride ozoSulfur Tetrafluoride otoDisulfur Dichloride and Sulfur Dichloride Selenium Tellurium amp Polonium Section 1627 Oxygen and Sulfur Reaction Flowchart Section 1627 H20 KC103 L120 HZ Li A 03 02 Na Nazoz MnOZ K C H202 K02 coZ Pbs SF6 stog Hso339 sof Lszof 34062 1313 F2 A H20 7 Q2 02 OZ H H H 2 H28 quot39 38 M 802 V205 H2804 3804 H C12 H O v t e JBRN H SO 2 r Nags 52012 2 L 7 820827 82804 Cl C L Na3804 31 Chapter 3 Covalent Bonding Lewis Formulas The interaction between two H atoms to form a covalent bond can be represented as shown below H39 H gt HH or H H The pair of electrons between the two atoms are said to be shared by the atoms and may be represented by a dash Here each h atom may be considered to be associated with two electrons giving each H atom a He electron con guration The interaction between H and Cl may be represented in a similar manner H gt H 1 or H 39C391 Here each atom again has a noble gas con guration H has a He con guration and Cl has an Ar con guration These representations of the H2 and HCl molecules are referred to as Lewis electron dot formulas or simply Lewis structures In Lewis structures electron pairs are represented by pairs of dots or dashes Bonding pairs shared pairs are those electrons that are shared by two atoms These are represented by a dash Lone pairs unshared pairs are the electrons which are associated 4 H 39C HCH with only one atom These are represented by a pair of dots Atoms of nonmetals tend to achieve a noble gas con guration by the sharing of electrons with each other Apart from hydrogen 3 H 39 T 39 39 H N H and a few other exceptions we will mention later atoms will tend H to be associated with 8 valence electrons 2 H O gt H g H This tendency for atoms other than H to acquire a total of eight valence electrons by the sharing of electrons is known as the Octet Rule H F gt H PT O en the number of covalent bonds formed by an atom corresponds to the number of single electrons shown in its Lewis symbol 32 Coordinate Covalent Bonds In socalled normal covalent bonds each atom supplies one of the electrons being shared There is another type of covalent bond in which one atom provides both electrons that are share This is known as a coordinate covalent or dative bond H I H H lII H gt H llH H H Once formed the coordinate covalent bond is indistinguishable from other covalent bonds All of the bonds in the ammonium ion above for example are identical in all respects Multiple Bonds Sometimes atoms cannot achieve an octet by the sharing of one pair of electrons There are some instances in which two or even three pairs of electrons must be shared for two atoms to each achieve an octet RV N N gt N EI gtNEN Double Bond A double bond is a covalent bond in which two electron pairs are shared by two atoms The atoms C N O and to a lesser extent S and P often have double bonds Triple Bond A triple bond is a covalent bond in which three electron pairs are shared Triple bonds are usually limited to C and N 3 3 Exceptions to the Octet Rule 1 Incomplete Octet An incomplete octet occurs when an atom has fewer than 8 valence electrons associated with it Incomplete octets are usually encountered in molecules with Be B or Al atoms For example 39 quot 6valenoe ebctiors Cl Be Eli 4va1eme electrors a101er B u quot arourrl Be However these atoms can achieve an octet by the formation of a coordinate covalent bond As shown in the example below for B 2 Expanded Octets An expanded octet involves more than 8 valence electrons being around an atom Expanded octets are not uncommon for atoms in the third period and beyond beginning with Si 516 XeF4 PF5 u ESE XRE quot 3 4 Electronegativity and Bond Polality When electrons are shared between two identical atoms the electrons are shared equally For example quot quot I when two H atoms are bonded together or two Cl HH 91 91 I atoms are bonded together there is equal likelihood of L nding the electrons near one nucleus or the other equal Sl aIing of electrons However when a bond is formed between atoms of two di erent elements the bonding electrons will not necessa1ily be shared equally A polar covalent bond is a covalent bond where the electrons are not shared equally but are more likely to be found nearer the nucleus of one atom than the other When a H atom bonds to either a F or Cl atom the bonding electrons will be more likely to be found near the F or Cl atom than near the H atom As a result the F or Cl atom acquires a partial negative charge while the H acquires a partial positive charge These partial charges are denoted by a lower case delta 6 6 F I l H F Na Eq39 nonpolar covalent polar covalent ionic A polar covalent bond is intermediate between a nonpolar covalent bond and an ionic bond This should also point out that covalent and ionic bonds are merely extreme idealizations on a continuum of bonding How do we determine if a bond is polar and if so which atom will bear the pa1tial negative charge We need a measure of how strongly an atom attracts the electrons in a chemical bond The ELECTRONEGATIVITY of an atom is just such a measure Electronegativity X A measure of the ability of an atom in a molecule to attract bonding electrons to itself 35 Atoms that have a large negative electron attachment enthalpy readily accept electrons and a large positive ionization energy reluctant to give up electrons should be expected to strongly attract bonding electrons to themselves when they are bonded to other atoms Accordingly Mullilmn proposed that the electronegativity of an atom be given by 12 the diiTerence between its ionization energy and its electron attachment enthalpy In other words when the ionization energy and electron attachment enthalpies are expressed in appropriate units eV mAnu T Linus Pauling proposed another more commonly used electronegativity scale shown below It should be kept in mind that electronegativity is a property of an atom in a molecule rather than of an isolated atom like ionization energy or electron af nity Therefore the values shown below cannot be taken too literally Nevertheless they are somewhat use ll Atoms of non metals tend to have higher electronegativities while metals tend to have smaller electronegativities They are said to be electropositive The greater the electronegativity difference between two bonded atoms the more polar the bond and the electrons will tend to be closer to the atom with the larger electronegativity As the diiTerence in electronegativity gets larger an ionic model of the bonding may become more useful 36 Formal Charge Suppose we have a compound whose molecules each contain one C atom one H atom and one N atom We can consider two possible H CEN arrangements of the atoms Why are these the only two possibilities and H NEC One of these is much more stable than the other Using the concept of formal charge allows us to select the more stable of two or more possible structures Formal charge the hypothetical charge that an atom in a molecule acquires when the bonding electrons are assumed to be shared equally between the atoms and lone pairs are assigned to a single atom Formal charge of valence electrons in the isolated atom 7 of non bonding electrons 7 12 of bonding electrons H CEN H NEC H1010 H1010 c4 0 40 c4 2 3 1 N5 2 30 N5041 When deciding on the relative stabilities of two or more possible structures 1 The Lewis structure which has the smallest magnitude of formal charges It is usually desirable to minimize the number of nonzero formal charges 2 When two or more Lewis structures have the same magnitude of formal charge the one that has the negative formal charges on the more electronegative atoms is preferred For a neutral molecule the sum of the formal charges of all the atoms equals 0 For an ion the sum of the formal charges of all the atoms equals the charge of the ion 37 Consider the structures shown below 2 H 2 B IS o 0 0 s 6701282 s 670712120 O 676712271 70 676712271 0 67471240 The second Lewis structure implies that two of the Si 0 bonds should be diiTerent than the other two However experiment shows that all four 8 0 bonds are identical and are somewhat shorter than a typical 8 0 single bond but longer than a 80 double bond In reality the double bonds are not localized between the S atom and two of the O atoms as shown but rather are delocalized over all four 8 0 bonds Delocalized Bonding a type of bonding in which electrons are associated with several atoms rather than being localized between only two The bonding in metals metallic bonding can be viewed as an extreme case of delocalized bonding The valence electrons are delocalized over an entire crystal rather than being localized between pairs of atoms For instance in a crystal of sodium metal the 3s electrons one provided by each Na atom are free to move throughout the entire crystal of what can be regarded as Na ions occupying xed sites The mobility of the valence electrons accounts for the electrical conduction of Na metal For molecules in which delocalized bonding occurs a single Lewis structure cannot adequately describe the bonding in the molecule We use the concept of resonance to describe the bonding in such molecules 38 Consider another example SO3 Ifwe draw a single Lewis structure the implication is that two of the bonds are different than the third This is contrary to experimental data which indicate that all of the bonds are l exactly the same within experimental error Furthermore the bonds are 39 intermediate between a single bond and a double bond To indicate this O we draw a series of resonance structures also called contributing or canonical structures separated by double headed arrows O S OI lt gt lt gt O O O The actual structure of 03 is a weighted average of all the contributing structures It must be stressed that the molecule does not uctuate between these structures An analogy might clarify things A traveler to Africa described a rhinoceros as a cross between a unicom and a centaur That is he was describing something that was real as the hybrid of two imaginary creatures The resonance hybrid is similar It is something that actually exists described as the average of things that do not exist A Resonance structures diiTer only in the allocation of electrons not in the position of the atoms Molecular Geometry Lewis structures indicate how atoms are bonded to each other in molecules but they are not meant to represent the three dimensional structure of molecules Molecules have de nite shapes and these shapes in uence many of the chemical and physical properties of the molecules For example the structures of BF3 and NH3 are somewhat di erent The angle between F atoms in BF3 is 120 and the angle between H atoms in NH3 is about 107 3 15 E HIKIH 120 H 107 Molecular geometry the general shape of a molecule as determined by the relative positions of the atomic nuclei The Valence Shell Electron Pair Repulsion V SEPR Model is a relatively simple model that allows us to predict molecular geometry The basic premise of the VSEPR model is that electron pairs both bonding and nonbonding will be arranged around an atom in a molecule in such a way as to the electron electron repulsions We must rst consider the ELECTRONIC GEOMETREES about the central atom that the repulsions In this graphic we should replace the term Number of pairs with the term Steric Number or Regions of High Electron Density In this context a single unshared electron counts the same as an unshared pair of electrons and a multiple bond counts the same as a single bond This table gives us the arrangement of the regions of high electron density around the central atom which the electronelectron interactions for the given number of regions These are sometimes referred to as the electronic geometry The molecular geometry is partly determined by the electronic geometry 1 Start with the appropriate Lewis structure 2 Determine the appropriate electronic geometry 3 From the electronic geometry and the number of bonds no distinction between single and multiple bonds deduce the molecular geometry The following tables will be extremely helpful m5 2 3 4 Arrangement nl mm Number m pans Arrangement m pans ELECTRON PAIRS ARRANGEMENT MOLECULAR Tulal Banding Lune P GEOMETRv EXAMPLE 2 2 o Lmeav 5 4 Be2 F Ee F Tngona I a a Ma a BF 3 Agtltd E E Tngona 3 p anav Eentov 2 1 angmay 502 AxZ amp 0 o ELECTRON PAIRS ARRANGEMENT MOLECULAR Tmul Banding Lune 0F PAIRS GEOMEI39RY EXAMPLE H Talahedval CH 4 o Ax WI H H Tngona 4 a 1 renamme mama NHa AK Bmov 2 2 angmay H20 AX ELECTRON PAIRS Tulal Bundlng Lune H M 532 H ELECTRON PAIRS Tm Bundlng Luna 6 0 e 5 I 4 2 ARRANGEMENT 0F PAIRS TngonaI pryIamIdaI ARRANGEMENT OF PAIRS DashedvaI Squaw wvamIdaI AX some planav MOLECULAR GEOMETRV MOLECULAR 397Lone paw EXAMPLE iI CI PCIS CI A CI CI E 53 T OIFG F Oquot F gtlteE2 EXAMPLE I EN F SE6 E f r E IEs XeF Steric Number 7 Pentagonal Bipyramidal Capped Trigonal Prism Capped Octahedron B B r l 4 B B B I B I B 53 x B 39 5 x s B B 4 I 39 4 B B Steric Number 8 Cubic Square Antiprism Triangualted Dodecahedron r r BT IBN B 39 P39 z393 A I I i l A l r r IIB I B l B B B next three limited to actinides and lanthanides Hexagonal Bipyramidal Bicapped Trigonal Prism Bicapped Trigonal Antiprism B 39I3 E 4 c B B er A39 B A Pc x B r an39 n I 39ZX Qr r B 3 B 9 B l B Q I 4 B B B Steric Number 9 B B B Tncapped Tngonal Prism I I I I I A I 3173 Fairly common eg ReH9239 and MH209339 of lanthanides B B 313 We can attempt to show the three H bond 39 b hind d1mens1onal structure of molecules herds gomg e on paper by us1ng a combmat1on 39 tl B plare Oftl E page of lines wedges and cross plam Of f page G di erent bonds A normal line indicates a bond that is in the plane of the page A wedge bOI Il out in ont indicates a bond that comes out in of E p1ane of B page front of the page while a cross hatched wedge indicates a bond that goes back behind the plane of the page hatched wedges to represent H quot39H H The VSEPR model allows us to predict approximate bond angles in molecules 1 Lone pairs require more space than bonding pairs Take for example CH4 NH3 and H20 All of these structures are based on a tetrahedral electronic geometry in which the ideal bond angle is 10950 The speci c electronic structure of a molecule will alter these ideal angles somewhat In CH4 where there are no lone pairs the bond angles are all within experimental error 1095O actually 1090 28 the ideal angle for a tetrahedral geometry The angle between nonbonding electrons and bonding electrons will be somewhat larger than that between bonding electrons The angle between two sets of nonbonding electrons will be even larger The presence of one lone pair in NH3 causes the HNH angle to be smaller than the ideal The presence of two lone pairs in H20 causes an even uther reduction in the HO H bond angle from the ideal H I 0quot HunfH H39 H H1H H 10730 10450 10950 2 A single unshared electron will generally require less space than a bonding pair N02 N02 N0239 1 800 1340 1 1 5 3 Multiple bonds require slightly more space than single bonds Take for example formaldehyde HZCO and ethylene C2H4 The ideal bond angles in each of these molecules are 1200 Steric number 3 no lone pairs The angle between the double and single bonds is somewhat larger than the ideal while the angle between the single bonds is somewhat smaller i H H 122 1170 g cC HCjH H H 1160 1215 Dipole Moment and Molecular Geometry We previously talked about the idea of the unequal sharing of electrons between two covalently bonded atoms This unequal sharing leads to a partial positive charge on one atom and a partial negative charge on the other The magnitude of this charge separation can be quanti ed A dipole moment is a quantitative 5 5 measure of the degree of charge separation in a bond or in a molecule HF has anonzero dipole moment Hi HF is a polar molecule F F F2 has no dipole moment quot quot F2 is a nonpolar molecule 5 5 E39 Even though it contaim polarbonds Be F Ber has no dipole moment Ber is a nonpolar molecule The dipole moment is given by u 6d where u dipole moment 6 charge d distance of separation The unit of dipole moment is the Debye D l D 334 X 103930 C m Each polar bond in a molecule can be treated as a vector Two vectors of equal magnitude but opposite direction cancel each other 0 09 quotOquotlt lt 39 gt H H We can also look at this from the standpoint of centers of positive and negative charge Ifthe centers of positive and negative charge are coincident the molecule will be nonpolar On the other hand if they do not coincide the molecule will be polar 6 6 6 oc H H Formula AX AX2 AX3 AX5 AX6 Molecular Geometry Linear Linear Bent Trigonal Planar Trigonal Pyramidal Tshaped Tetrahedral Square Planar Seesaw Trigonal Bipyramidal Square Pyramidal Octahedral Dipole Moment Nonzero Zero Nonzero Zero Nonzero Nonzero Zero Zero Nonzero Zero Nonzero Zero Assuming all X atoms are identical and A and X have di erent electronegatiVities Models of Chemical Bonding I Valence Bond Theory According to Valence Bond Theory a bond form between two atoms when 1 An orbital of one atom occupies the same region of space as an orbital of another atom These orbitals are than said to overlap 2 H A pair of electrons simultaneously occupies both orbitals H2 rlt Because of the overlap of the two orbitals the electrons are simultaneously attracted by both nuclei holding the atoms together The strength of the bond depends on the degree of overlap Orbitals other than s orbitals overlap only in particular directions These directions are such that maximum overlap is obtained Hybrid Orbitals Based on the bonding model we have considered so far it might seem as though an atom will form only as many covalent bonds as it has unpaired electrons in its valence shell F L i JLLL HF ls 2s 2p 0 L i LL ls 2s 2p N L L LLL ls 2s 2p 318 However C which has only two unpaired electrons in its valence shell in its ground state almost always forms four covalent bonds H H CH4 1s 2s 2p We can rationalize this behavior by considering the process of promotion Keep in mind that in the n 2 shell the 2p subshell is higher in energy than the 2s subshell Moving an electron from the 2s to the 2p subshell requires the input of energy 2p lL L L ZS L promotion ZS L We might now expect that three of the bonds in CH4 might be at 900 to each other with the fourth bond being at any arbitrary angle corresponding to the p orbitals which are mutually perpendicular and an s orbital which has no directional character However we know that in CH4 all the bond angles are 10950 and all the bonds are equivalent The three 2p orbitals and the 2s orbital may mix resulting in four new hybrid orbitals which are equivalent to each other in all respects 2p hybridization L L L L sp3 2sl 3 2p orbitals 1 2s orbital 4 sp3 hybrid orbitals In CH4 a H ls orbital overlaps with each of the four sp3 hybrid orbitals resulting in four equivalent covalent bonds Orbitals must be conserved in the hybridization process so of hybrid orbitals formed of atomic orbitals mixed An s orbital can mix with various numbers of p and d orbitals to give various hybrid orbital sets each with its own distinctive arrangement of lobes Atomic orbitals hybridization of hybrid orbitals arrangement 1 s l p sp hybrid 2 linear l s 2 p sp2 hybrid 3 trigonal planar l s 3 p sp3 hybrid 4 tetrahedral l s 3 p l d sp3d hybrid 5 trigonal bipyramidal l s 3 p 2 d sp3d2 hybrid 6 octahedral These arrangements should look familiar to you These correspond exactly to the electronic geometries used in the VSEPR model Therefore if we know the electronic geometry of an atom in a molecule or ion we can infer its hybridization 1 Write the Lewis structure 2 Use VSEPR to predict electronic geometry 3 From the electronic geometry deduce the hybridization Multiple Bonding One hybrid orbital is required for each bond to another atom in a molecule and for each lone pair H H C H H Consider ethylene C2H4 Here each C atom is bonded to one C atom and two H atoms with no lone pairs Therefore 3 hybrid orbitals are required and sp2 hybridization is indicated Altematively we could rationalize that since the steric number for each C atom is 3 the electronic geometry is trigonal planar and the hybridization is therefore sp2 This leaves an unhybridized p orbital on each C atom perpendicular to the plane of the molecule We now need to consider two different types of bonds Sigma 0 bond A sigma bond has cylindrical symmetry about the bond axis and results from the headtohead overlap of two orbitals GD 0 9 9 GO 0CD 401 9 9 One component of the C C double bond is a sigma bond and each of the C H bonds is a sigma bond Pi 11 bond A pi bond has electron density above and below a plane that contains the bond axis and results from sidetoside overlap of parallel p or d orbitals w unhybridized p orbitals One component of the double bond is a pi bond The formation of the pi bond locls the molecule into this planar structure Molecules are free to rotate about single bonds but rotation about a double bond cannot occur We can compare the r 39 10 I I u and 1a J 39 d H 7 H S CC C C c1 free c1 rotation Cl H Free rotation about the single bond occurs in l2dichloroethane shown above There is only one compound with this name On the other hand 12dichloroethene has a double bond about which free rotation cannot occur H Cl H H C C Cl H C1 C1 trans 12 dichloroethene cis l 2dichloroethene Two compounds with the name 12dichloroethene exist and they have diiTerent chemical and physical properties They are said to be isomers of each other The isomer in which the Cl atoms are on the same side of the C C bond is the cis isomer and the one in which they are on opposite sides is the trans isomer These isomers cannot be interconverted without breaking the pi bond Next we will consider acetylene C2H2 H C C H Each C atom is sp hybridized and each C H bond is a sigma bond The triple bond between the carbon atoms consists of one sigma bond and two pi bonds H370le 05 Molecular Orbital MO Theory Molecular orbital theory is a more sophisticated approach to chemical bonding than Lewis structures and valence bond theory MO theory may explain some properties of molecules that the other approaches cannot For instance experiment demonstrates that 02 is a paramagnetic molecule with two unpaired electrons The Lewis structure that is commonly drawn for 02 does not demonstrate this fact Using molecular orbital theory it is very I 39 39 I easy to rationalize the paramagnetic behavior of 02 In atoms atomic orbitals AO represent regions of space where there is a high probability of nding an electron Atomic orbitals have discrete energies as well In molecules molecular orbitals MO are regions of space where there is a high probability of nding an electron Unlike atomic orbitals however molecular orbitals are o en delocalized over two or more atoms 1 Each MD has a de nite discrete energy 2 A maximum of two electrons may occupy any MO and then only if they have opposite spins 3 MOs are lled with electrons lowest energy to highest energy in accord with Hund s rule What do these MOs look like We can make an approximation that MOs are formed from a linear combination of AOs Orbitals are conserved so for every AO which combines one MO results Linear Combination of Atomic Orbitals LCAO Approach 1 MOS are formed by combination of parent AOs on two or more atoms 2 Only valence orbitals and electrons are considered 3 Orbitals are conserved in chemical bonding 4 MOs follow the same rules as AOs Aufbau Principle Pauli Exclusion Principle Hund s Rules etc 5 Only AOs with identical symmetry properties can interact 6 Orbital mixing is most effective when there is a good energy match ElTectiveness of overlap decreases as energy mismatch increases Overlap of Atomic Orbitals AOs are said to overlap when they occupy the same region of space Overlap may be positive negative or zero Positive overlap signs of wavefunctions of overlapping orbitals are the same Negative overlap signs of wave lnctions of overlapping orbitals are diiTerent Zero overlap equal amounts of positive and negative overlap positive overlap negative overlap zero overlap 324 Consider HZ Each H atom contn39butes a Is atomic orbital As a result 2 MOs result from two di amt linear combinations of the atomic rbitals ls ls giveIisetoZMOs Addition of orbitals ouiios up eieotron de Slly in overlap region Subtraction of orbitals results in low eieo ron natty in tne overlap region t L i 3 l at m Antibonding orbital The bonding MO designated a s is the lower in mergy of the two molecular orbitals while the antibonding MO designated 0 is higher in magy Molecular Orbital Energy Diagram We m 39 L corui Llldllun Jul we do ful an atom For H2 the electron cm gjration is written as 052 since thae are two electrons in the 05 MO Bond Order The bond order in a molecule may be determined from the MO diagram and it corresponds to 12 the di erence between the number of electrons in bonding MO s and the number in antibonding MOs Bond order 12 e39 in bonding MOs e39 in antibonding MOs A higher bond order indicates a shorter stronger bond If a bond order is calculated to be 0 or negative the species is not a stable molecule Integral and halfintegral bond orders are possible The bond order for H2 is 12 2 0 1 For He2 we can use the same diagram as for H2 The additional two electrons go into the 015 MO and the electron con guration is 015 2 0152 The bond order is given by 12 2 2 0 He2 is therefore not a stable molecule In order for two atomic orbitals to combine to form molecular orbitals two conditions must be met 1 There must be a reasonable energy match between the two atomic orbitals 2 The symmetry of the orbitals must be the same In the case of Liz the energy match between the ls atomic orbital and the 2s atomic orbital is not very good so the ls and 2s orbitals do not interact gtxlt 0 25 I z39 ll 25 l 3 25 I l I Li Li2 Li 326 The electron con guration of LiZ may be written as Ols 2 0152 025 2 We can abbreviate the rst part Ols 2 0152 as K to indicate that the K shell of each atom is lled The con guration may therefore be written as K 025 A homonuclear diatomic molecule is a diatomic molecule which is composed of two like atoms for example H2 02 F2 etc A heteronuclear diatomic molecule is a diatomic molecule which is composed of two unlike atoms for example HF CO etc For the rest of the second period diatomics we must consider the MOs that can be formed by the overlap of the p atomic orbitals px py and p2 Two types of overlap are possible By convention the internuclear axis is taken as the Zaxis By this convention the pZ orbitals point at each other and can overlap in a headtohead or sigma fashion Additive overlap results in a sigma bonding MO 02p while subtractive overlap results in a sigma antibonding MO 0 2p These are shown in the gure below The px and py atomic orbitals line up in a sidetoside or pi fashion Additive overlap results in a pair of pi bonding MOs 1121 while subtractive overlap results in a pair of pi antibonding MOs 112p The two bonding MOs are at the same energy and the two antibonding MOs are at the same energy Orbitals of the same energy are said to be degenerate The ip and 112p MOs are shown in the gure below Relative Energies of Molecular Orbitals The energies of molecular orbitals do not vary a great deal from the energies of the atomic orbitals from which they are formed There is some spreading of the bonding MO and anti bonding MO below and above the parent atomic orbitals though In general MOs formed from2p atomic orbitals will be higher in energy than those formed from 2s atomic orbitals A molecular orbital diagram appropriate to the second row homonuclear diatomics up to N2 is shown below The MOs arising from the ls atomic orbitals are omitted We can use this diagram to predict the bond order stability and magnetic properties of the homonuclear diatomics through N2 Molecule Stable Bond Order Paramagnetcdiamagnetic Liz Yes 1 Diamagnetic Bez No B2 Yes 1 Paramagnetic C2 Yes 2 Diamagnetic N2 Yes 3 Diamagnetic The diagram above does not rigorously apply to the rest of the homonuclear diatomics of the second period The reason concems the relative energies of the 2s and 2p orbitals in atoms of the 2quot 1 period elements 2p level demon ln an s orbltal ls more strongly m uenced by llne lnereaslng nuelear charge However amne begmmng of llne penool llne levels diffenn energy by only about 0 2M nnolquot In these Circumstances the wave funenons for the 2s and 2p orbltals beeonne mlxed One result of llne than llne 17 orbltal Ins ordenng of orbltals applles o dmltxogen and llne precedlng elements ln Penod 2 llne orn erossover occumng between dmltxogEn and dloxygm Li2 Be2 52 c2 N2 02 F2 The appropriate MO diagram for the remaining homonuclear diatomisc of period 2 is 2p 2p f j 25 25 39Z 25 0 25 Molecule Stable Bond Order Paramagnetcdiamagnetic 02 Yes 2 Paramagnetic F2 Yes 1 Diamagnetic Nez No We can also use the MO energy diagram for cations and anions of the homonuclear diatomics For example each of the following ions is known 02 0239 and 02239 What are the bond orders and how many unpaired electrons does each contain Moleculeion bond order unpaired e39 of 25 l 02 2 2 0239 l 5 l 02239 l 0 Knowing the bond orders allows us to predict the relative bond lengths and strengths Bond length 02 lt 02 lt 0239 lt 02239 Bond strength 02 gt 02 gt 0239 gt 02239 Heteronuclear Dxatomms appreerably The emerges of the corresponding atomic orbitals m N and o are not exnernely far ap dLhe M dAagmm ofNO somewhat resembles that of 01mm the 0 atomic orbnals e corresponding N atomic mbxmls Usrng Arum Molecular News mm mhlmk the M o diagram for NO does not 0 rnnoduee smous err ener Even Justtwo elements awa the q diagmm atall In the case ofHZF the H15 orbnal xs slgm candy hghenn energy than even theF 2p was As a rs PM L 2p pr zpy Note um the 0 bondmg M 0 rs rnuen closer m energy to 1131 2p orbitals than to the H 15 but H H F whde the 0 Mo rs closer m Energy to the H 15 orbnal y otherwords the HI bondxs polar wrun the F atom nanmg une pamal neganve charge Molecular Orbitals and Delomlized Bonding Another advantage of MO Theory over the valence bond approach lies in the Way that it deals with 39 39 39 39 39 vaoltrecall 39 l39 39 39 39 i Ll 39 m H 39 MO 39Iheory describes this bonding With a single electron con guration In the ralmce bond approach 31s ue uiueu u VAOO39 H 39OOAv The implication is that each 0 atom has three localized electron pairs associated With it and sp2 delomlized electrons may be thought of as one of the pair in the double bond and one of the e C m r quot quot m uimLonebonding quot quot aiiu unciiun bonding as shown below l Fa A o l l l l Antibonding 1c orbital Nonbondlng 1c orbital l l Bondlng 1c orbital Two of the electrons occupy the bonding MO Which is the lowest of the three in energy and two electrons occupy the nonbonding MO Whld39l is intermediate in energy The antibonding orbital Whld39l is highea in energy is unoccupied Intermolecular Forces IMF Intermolecular forces are relatively weak forces of interaction between molecules They are typically much weaker than ionic or covalent bonds In the gas phase we normally neglect the intermolecular forces however they are quite important in the condensed phases and may be used to explain some of the physical properties of liquids and solids Intermolecular forces are responsible for keeping the molecules in the liquid phase The strength of the intermolecular forces affects the ability of the molecules to escape into the vapor phase The strength of the intermolecular forces therefore has a profound effect on the boiling point heat of vaporization and vapor pressure of a liquid Stronger llVJFs result in higher boiling points greater AH3p lower vapor pressure at a giver temperature Van der Waals force is a general term for two types of Ms dipoledipole forces and London forces These forces are described below 1 Dipoledipole Forces Dipoledipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule Polar molecules ten to align themselves in a to 7 fashion The molecules of HCl in the solid phase are wellordered in this alignment In the liquid phase these interactions are only partially disrupted In general the more polar the molecules are the stronger the dipoledipole forces will be 333 2 London Dispersion Forces Consider a normally nonpolar substance such as neon or argon 12 7 10 r i On average the electrons are distributed in a symmetric fashion around the nucleus At some instant in time this distribution may not be perfectly symmetric and an instantaneous dipole moment results This temporary dipole moment can affect the distribution of electrons in a neighboring atom or molecules as shown Thus an instantaneous dipole is induced in a neighboring atom London or dispersion forces are the weak attractive forces which result from these small instantaneous dipole moments which arise due to uctuations in electron distributions London forces increase with increasing molecular weight Gas BPK AW Hydrogen Bonding Ne Ar Kr Xe 27 87 120 166 2018 3995 8380 1313 A third type of IMF commonly encountered is the hydrogen bond Hydrogen bonds vary widely in strength They are responsible for some of the properties of liquid and solid water as well as the secondary structure of many proteins Hydrogen Bond A weak to moderately strong attractive l 0 force between a hydrogen atom covalently yquot r bonded to a highly electronegative atom thiHydrogen bonds H such as F O or N and a lone pair of g l electrons on another highly electronegative piss 4 H O atom H TKOH To appreciate the effect of hydrogen bonding consider the A B compounds uoromethane CH3F and methanol CH3OH Both of these have about the same molecular weight and CH3 CHBOH dipole moment We might expect the magnitude of the lMFs I in each of these compounds to be roughly the same and molecular We gll mm 3403 3204 therefore might expect their boiling points to be comparable dipole momem D 181 170 However as you can see there is a large difference in their boiling points so there must be some type of IMF present in boiling point C 78 65 methanol which is not present in uoromethane 334 Let s consider valiations in boiling points of some hydrogen containing compounds of groups IV to VII There is a trend observed where as the atomic weight of the nonhydrogen atom decreases the boiling point decreases This trend does not hold for the compounds containing F O and N These are the atoms which allow for hydrogen bonding Hydrogen bonds are substantially stronger than van der Waals forces but in most cases considerably weaker than covalent or ionic bonds aamng pmth a C aamng pmth a C HTE 2 mm EIZEI mm EMU DMD EIBEI man 2m 4m an an mu 12D 2m 4m an an mu 12D Mulecularwelght Mulecularwelght A E Introduction to Covalent Bonding Section 31 Molecular Orbital Theory a more sophisticated approach to chemical bonding than Lewis structures and Valence Bond theory Why p 9 P enigma m Mg Introduction to Molecular Orbital Theory Section 32 otoln atoms atomic orbitals AO represent regions of space where there is a high probability of finding an electron Atomic orbitals have discrete energies as well otoln molecules molecular orbitals MO are regions of space where there is a high probability of finding an electron Unlike atomic orbitals however molecular orbitals are often delocalized over two or more atoms 1 For orbitals to overlap the signs on the overlapping lobes must be the same 2 Each MO has a definite discrete energy Antibonding MOs are higher in energy while bonding MOs are lower in energy 3 For significant mixing to occur AOs must be of similar energy 4A maximum of two electrons may occupy any MO Hunds Rule 5 MOs are filled with electrons lowest energy to highest energy What do these Molecular Orbitals look like Bonding and Antibonding Molecular Orbitals MOS Bonding Molecular Orbital Antibonding Molecular Orbital Addilion of orbilals builds up electron denslly in overlap region 0 0 amp ls I Bonding orbital Subtraction of orbitals rcsulls in 10 w elachon density in the overlap region opc 0ogto ls Antibondlng orbital Molecular Orbital Energy Diagrams and Bond Order H atom H2 molecule H atom CD 15 13 Energy 013 Bond Order BO B0 1 of e 39 in bonding M05 of e39 in antibonding M05 What is the bond order for H2 Molecular Orbital Energy Diagram Add More Electrons Li atom Li2 molecule Li atom Electron Configuration for a Li atom 395 How man electrons are in Li D V 2 2 lt 2x a Energy What is the molecular orbital configuration for Liz What is the bond order for Liz Molecular Orbital Energy Diagram Nitrogen N2 Electron Configuration for a N atom Number of electrons found in a N2 molecule A l 6 What is the molecular orbital configuration for N2 What is the bond order for N2 Molecular Orbital Energy Diagram Dioxygen 02 Earlier you saw that 02 was attracted to the poles of a magnet This is because 02 is paramagnetic it contains unpaired electrons The Lewis structure and Valence Bond Theory both incorrectly predict that 02 is diamagnetic it contains no unpaired electrons What does Molecular Orbital Theory predict 1 Electron Configuration for an O atom 2 Number of electrons found in an 02 molecule Molecular Orbital Energy Diagram Dioxygen 02 A 039 2p 13 P H r X r 2 I I r r 4 p bu 7 x 3 What is the molecular orbital configuration for 02 What is the bond order for 02 Is 02 paramagnetic or diamagnetic Energy Summary of Period 2 Homonuclear Diatomic Molecules A 0219 7r2p a2p 0217 zp mp H39 L4 H H 44H nap 023 H C21 025 623 725 B2 c2 N2 02 F2 Fundamental Properties of Group 13 Elements Boron Aluminum Gallium Indium 39 lallium Symbol B Al Ga In T Atomic number 5 13 31 49 81 Natural isotopes I0111978 27A1100 69Ea604 n428 203112950 A abundances IS8022 Ga396 Sin9572 205117050 Total no of isotopes 6 7 14 19 21 Atomic weight 1081 2698 6972 11482 20437 Valence electrons 23221 352301 4324171 532511 6526p1 mp bp C 2300 2550 660 2467 29782403 1566 2080 3035 1457 Density gcm3 234 270 590 730 1185 Atomic radigxs 143 153 167 171 metallic A Ionic radius 025X4 0534 0766 0946 3 02 6 Shannon Prewitt A cm 39 11646 Pauling EN 20 J 15 16 17 18 Charge density Chargeionic 120 57 40 32 329 madius unit charge x 1061 E0 a V 090 m 166 056 034 033 Oxidation states 3 3 1 3 1 3 1 3 covalent Ionization energy kJmol 801 578 579 558 589 Electron Af nity kJmol 23 44 35 34 48 Discovered by date GayLussac Wijhler de Boisbaudran Reich Crookes 1808 1827 1875 1863 1861 pr o2 8203 A1203 33203 111203 T120 Acid base eharacter of oxide Acid Ampho Ampho Ampho Base rpw N2 BN AlN GaN lnN None rpw halogens BX 3 Alzx 3212K6 In ZX TlX rpw hydrogen Crystal structure Hexagonal fcc Cubic Tetragon a1 Hexagonal Group Trends Sections 131 ozoCommon Features otoMelting and Boiling Points ozoCovalent Character Table 131 Melting and boiling points of Melting Boiling Element point C point C B 2180 3650 A1 660 2467 Ga 30 2403 In 157 2080 T1 303 1457 Table 132 Charge densities of Period 3 Charge density 3 Group Ion Cmm 1 Na 24 2 Mg2 120 13 Al3 364 Boron Section 132 Figure 131 lcosahedral arrangement of boron HBO 0amp0 Figure 132 Actual structure of Figure 133 Structure of the the borate ion in borax DeFOXObOFate Ion Boranes Section 134 otoPreparation and Properties otoApplications Figure 134 Structure of tetraborane B4H10 The boron atoms are shaded Figure 135 Structure of the Bsz anion Figure 136 Structure of the Blngf anion Sodium Tetrahyridoborate Boron Trifluoride Boron Trichloride Sections 135 137 t Sodium Tetrahydridoborate Filled np orbital of halide t Boron Halides BF3 amp BCI3 Unhybridized 2 p orbital Aluminum Section 138 ozoPreparation and Properties ooooo meeeee Figure 1310 Formation of a single oxide layer on the surface of aluminum metal The small A aluminum 3 ions are indicated Appllcatlons by the solid circles Aluminum Cont t Solubility of the Aluminum Ion Solubility mmol 39 Ifl 10 DrCf OCO A1OH2613 and other cation species WHO S 00 AKOHLKH 20 quot aiuminate A1HZQ63 Aluminum Halides Section 139 otoBonding Coordination Numbers of Metal Atoms in Group IIIB13 Halides F Cl Br I Al 6 6 4 4 Ga 6 4 4 4 In 6 6 6 4 T1 6 6 4 Gallium Indium and Thallium ozoPreparation and Properties ozoApplications Biological Aspects Section 1312 otoBoron otoAluminum 3 Thallium Boron Reaction Flowchart Section 1313 NaBH4 NaH BF NaH BZH6 OZ 3203 C B4C TiCOZ TiBZ NH3 H20 Mg F3BNH3r 1131303 B Aluminum Reaction Flowchart A1113 HF AlzBr6 A1203 Na3A1F6 Brz V w C12 OH A1 A1OH4 A1c13 H H OH H A1OHZ63 A1OH3 OH Quantitative Aspects of HalfReactions Section 87 Voltaic Cells Galvanic Cells electrochemical cell in which a spontaneous reaction generates an electric current Oxidation HalfReaction Cds 9 Cd2 2 e39 Anode Ag cathode W Ag 1 equot Ags 39 Cathode 2Sum of the two halfreactions Cd gtCd22c Age Ag Cds 2Ag a Ce2 2Ags 2Cell Notation Standard Potentials t Standard Electrode Reduction Potential E Ag 1 equot ltgt Ags E 0799v 2 H 2 equot gt H2g E ooov Col2 2 equot 2 Cds E o402v ozoPredicting Spontaneity for Redox Reactions E ceu Reduction Cathode HalfReaction Oxidation Anode HalfReaction Table of Standard Potentials um mu sund summ mam Pnlnnllal m Annam 5mm a 25 sundvd mud ledur un mum1 milRum E39N 2mm z mun 20mm 7013 1 nm c now on m m 0m L El tuq 201mg nun nang v w e mm mm Nom A 21mm 093 M i an mm an 4 mpg anewl39m NH we v at 2 aay 711200 L33 mm z cram r as Mno aay mm s Mum mom ms 39 mom quot2 u Fzm a 2 2mm 2 a1 Oxidizing and Reducing Strength Standard Cathode Reduction Potential HalfReaction EM L1aq e L10 304 Naaq e Nas 271 Mg aq 2e 2 Mgs 238 A13aq 3e A1s 466 Su2uq 26 Sum 014 Pb2uq 2e Pm 7013 Fe3aq 36 Fes 004 21mm 25 2 Hzg 000 Sn4aq 2e 2 Sn aq 015 Cu2aq e Cuaq 016 Cu2aq 2e Cus 034 Mnomaq 8Haq 5e Mn aq 4HZOI 149 H202aq 2mm 2cquot 211200 173 s2ogz aq 2 2so42 aq 201 F2g 257 2 2F aq Effect of Concentration on Ece The Nernst Equation RT Ecell E2611 1n prOdUCtS nF reactants where R 8314 V G mol391 K 1 T temperature in Kelvin n moles of transferred electrons F Faraday constant 965 x 104 0 mo 1 ozoCalculate Ece for the reduction of MnO439 to Mn2 when the pH of the solution is decreased to 400 and MnO439 10 M and Mn2 10 M Electrode Potentials as Thermodynamic Functions Section 88 AGO nFEoce F Faraday constant 965 x 104 J V391 mol 1 Successes of Crystal Field Theory and Electronic Spectra Sections 198 199 00 Magnetic Properties dxl yl dzz dxzyz dxy dxz dyz dxy dxy dxz dyz dzz dxziyz d 2 Octahedral field Tetrahedral field Z dyz dxz Square planar field ozoColors of Transition Metal Complexes Electronic Spectra 1 3112 d9 2 Ti d1 1 1 1 1 dzz d Mz Energy A 1L 1L 1L 1L 1 quot dxy dXZ dyz Ground EXClted Octahedtal field state state
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