INORGANIC CHEMISTRY CHEM 462
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Ebz L Carbon and Group 14 Elements Chem 462 NOV 17 0 All bonds staggered ideal o dC C 154 A like ethane Trends in atomization energies a G a I Cohesive Energies klmm E 8 a e u 1 2 3 4 5 6 7 Number ofs 8 p valence chattrons Explanation 1 ij mqu ur Ulmulaluun m hm c alumna h u 4 um mm u mum w n mnlwulmeuLK m 39 awnml p rmm mrkmiuilmllcmclpl mm mylllh up arm 0 439 H m m f 1 m 26 quotu 1 cO 47819 0 eg BeZC anti uorite 5mm Ach 7 more Complex but still no Crc contacts less than 316 A BeZC or Al lC3 H20 gt mostly CH41 BeOH or AlOH3 largely ionic v v Curbides b acetylides containing czcz o CaC2 s most important agam largely tome 3210 3c gtCaC2 co gt zzoo39c AH 4657 kImoH x dustnal Cathq mg zcsz CacZC2I2gt Hztxl reo39c gtCaC2 csz azs39c lab Cac2 ls mdusmally important 9 CaC2 N2 gtCaCN2 c gt 1000397120039c N CO2HZO ACaCOS H NCN Cyanamlder can be tnmenzed to melamme for melamine ere me carbonrmetal electronega vity J i Carbon as Industrial Reductant Carbon as cokeis often the cheap reducing agent of choice in metallurgy eg SnOZs C gt Sn 2COg ZnOs C gt Zn COg 2C coke 02 gt 2COg 1200 2000 C Fe203s 3C0 gt 2Fel 3COZg 1000 C w V Simple Key Reactions Generation of hydrogen from natural gas or carbon monoxide from coke eg CH4g 2H20g gt COzg 4Hzg CH4g H20g COg 3Hzg COHZ mixture is syngas C0g HzOg C02g Hzg watergas shift reaction M in CCl4 industrial preparation CH4g ZSl gtC52g H25g N 700 C 32 3 Clzg gt CC14g SZCIZU 32 2 SZCIZU gt CC14g 6 55 o CCl4 is generally inert making an excellent nonpolar solvent unfortunately carcinogenic 0 potent greenhouse gas and catalyst for ozone decomposition 1 Freons mixed Chloro uorcurbons lt3 9 CCLA HF gt CFC3 CFZCl2 HCl These are very stable nonpolar volatile 7 excellent refrigerants 9 Unfortunately the scienti c case for their role in ozone layers degradation is strong 7 the work was the basis for the 1995 Nobel prize in chemistry 1 8 zone degradation oversimpli ed CF3C1 hv gtCF3 Cl stability of freons allow them to survive transport to upper stratosphere 03 C1 02 C10 C10 gtO C1 0 C10 02 C1 CCLA HF controlled gtHCFC12 HCl HCFClZ heat 70039CPtcatalyst gt FZCCF2 FZCCF2 free radical initiator gt F Teflon F l Freons and Te on owe stability to strong CiF bond 486 kImol NZ in Carbonates 03 is the major form of carbonate in nature limestone CaCO3 HBOquot gt Ca2 HCOB HBOquot gt H20 CO2 In water llcarbonic acid exists mostly as dissolved C02 HZCO3 CO2 H20 K COZH2CO3 600 This equilibrium is responsible for lower apparent acidity of carbonic acid N V Hydrogen Cyanide O HCN weak acid pKa 921 2 CH4 302 2NH3 a2 HCN 61420 800 C Wcatalyst 2 HCN N m C N A N l mm H2 Nylon 66 H2NNH2 adipommle xx I mgwsAW la HO 0 C H 3 adipic ma f Silicon Germanium Gray Tin O quotquotquotquot quot3513 Diamond structure 7777777777 7 Si and Ge semiconductors Si is the purest element manufactured 0 Industrial grade prep SiOZ excess C Si 2 CO excess used to prevent formation of SiC SiC Si02 3Si 2C0 0 High purity Semiconductor industry Scrap Sl smgle crystal 5 Zone C12 Iefming v Mu or Zn SiCL may Snponge2MHClz b ll 57b disnllarlon Silicon vs Carbon 0 CX TL bonding important 0 3 and 2coordination is common in multiple bonding to X C N O and S 0 Si X TL bonding is weak 0 4 coordination is overwhelmingly preferred eg compare 0 0 OC0 vs Silicon vs Carbon 0 for carbon bond strengths to H amp O are comparable to CC bonds 416 amp 336 kImol vs 356 kImol for single bonds 0 occurence of coal hydrocarbons is reasonable 0 for silicon bonds to oxygen are much stronger than SiSi bonds 368 vs 230 kImol 0 no natural occurence of Si Si bonds canbe synthesized however Silicon vs Carbon 0 Reactivity at tetrahedral Si is greater eg SiCl4 HZO gt SiOZ 4 HCl AH 78 kcalmol CCl4 HZO gt COZ 4 HCl AH 84 kcalmol but reaction is rapid for SiCl4 and never occurs under normal conditions for CCl4 0 5coordinate transition state much more accessible for Si dorbital participation and larger size thought to be important factors in Si case Chapter 8 Hydrogen CHEM 462 Wednesday Nov 3 T Hughbanks What comes now 0 From this point on we will be discussing more descriptive chemistryquot 7 this requires integration of the principles coveredup to now 0 I ll try to explain how to use the principles to help organize the descriptive facts but this takes practicing chemists a long time to do well 0 There is no way we can do more than scratch the smface of selected topics sol will have to pick and choosequot whatl talk about You will see more divergence from the text in several places Remember general trends and Periodic Table 0 Electronegativities and related ionization energies and electron af nities 0 Radii cation and anion 0 Polarizabilities 0 Bonding summary 7 Octet rule for sp7block elements 7 Hydridization mbonding 7 Covalency ionicity Hydrogen is Unique 0 Fairly large ionization energy small electron af nity never is present in compounds as a naked cation forms an anion only With highly electropositive cations in the solid state 0 No core electrons highly polarizing makes fairly strong bonds With many elements 0 Most bonds have substantial covalency even if polar Hydrogen doesn t belong O to the group I metals Although hydrogen makes a monocation in aqueous solution Ht isn t at all like a group I cation Reactivity of elemental hydrogen is not at all like alkali metals either 0 to the halogens Although the element exists as a diatomic molecule like the halogens it is much less reactive rarely reacts to give Hi and the H7H bond is quite strong 436 kl mol1 Covalent hydrogen compounds have structural similarity to halogen compounds but very different reactivity Hydrides 0 Traditional categories 7 ionic 7 covalent 7 metallic 0 Divisions between categories are not entirely sharp Hydride periodic geography Examples to be discussed 18 I3 M l5 16 I7 sum I Molcuxlur 1cmllr nmhmmimi nr unknown Interstitial hydrogen in metallic hydrides Basic physical and chemical properties 0 H2 is lowboiling 26 K low melting 20 K colorless gas Strong bond 436 kJmolrl 0 Basic reactions to make covalent hydrides 2H2 g 02 g gt 2 H20g AH 488 kJ H2g F2 g gt 2 HFg AH 542 kJ 3H2g N2g gt ZNH3 g AH92kJ Bonds broken are strong but strong bonds made are too and are more numerous Slow rxns at rt Covalent hydrides have low bps boiling points CH4 7164 C colorless gas SiH4 7111 C colorless gas PH3 788 C colorless gas BZH6 790 C colorless gas some Hbond vaporization Strengths enthalpies F H F 30 165 kJmol F HFH 2 30 kJmol HO H OHZ m 25 kJmol HZN HquotNH3 17 kJmol 7 z 3 4 5 Period m rnmlvlmng lemeur F Hydrogen bonds and water 0 H bonds give water some special properties 7 high melting amp boiling points essential for our existence 7 structure of ice is very open and very stable 7 density of water increases slightly on melting so ice oats 0 H bonds also important in structure of bio molecules like proteins DNA Hydrogen bonding in Ice JL quotv taint q I 3 m Ice has a very open diamondlike structurein liquid H2O collapse occurs when a few Hbonds break Liquid H2O is therefore denser than ice A Personal Favorite Zr6B C1185 ions trapped in ice cages Two Forms of Hydrogen H2 Ortho and Para Two special circumstances peculiar to hydrogen conspire to give quantum mechanical wavefunction symmetry a dramatic effect on the macroscopic thermal properties eg heat capacity of H2 Rotational levels in H2 are relatively Widely spaced and because H2 is a homonuclear diatomic they alternate in symmetry as a function of J the rotational quantum number J even symmetric J odd antisymmetric 6 The nuclear spin states have opposite symmetry too I 1 symmetric I O antisymmetric O 11 Nuclear Spins m1 A 1 O The nuclear spins couple into triplet I 1 and singlet I 0 total 0 nuclear spin states 0 The triplet state is 1 symmetric ortho the singleti 39 0 v5va Rotational Levels 7 m 97 O The ortho and para forms correspond to the odd 9 and even rotational levels 0 As long as the interconversion between or levels is slow the two mo l 7 forms have distinct thermal properties 6 s Some Statistical Thermo Relations JJlh2 E 21 l g 2l 75 CT qm zgje 1 partition fct sum over states I 2 qpm ZgJEJ 2lt2J1gte h W odd odd Mo 3 ZgJEJ 3 2lt2J1gte hzW even I even U RTz 91on C aUm rot rot m Heat Capamues Hydmgen as a reducmg agent 3 mdhydm zxemwmuuamvg 25 m m wmmm wwwmm m velysvmvgmdmms me mdmzny es cu0s u g cm moer 14 amsyzigog mm w HA1045 M a 1445 o was o lms szH Namx reducmg agen39s was mews a Hangs 1 mew t mm mm my mum hhMHAswEggazmH mum uzkog mgm mums 2m 55 mm seismme mm m kmxsdmms uni mwmu seismve Class 92 Redox Chemistry 111 Wednesday October 20 CHEM 462 T Hughbanks Z U Latimer diagrams Nitrogen E n acid solution pH 0 171 117 1111 159 177 117 11 127s H39NO N01 b H39NOz D No D N20 gt N2 gtNH30H b N2H539 DNHA39 m 1111 basic solution pH 14 4116 41quot M 0 76 09 30 0 73 01 NOf p N101 p NOT DNo DNzo gt Nz gt NHzOH gt NzIh gtNH2 D 1 ifFrost diagrams Nitrogen J Frost dizgmns aqueous Nitrogen species nuo Lownwbmmq oxidation number ome important NO reactions 9 NH4aq NOaq gt NZOg HZO1 spontaneous NZO Useful anesthetic laughing gas Nearly thermoneutral rxns 9 NOg NOZg HZO1 gt ZHNOZaq rapid 9 2 HNOZaq gt HNOgaq 2 NOg HZO1 rapid 9 2 N02 H200 gt 2 HNO3aq HNOZaq slow These or similar rxns important for nitric acid synthesis M u 3 fExample Problem 0 Write balanced halfreactions for the reduction of NO to N20 and N20 to N2 in basic solution 9 Is N20 thermodynamically stable With respect to disproportionation to NO and N2 in basic solution What is AG 7 3K ll fiatimer diagrams Manganese acidsolution pH 0 09 12 29 09 15 1quot MnOf pHMnOf p HgMHOA gt M1101 p Mn gtMn p Mn basic solution pH 14 056 027 093 mm 4223 456 MnOr39 p Mnot39 MnOf39pMn0 an03 Mn0Hz gtMn HzozHzo 02H20 HVHZ aIanoi oxidation number n E Z 39 fManganese VS Rhenium wNHop wamm 3 4 oxidation number Manan vs Rhenium acidic soln Mnor Cyclic Voltammetry 0 Most COIHIHOII tool of inorganic chemists to survey re ox properties of new moleculesquot 0 Voltage at the working electrode is swept against time in a sawtooth function CU rre nt I time Voltammo gram of a Cluster IMOGSCXCN6I in 025M KOH I I I I I I I I I I 716 714 712 710 708 706 704 702 0 02 V VS SHE Selections from Chapters 9 amp 16 The transition metals IV CHEM 462 Monday November 22 T Hughbanks Jahn Teller Distortions 0 JahnTeller Theorem Nonlinear Molecules in orbitally degenerate states are inherently unstable With respect to distortion 0 Explanation If a molecule has suf ciently high symmetry that it is possible to have two or more molecular orbitals With the same energy by symmetry If a degenerate set of orbitals is occupied unequally With electrons the molecule Will distor t TMs Some Oxidation State Trends It Maximum can never exceed the group number It Highest Ox State Ti V VV CrV MnV 7 after Mn for 1st row ox states higher than 3 hard to achieve It Early metals tend to be Sour ynr LnrnY TirvY Z I Oxides are more acidic in higher oxidation state most stable in highest oxidation state rvy Hfrv I Some broad chem similarities are found for ox states 11 amp III 6 or 4 coordinate for complexes in solution or in crystals aqueous chem etc It Except for Cu macid quot quot quot 39 quot for ox states lower than 2 Acidic High Oxidation State Oxides 0 CrVI is acidic Cro3 Hp gt HzCrO4 gt HCIOJ H3O 0 Green Crm is amphoteric CrZO3 H3Oaq gt CrHZO63 acid base properties similar to AlHZO63 ocCrzO3 OH aq gt soluble chromites 0 HMnO4 H2CrO4 H3VO4 acidity trend similar to HClO4 H2504 H3PO4 Oxidizing Agents acidic soln MnO4 8 H 5 e gt Mn2 2 H20 E 151 V CrzO72 14 H 6 e gt 2C15 7 H20 E 133 V V02 4H e gt 2V3 2 H20 E 0668 V Common Oxidation States Typical compounds H state Compounds of all elements TiCu are found 0 in water all MHZO62 are known except Sc 0 in aerated acid solution V Cr and Fe are oxidized to 3 0 solid state oxides and halides known for all M11 metals with varying degrees of ionicity 7 MEX2 all have CdXzitype structures 7 MHO compounds surprisingly complex Common Oxidation States Typical compounds III state All T M elements I for Ni amp Cu ox state III is fairly rare FeCl3 some similarity with AlCl3 Both form Mpg molecules in liquid gas phases Both can act as Lewis acids eg as Friedel crafts catalysts but AlCl3 not much ofan oxidant Tia Co form MmHZO63t ions MnmgtIVInH amp Comgt Co a quite madin Fe HLO5339 only forms in strong noncomplexing acids Common Oxidation States Typical compounds 111 state is less stable for later TM s eg ComHZO63 e a COIKHZO62 184 V but CoDI canbe stabilized by complexation by strong eld ligands eg ComNH363 e a C0HNH362 01 V synthesis CoClzsoln 4 NH4 20 NH3 02 a 4 ComNH36C13 2 H20 4enH 8enO a 4 Comen3C13 2 H20 Common Oxidation States Typical compounds IV state important for Ti V TiOz TiClA extensive chemistry of VOZquot but for later rstrrow TM39s this tate is found mainly in uorides oxides oxo complexes actual c harge Higher ox states only oxides uorides oxy uorides For a given ligand type higher ox s tes are characterized by increasing covalence eg including extensive M70 7 bonding oxidation state First ROW TMs vs 2nd and 3rd TM Series I Radii 2nd and 3rd row N 01 a 02 A larger I But 2nd and 3rd row radii close to each other lanthanide contraction I Some 2nd and 3rd row elements have very similar chemistries eg Zr and Hi Overall 2nd and 3rd row elements are more similar to each other than to lst row I Metalrmetal bonds stronger and more numerous examples are known for 2nd and 3rd row elements This leads to substantial differences in lower oxidation state chemistries I Higher oxidation states generally more stable for 2nd amp 3rd row Examples of oxidation state differences CrOi 7 very powerful oxidant oxidizes most organics forms more in acidic so1uu39on M003 we r mi1der oxidants will oxidize H2 amp NHi but not most organics RuON OsO r are strong oxidants but no Fe analog exists at all Mno disproportionate to form Mno2 and Mno suong oxidant while ReO3 forms a stable metallic solid T0207 Re207 e isolable so1ids with reasonab1e siabi1iry Mmo e explosively oxidizing o Niw me uorides eg M ZNiFS and oxides BaNiOi Pd fairly uncommon but range of compounds 1arger man for Ni Pd H0 mo 7 aqua regia gt PdCl or PdCl Cl2 gt PdClsz Pt fairly common Pix have an been made MZPLX5 easin preparedX Cl Br I amine complexes as salts Pdam2X22 X2 ermamx oo 2nd amp 3rd row aqueous chem O Genuine aquo ions conmining M and H20 only are relatively rare Exceptions I MoHZO63 can be isolated but even for Mom binuclear ions 2 also occu In ClOA solutions PdHZOAZ amp PtHZOAZ can be made eg from PdO or KZPtCIA I RhHO63 amp IrHO63 reasonably stable but subject to air oxidation 0 Why not more I For decounts less than d6 maybe d5 bridging and terminal oxo complexes are common I Late TMs in lower oxidation states are oxophobic Chapter 4 Ionic and Other Inorganic Solids CHEM 462 Wednesday September 22 T Hughbanks Structures of Solids 9 Many dense solids are described in terms of packing of atoms or ions 9 Although these geometric descriptions are often used and can appears as though solids are assembled from atomic marbles the forces and physical laws that govern solid state structure and molecular structure are exactly the same 9 These are just useful ways to Visualize and classify structures 9 See the class web site for longer animations Also useful In 3 wwunluttaa uk ohnrmnt Unit Cells Not 7 9 Text describes the shaded l areas as unit cells Are either of them really unit W cells l 9 Can you offer a better a 0 choice Fig 21 p 36 Graphite Unit Cell Concepts eb site for this animation Hexagonal Close Packing Unit Cell o All crystals are built up from units that are repeated throughout their structures known as the unit cellquot 0 How many atoms are in the unit cell of an hcp element Cubic Close Packing FaceCentered Cubic o The ABC stacking sequence generates packing of atoms with cubic symmetry 0 When Viewed from the perspective of the cube the cube is seen to have an atom in eye face face centered cubic Eady cgva cm Fm Hales m muse Whngs 4 Rocksalt NaCl type ccp stacked Cl all octahedral holes lled Zinc Blende sphalerite ZnS Type o ZnS zinciblende is described as fee arrays of 52 ions with half the tetrahedral holes lled with an cations o antistructures have own antistructure Zinc Blende heteroatomic diamond Flourite Can Type o CaF2 uorite is described as an foo array of Ca ions with all tetrahedral holes lled with F anions o antistructures have anti uorite structure CdIZ type hcp stacked I half lled CdClzitype cop stacked Cl halfi lled CdIz BiI3 type hcp stacked I 13rd lled l V AC BiI3 type ball amp stick View Interpreting Polyhedral Pictures x o 0 M Mgtlt6 i d x a 132 K i Xiz MoSz type Trigonal prismatic Mo MOSZ type A Lamellar Structure Molybdenite Perovskite ABO3 ABOsg Examples BaTiO3 KNbO3 R603 2 D1303 Perovskite With a Shifted Origin Perovskite ABO3 ABOsg Examples BaTiO3 KNbO3 R603 2 D1303 BaTiOS is a ferroelectn39c with spontaneous wi Lamoe Energwesr Dem vvlt mmbVIrlesxn i2 lonrlon Coulomb repulsions TL Nquot mtemcnons betweanon cores I NE quotA 9 3 i5 INS eg Nacnype cgsumts rzfzctorof mm m 73 39m mun be xnduded to convert to 51 m m Ns VC 1 mb o e1z z 12 eZmxz 2 2 2 x e warm z o e z Nzcnype vc 1 mb NAe1rzzx A A 75 12 7W3 wry15 The Madelung consunt a geomemml parameter that 15 the Ems forzll compounds of a gwa structure type Ion core Ion core repulsions o vrepmsm NA c exppum Allemalive estimale from compressibilities r W N MM is always muchsmallerthan atypical 39 39 34s A Internuc ear Islance 0 m E v NAeZrZ 271A He 5 NC39exprr Ne 7 Min39 39 dVdr 0 to 0 Iain A 9 Ilee 5a C 39exprltrmmrgt r e2 2 2 A mummy me NAeZrmmZ mu 7 gomm vmm N clmm 2 mu 7 1 A Thermochemical Born Haber Cycle M EX Mg gtlt9 gt Mg gtlt g 5M l2Dx2i i39AHL AHf Ms l2gtlt2g gt Mgtlts Values NaCl klmoi SM sublimation Enthalpy of Metal 298 K 108 DXZ dissociation Energy of X2 bond 298 K 242 M Ionization Enthalpy of Metal M 496 EX electron attachment Enthalpy ofgtlt atom 349 AHL Enthalpy for separation of salt to ions theory Thermochemical Born Haber Cycle I M E M xltggt M Ma x7ltggt 5M V2Dgtlt2T l39AHL AHf Ms l2gtlt29 a Mgtlts AHL me NAezrmlnz Zl1 39 Wrmmm for NaCl rmm 2731 A 5162 au 776kJmol Experiment BornHaber cycle AHf SM V2 DX2 M 7 EM 7 AHL for NaCl measured value of AHf is 411lltJmol 411 108 2422 496 349 r AHL H 2 for C correction 395 IJ mol Consequences of Lattice Enthalpies O Electrostatic component stabilizing ionic solids gives us ZA ZE Lattice energy x d a lattIce spacing O TheImal stabilities Large Cations stabilize large Anions Effects of charge and size on lattice energies mps melting points C Can 1423 SrCI2 872 L120 gt1700 N320 2800 2580 2430 1923 Classes 7123 Intro to Coordination Complexes MWF October 25 27 29 CHEM 462 T Hughbanks Coordination Compounds I Even in cases where we think of solvated cations evidence for solvent cation interactions shows that more than electrostatics is involved I Example ionic vs hydrated radii A of alkali metals Li Na K Rb Cs ion 090 116 152 166 181 hyd 340 276 232 228 228 Real Coordination Complexes I For higher oxidation state less electropositive transition and main group metals the ionic picture is not adequate I Metal ions can be viewed as Lewis Acids ligands as Lewis bases Types of Ligands I Monodentate X H20 ROH ROR THF OH OH DMSO NH3 NR3 py PR3 CHS CN CO I Multidentate en dmpe glymeacac bpy phen dien terpy porphyrin I Some which can be mono or di carboxylates nitrate sulfate dithiocarbamamtes Stability Constants Stability constants are defined in aqueous solutions as sixcoordinate case MOH2eiquot L 2 MOH25L1quot H20 K1 MOH25Lquot L2 MOH24L2quot H20 K2 K1 MOH25Lquot MOH26quotL etc b6 ML6quot MOH26quotL6 K1 gtlt gtlt K6 AG AH and AS can be defined for each process and related to each equilibrium constant Chelate Effect 0 ammine ligands replace aquo ligands on NiH in an enthalpy driven reaction b 2 1086 NiHZO62 6 NH3 a NiNH362 6 H20 0 The free energy of a similar reaction with en has an even more negative b5 2 10133 NiHZO62 3 en a Nien32 6 H20 0 Direct comparison K 2 1097 NiNH362 3 en a Nien32 6 NH3 For this AH 712 kJmol Bonding and Electronic Structure 0 Molecular orbital picture for Lewis acid base interaction 0 d orbitals acceptor orbitals on transition metals 0 MO picture for octahedral complexes oLow spin vs high spin complexes ligand eld splitting vs pairing energy Metal Ligand bonding I Concepts of hardness and softness already discussed qualitatively Pearson s 1963 JACS paper introducing these concepts is posted in the web site handouts section I Intro to Ligand tields an M0 scheme for a typical octahedral transition metal complex symmetry adapted LGOs are shown in Figure 424 The Spectrochemical Series CO CN gt N02 gt en NHZCH3 gt NH3 gt NCS gt H20 oxalate2 gt OH gt F gt C1 gt NCS gt Br gt 1 Factors affecting dorbital splittings 0 Lewis basicity 0 Electronegativity of donor atom 0 ndonating or naccepting character of the ligand 0 Metal atom charge amp row 15 lt 2 d lt 3 Properties of complexes 0 Structural and Thermodynamic effects of d orbital splittings 0 Examples Hydration energies and radii of 2 ions in 1st row TMs Data to Explain for 2 ions Hydration Energies kJmol HighrSpin ion radii A 3000 10 2500 Properties of complexes 0 Magnetic properties Highspin and lowspin comp exes 0 Optical properties 7 measuring Am 7 effect of spin on transition probabilities Spin Only Magnetism O c is the magnetic x Cmolar Mm and obeys T the Curie Law for simple paramagnetic compounds 0 MB is the Bohr Magneton MB 57884 1075 eV T39l F sx 851136 1 39 WNW 5585 J T391 moi1 gs 20023E 2 T Tesla 8 8h ZmEc N 2 C mohr 3kg 2 Carbonyl Compounds and the 18 e rule 6 Carbonyls and other strong eld ligands tend to form transitionmetal compounds that conform to a socalled 18 e rule 0 When e count is low MM bonds 0 Other important organic ligands can be included among those that form complexes that tend to conform to the 18 e rule H CH3 and other R 7 all 2 e donors C5H5 Cp 7 a 6 e donor other organics Class 31 Basic Molecular Orbital Theory Diatomics Monday Sept 13 CHEM 462 T Hughbanks Intro to M0 Theory l Usual approach Linear Combination of Atomic Orbitals Atomic Orbituls w K N Molecular Orbitals Wu K M N w rm Him cwzL wt l Coef cients give the contributions that each of the individual AOs make to the M0 Readingt etc I Continuation of Chapter 8 Sections 37 310 in this set of notes I Download the resume of MO theoryquot from the class Web site for use in conjunction with today s class I Next Appendix I Read ah of the Appendix this is material on symmetry MO Theory v General Features Starting with N AOs we can make N molecular orbitals No more and no less M05 are mutually orthogonal and normalized overlap between any two MOS is zero and overlap of a MO with itself l Pauli exclusion principle an M0 can accommodate 2 electrons One is spin up m l2 one is spindown m5 12 Liz Beg BI C2 N1 I v wlkv39171wxquot0z O My q rw W uvzuV IWH my 51 viz m Liz Bel B2 C2 N2 s p mixing lanes any 3 masks mm HF polarity in 21 heteronuclear diatomic Fig 317 energies relative to In PR Covalent Insulators and Semiconductors Chem 462 September 24 2004 Diamond 0 Pure sp3 carbon 0 All bonds staggered ideal 9 dCC 154 A like ethane I Silicon Germanium Gray Tin 1 TR Diamond struct ure Si and Ge semiconductors Si is the purest element manufactured V 1 Silicon and quot3 5quot Semiconductors 0 Silicon diamondetype is related to the octet 375 semiconductors by replacing Si with eg GaAs M Silicon 3905 Carbon O CC single bonds 356 kImol are much stronger than Si Si bonds 230 k mo 1 o occurrence of coal hydrocarbons is reasonable 0 no natural occurence of SieSi bonds can be synthesized however 0 Stable under standard conditions but unstable wrt 37 Sn below 286 K at 1 atIn pressure L I lt rade picture for Bonding in Diamond 0 Pick one carbon atom and look at its bonds to four neighbor atoms 0 Mix 4 spa orbitals from central atom With one spa orbital from each of the other 4 0 Get 8 new orbitals 4 bonding and 4 antiboncling O Bonding orbitals filled antiboncling empty Why an insulator 39 O A quotbandgapquot exists between the filled and unfilled orbitals 5 sp3 antibonding PW conduction bandquot Band gap energy Filled valenceband SlySpa bondmg O The gap is big the bonding and antiboncling interactions are strong Band Diagrams Energy FEE Metal Semiconductor Insulator xi 9 1 Insulators 0 With a large band gap a lot of energy is needed to promote an electron O Visible light photons too low in energy s diamon is transparent O Electrons can t readily move through material so no electrical conductivity 0 Similar idea for thermal conductivity at normal T only low energy excitation V possible I figMeasuring the Band Gap v Absorbance E9 E9 Energy Wavelength 4 lt7 Energ XL f if Semiconductors 0 If the band gap becomes small enough some conductivity can be achieved 0 Band gaps diamond 580 kImol I N 206 nm silicon 105 kImol I N 1140 11m germanium 64 kImol I N 1870 11m 0 Pure Si or Ge can conduct at high T or if exposed to light V Semiconductors S C Energy Thermal energy 4 promme 6713 i 0 Energy from heat thermal equilibrium gives virtu y BoltZInannelike distribution of excited states where some electrons are promo ed 0 When electrons have been promoted heat light the material will begin to conduct Intrinsicquot pure undopcd Semiconductors 0 Moderate band gaps conductivity is low but increases with temperature enh Keq e AGIRT eASIRe AHIRT AHa AB Egap also e e h z eAS392Ro e EgapZRT O Conductivity is an activated process E3 EgapZ in a pure semiconductor Plot Inc vs lT to get slope EgapZ 1 1 It i Wctals vs Semiconductors Resistivities p 16 a o Metals resistivities increase wi T o Semiconductors resistivities decrease with T l39 i Extrinsicquot Doped Semiconductors 0 Pure elemental semiconductors Si Ge etc can only be used for devices Where light or heat can be supplied to promote electrons 0 More useful devices are made using doped semiconductors appropriate impurities are intentionally added to supply electrons eg P or holes eg Al nType Semiconductors n Dope With phosphorus An electron is leftover after forming SiP bonds 0 The added electrons are easily promoted from the 11donor levels at norma temperatures so they can serve as charge carriers 0 Typical ntype devices contain on the order of 000001 P V iyf Phosphorus doped mto Si O nTypc Semiconductors Adde m donor 4 levels re silicon 0 Initially valence band is full conduction band is em t Energy 0 Added equots must go in conduction band 0 Extent of conductivity depends on of electrons added pType Semiconductors n Dope with aluminum Formation of AleSi bonds steals an electron from Si 0 The holes allow 1 places for electrons to move into within the valence band so they serve as charge carriers 0 Shallow impurity levels as for netype 7 electrons easily promoted at normal temperatures 0 Properties of n amp p type differ slightly Most devices contain combinations of both pTypc Semiconductors e acceptor m Energy pure silicon 0 Initially valence band is full conduction band is empty 0 Removing equot s leaves holes in valence band 0 Number of electrons removed determines conductivity Solid State Synthesis Aspects of Thermodynamic Control Chem 462 September 28 2004 Phase Rule WFC2 o P of phases present at equilibrium 0 C of components needed to describe the system 0 F of degrees of freedom of thermodynamic variables 7 taken from T P compositions of components This full form of the phase rule is used When all three common physical states 5 l g are important V Condensed Phase Rule PFC1insteaclof2 0 When the vapor pressure of all components is negligible the effective number of degrees of freedom can reduced by one 0 Pressure is not an important thermodynamic variable leaving only two kinds variables necessary to describe the solid solution 7 temperature and compositions of the components VA Boring diagram LiF NaF y Discuss the system in every region 9007 m 60 LlF Mole Va NaF NaF VHypothetical T C diagram AQB 7 conquently 00 melting 7 easily synthesized 900 e 7 eutectic point E00 11 7 peritectic point ABS 7 how 500 should it be synthesized w a Wha is likely to be the best way to prepare Ta s z b Whal is likely to be the best way to prepare Tazs z c Whal will occur if anyan when asample ofTaZS is healed from 25 to 1100 quotC a Whal is likely to be the best way to prepare T3552 Explain quotC 31 it CczO SiO2 system This is effectively a two component system even though there are three elements present Why 2L What are the conguently 0 melting compounds What is compound D and Cnsu L ocA T d how might it be prepared r W Suggest structural features that may be present in A CEO 502 C and D m wos og T d dermt n51 baht on quartz C S CCIOSi02 305an 352 3000102m3i207 C25 2000102 mZSIFlA 03 BCuOSIOz 303805 ATOM Cu 00 O Binaryjoin CaOSiO2 in the ternary system Ca SiO Note the method used for labeling ofphases C CaO S SiO2 This type ofabbreviation is Widely used in oxide chemistry 0 ldealized cubic structure of cristobalite form of SiOZ Class 62 Acids and Bases Friday October 8 CHEM 462 T Hughbanks Table 83me folly Modem Imrgrmzc Chemzstry Aquzous pKa values ofthe bmmy hydrzdes ofthe mmmzm s CH4 NH3 H20 HF 44 39 1574 315 SiH4 PH3 H28 HCI 35 27 689 63 GeH4 ASH3 H28e HBr 25 s 23 37 87 H2Te HI 26 93 A cidic and Basic Oxides I The oxides that one uses to form acids and bases in aqueous solution often have reactivity that reflects their acidic or basic character I Examples LiQO CaO and BaO react with water to form basic solutions and can react with acids directly to form salts Likewise SOS 002 and N205 torm acidic aqueous solutions and can react directly with bases to give salts Oxides ds Acid and Basic Anhgdridcs Basic Oxides usually ionic Ca0 2H20 gt Ca2 ZOHi moderately strong base 02 H20 gt 20 K gt1022 Alkali metal and alkaline earth oxides are basic dissolve in acid Acidic Oxides Acid Anhydrides elementoxygen E O bond not broken on dissolution either an E O E group is hydrolyzed by water or water is added across a double bond Acidic Oxides not soluble in water will dissolve in basic aqueous solutions to produce salts eg AsZO3 2NaOHaq gt 2NaH2AsO3 Often seen for anhydrides of weaker acids Examples E quote IE ejg Eo igj H20 265C VH gt 0 dissociation 5H20 gt 4 DO523 HSJEOH OH H20 74 J5O H H Amyhoteric Oxides Dissolve in acids or bases if strong enough Eg BeO SnO certain forms of AIZO3 In strong acids ZnO 2HCIaq gt ZnCl2aq ZnO 2HNOSaq gt ZnOH262 N03 ln strong base ZnO 2NaOHaq gt 2Naaq ZnOH4lZ 61 Lux Flood Concept Oxide Solids I Acid Oxide ion acceptor I Base Oxide ion donor A generalization that includes reactions between solids when water never gets involved Eg Ca0 Sio2 gt CaSiOs 3 NaZO P205 gt 2 Na3P04 NaOH 002 gt Ncho3 Other Oxides Many oxides paiticularly ot the transition metals are difficult to classify as acidic or basic because redox chemistry is more important eg MnO2 4Hl aq conc gt Mn2aq 2 2H20 2 I Hydrolysis of Metal nmpieme can give acidic solutions mil lIumn plg s for MHQOel 3 246 mum ini CHIHzO 6 3 NiH20e 2 1mm livinwi mmi Examples ofACids From S olvolyzed Metals Agua Acids solvolysis Al3 solutions are acidic AIC3s 120 4 AlWaqj criaq wsm arnts of water HCI gas is evolved AICIB HZO 4 AIOH3quot 3HC a AC2OHnH20 complex AICIOH2 mHZO Highly Charged cations with small radii make for stronger acids FeOl 126 2 fairlyweak FeOHZ63 is much stronger rFe3 lt r Fe2 the smaller more highly charged more polarizing cation 39 ws more e densin from coordinated wa er More than size is involved rAl3 lt rFe5 ionic radii but FeOH263 is stronger than AIOH263 Fe LO bondin robably more covalent smaller eledronegativity diff than AlO Acid Base behavior ofMetal Ions not so simple Ions like Al3 and Fe3 are often used as examples of cations that form acidic aqueous solutions however acidbase equilibria tor the ions are not simple Compare CICHZCOOH pKa 285 has a normal titration curve pH vs added base What does it look like An acidic AIOH263 solution will not titrate this way Why AlOH26 3quot is deprotonated then eg AOH263 CHSCOO 2 AOH25OH2 CHSCOOH condensation occurs HZT 2 2W3 iOH2 2HO 20 0 2 H liz H H2 H2 H OH H20Ilt OH2 AIOH3 0 H20 1 a i oH2 H2 H2 Transition Metals in High Ox States Acidic Metals in very high oxidation states form strong largely covalent bonds with oxygen gt weakens OH bonds eg CrO42 weak conjugate base of chromic acid 2 Had gt02Er o cfoo dichromale eg MnO4 very weak conjugate base of permanganic acid both are powerful oxidants Lewis acids and Bases orbital viewpoint 1 Metal cations Lewis acids amp Ligands Lewis bases 2 pblock Lewis Acids Incomplete octets eg BCH33 AICI3 Review of bonding Lewis acidity Trend tor BX3 X F Cl Br A1C13 exists as A12C16 in gas phase Familiar application FriedelCratts acylation To be discussed later Chloroaluminate based Ionic Liquids Molecules with low lying LUMOs I Eg review CO2 and SO3 examples 302 can function as a Lewis acid or Lewis base39 I SO2 as an acid involvement of LUMO I SO2 as a Lewis base use of S localized HOMO or Olocalized lone pairs S O3 is very electrophilic strong Lewis acidity I electronic structure review I Oleum formed by adding SO3 to H2304 Chapter 10 The sp block elements 111 CHEM 462 FriMonNov1215 T Hughbanks Group 13 Elements 6 Electropositive character of elements still very importan 0 Atomic con gurations ns2 np1 0 Bonding in the elements is much stronger and p orbital participation in bonding is important sp2 and sp3 hybridization is common in compounds except for In and especially Tl in oxidation state1 inert pair effect A highly variable group melting boiling Density Form pts C pts C g39cm properties B 2030 3700 247 Brown semicond A1 660 2350 270 Silvery metal Ga 30 2070 591 Silvery soft metal In 157 2050 729 Silvery soft T1 304 1460 1187 soft metal Sourcesuses of Group 13 elements 9 Boron from Na2B4O7nHZO barax 0 Al from electrolysis 0 Ga is is a byproduct of Al production In is a byproduct of PbZn production Tl must be separated from other ue dust elements Ni Zn Cd In Ge Pb As Se Te 0 Ga and In have uses in specialized semicond devices LEDs junction metals dopants in Si amp Ge photoconductors and lovv T solders Al production 553m main Al is made on a massive scale by electrolytic reduction of A1203 from bauxite in cryolite Na3AlF6 at 950 C Al3melt 3e gt All 02 melt Cgraphite gt C02g 4e Mohen cryolite M01 6 and aluman m m nlmlm Some general group trends 0 Boron has chemistry that more covalent With some similarities to silicon in compounds With halogens and oxygen The chemistry of boron rich compounds is unique 0 Al Ga and In chemistry is dominated by oxidation state III Bonds With ligands have polar covalent character 0 T1 has a prominent ox state I chemistry closo n1 cage bonding e7 pairs nido n2 cage bonding e7 pairs arachno n3 cage bonding e7 pairs closo deltahedra Diborane and other boron hydrides o BZH6 mp 465 C bp 93 C Industrial Prep 2 BF3 6 NaHs gt B2H6g 6NaFs BZH6 amp many other boranes are spontaneously ammable in air BZH6 3 02 gt B203s 3H20 AH 72165 kJmol 0 Hydroboration W BZH6 extremely important in organic chemistry HC Brown Nobel 1979 Borane Synthesis KP in 5atz539c dnmclananemigy z l55 hlnnl hp 1839 hp 55 3sz t 203113 EtHiu EH3 392 35H H2 EZH51 8sz Jew ls EiHa EsHv H2 etc Various stable 3sz 1 intemiediates E Hm isolated by careful control ofT amp P Ean winl Unique Boron Compounds 27 H 39Blzmz39 H 239 Bomnrnchcompountk tendency to form BoronIich solids Boron takes on many forms in the solidrstale The iccsahed is l c buildingrblock of allthe forms uululrliul solids This shows acutravmy ofclnsters from 4 layers ofthe simplest boron form The e e Vlgt W boron Very hard and resistant to chemical attack Counting Electrons in oc Rhom B 0 There are two 3c72e bonds N 20 A per B12 unit within each layer 3 O Layers stack over each other in an ABC stacking mode Counting Electrons in oc Rhom B Each of six Batoms in each B12 unit is involved in one 2c 2e bond 176 A Since 10 e are used in intercluster bonding 26 e 212 2 remain for cage bonding CaB6 is a semiconductor 7 Wade s rules satis ed for each B6 unit LaB6 is useful thermionic emitter metallic Boron chemistry general comments 0 Boron chemistry more like silicon than Al eg 7 BOH3 baric acid not at all basic AlOH3 amphoteric but mostly asic 7 3203 and SiO2 more similar in their network structures 7 both acidic react with basic metallic oxides to form silicates and borates 7 BX3 and SiX4 X Cl Br I all readily hydrolyzed AlX3 only partially hydrolyzed 0 Unique boron features delocalized bonding in boron hydrides and remarkably diverse solidstate elemental and boronrich chemistry Sources uses of Boron O BOH3 O Na2B4O7010HZOb0rax Na2B4OSOH4 2HZO kemite 7 huge deposits in Boron CA 0 Detergents soaps enamels herbicides refractory borides Boron is used to strenthen plastics it is much stiffer and lighter than Al Borates used in borosilicate glass eg Pyrex 7 these are lower melting and much more workable than pure SiOz Preparation of Boron 0 Reduction by metals eg by Mg from electrolysis of MgClz 7 contamination by refractory metal borides a problem 0 Electrolytic reduction of borates or tetra uoroborates 7 eg KBFA in KClKF 800 C 7 cheap but purity only N 95 0 Reduction of volatile boron compounds by H2 7 2 BBr3 3H2 7 GHBr B gt 99 Ta lament1000 C 0 Thermal of boron hvdri de amp halides 7 2 BI3 7 mrhombohedral B 312 gt 99 hot Ta lament Boron Trihalides 0 Halides of pblock elements are key intermediates boron no exception 0 mp direct reaction for X Cl Br I 0 B133 13203 3cm2 6sto4 a 2BF3g 3 H3OHSO4quot 3CaSO4s strong BiF bond insoluble CaFZ acidibase rxn 0 BX3 all good Lewis acids acidity BF3 lt BCl3 lt BBr3 Boron Trihalides O Halides exchange is facile Why so easy eg 7 BF3 Bc13 BF2C1 BCIZF typical K N 05 r 2 O BF3 is most versatile reagent in presence of H20 because others completely hydrolyzed BX3 H20 gt BOH3 HX hydrolysis X c1 Br I BX3 ROH gt BOR3 HX alcoholysis X Cl Br I BF3 xs H20 gt 3H O 3 BF 13OH3 heat BFA Very stable anion O BF3 especially useful in FriedeLCraftsilike acylations and alkylations OX0 compounds of Boron O Boric acid BOH3 H20 gt BOH3HZO BOHA H30 at neutral pH get condensation similar to Al chem 3 BOH3 HZO 13303OH4 H30 K 14x 107 Na13303OH4 ZNaOH heat gt Na3B3O3O3 Pyrexprecutsor O Borate Esters in acid 13OH3 3CH30H gt BOCH33 3 H20 BOCH33 is a weaker Lewis acid than BF3 Why BoronNitrogen compounds 9 Other compounds aminerboranes H2 82H5 NEl3 9 H383 NEl3 Solld adduel o BNalkene analogs very reactive HZENH2quot a BENEHU cyclohexane analog 9 Borazlrle 1925 Stock 2 BZH6 NHi a BENEHE rlol benzenerllke e g BENSHS 3 Hcl gtNHZEHC13 NHBCl3 3NaEH4 gt NHBH3 52EZH5g 3NaCl NHBCl3 3RMgCl gt NHBR3 MgC12s BoronNitrogen compounds 9 BomnNitride 120039C 3203 2 mag 92 ENs 3HZOg 1800390 hlghrP ENs hex gt ENs cublc Hexagonal Boron Nitn39de o Layels are graphmc W but stacking places N atom in one layer PW above and below B layels and vice versa Reaction Summary Borie Acid Na2B4O5OH4 39 8 H20 Borax Agt Na2B4O7 BOCR3 H NH4BF4 H20 RCOCI HC1 NH4HF2 V B A Mg or Fe 13203 ROH 02 B2H6 adapted from Figure 1213a CampW Reaction Summary BCl3 B3C13N3H3 B3N3H6 1 B OH PII 4C1 2 HCl C1BNCH o 2R0H 2 3 2 C2H5OH HC1 BOR3 ROH HC1 C13B 1ICH3 HZNCH3 BC13 C12 B2H6 H A12C16 CH33N CH3NH2 C12 C12 H2 C BF3 B3C13N3CH33 13203 B B3C13N3H3 adapted from Figure 1213b CampW Class 63 Lews Acids and Bases Monday October 11 CHEM 462 T Hughbanks Halogens X molecules as weak Lewis acids I Bromine is a deep red liquid and red gas in Lewis basic solvents color shitts I I2 has a fairly low lying 5 orbital formation of I as evidence of Lewis acidity brown color of aqueous solutions Lewis Acidity some Group 14 Compounds I Si X4 very different from carbon analogs eg SiCl4 4H20 gt SiOH4 4HCI rapid Lewis acidbase adducts SiCl47nOH2m likely to be key intermediates eg SiF4 2HF gt SiFB 4H SnCIz Lewis acidity compare HOMO and LUMO with carbenes consider adducts eg SnCIy Ionic Liquids Aprotic but dissociated I A larger number of Ionic Liquids are mixtures of organic salts most often N e N halides and a 3sz V CHs Lewis acid C17 Systems using RlelmCh salts A C393 are especially popular these days AlClSIVIeEtImCl Phase Diagram Liquid range at or above room tempemmreu 100 liquid k 50 it mdi o solid Est T7100 7 7150 exp glass luminous 0 0 0 1 0392 0393 04 05 06 017 0 3 XAICI3 AlClgIVIeEtImCl Mixtures like H30 OH e aq aq AICI3 Cl gt AICI4 gt H20 Key Lewis acidbase equilibrium 2A1Cl439 A12Cl739 c1 Kz103916 71017 21140 quotc Karpinsky amp Osteryoung Inorg Chem 1935 24 2259 I What species dominate when XAICIS XImCI I What species dominate when XAICIS gt XImCI I What species dominate when XAICIS lt XImCI Sugeracids Superacids are defined to be acids as strong as or stronger than pure H2304 Most superacid systems involve more than one component and combine both Lewis and Bronsted acidity ie Bronsted acidity is enhanced by Lewis acidity How to measure acidity in media other than water Detine Hammett aciditytunction H0 defined in terms of indicator bases B for which there is the protonation equilibrium B H k BH conjugate base conjugate acid KBH is the Ka tor the acid BH l 0sz w 1 BH 0g 3 ln dilute solutions K BH 1 gt H 10H H BH BH 0 g p The ratio BHB can be measured spectroscopically tor suitably chosen indicator acidbase pairs eg nitrosamine In aqueous solution S 8M various strong acids have similar values ofH0 Suggests that acidity lt 8M is independent of Anion H90 probablythe main acid species Some superacids pure H2304 H0 12 10 2x 1M H2304 Oleum mixture of HQSO4 and SOS H0 15 HSOSF Fluorosultonic acid H0 15 Using Lewis Acids to Enhance Bronsted Acidity I Group 15 Lewis Acids are important egg PX5 o CH 0 gt PX4 PX I In pure HF SbF5 is F ion acceptor SbF5 2 HF gt SbFS HFH HFH is the acid species a powerful Bronsted acid I F donors CIF3 BrF3 yield bitluoride FHF 11 F species 7 F use Leth F acid to bmd o 50517 F quotMagic Acidquot Example protonations H303gt OH Q F F H303o H2 IO f H H30 H307 H3CHH H O H 30 3 O H 1995 Nobel Prize I gt IL George Olah HCH H H Published extensive series Carbonium i quot Class 2 Electron Con gurations Periodic Properties amp the Periodic Table Wednesday Sem 1 CHEM 462 T Hughbanks Announcements I If I do not have your correct email address please email me quotLL mmailrch nlemdw I 131 homework set is posted mmadumoyplmiohnanks a mu I Lecture notes handouts old exams etc posted in me same place Reading etc I We will continue through Chapter1 loda i I Material related to Chap1er 3 will begin soon b F may e riday Headi oflhe this material Orbital Filling Electron nnligurzniuns I Low energy orbitals iill first I Orbital energy increases as n increases amp I increases I Pauli exclusion principle Electrons can39t have identical quantum nos 2 equots per orbital opposite spins I Hund s rule For lowest total energy all unpaired e s will have the same spin ruximate Electron Con gurations I You should be able to write these easily for any representative elementquot From 5 or p blocks of periodic table Write e configurations for O 8 Ar Use both 152 quotraregasquot and quotarrowquot notations I Some irregularities in transition metals Examples Ti Co Zr W 9 z n a 2 s 4 I z 2 2 z I 2 z z 2 I 1 3 5 s a 7 a III In RI Sr 2 Nb In TI Ru I39llI I39d Ag CI Se 1 1 I I I I I u I z kl I z 4 5 a 7 K II In In 5 a E 39 R m I n u Hg 0 I 65 2 2 5d IJJJ 57mm PzII IImIIgnclism Electrons have magneIic properties 5 have oppos e magnetic momenis so they cancel oulquot each umers magneu sm Molecules wilh y by a magnetic ileld diamagneiism I by magnelic elds and are said to be paramagnetic The Periodic Table I Elemems in a column of lhe periodic table have analogous electron con gurations I Elements in a column of the periodic table lend to have similar chemical propenies Valence electans determine the chemical behavior of an alum Valence Electrons I Valence electrons are those with the highest n value plus any in partially filled dor fshells I These electrons are the farthest from the nucleus and they have the highest energies hus t ey are t e most quotaccessiblequot to other atoms 7 Valence electrons determine the chemical properties of an atom Periodic Properties I Atomic amp Ionic Radius I Ionization Energy I Electron Affinity Atomic Radius I Size of an atom is determined mainly by valence electrons Why I Hard to define and measure I Which would you expect to be larger Na or K N or F Atomic Radius Variations I Moving across a row Z increases while valence electrons are added to the same nshell gt size decreases I Moving down a column the n quantum number of the valence electrons increases a size increases Ionic Rudii I Think about these the same way as for atoms electron configurations I Be sure to use correct of electrons I Which should be larger Mgz F or Fquot Ionic Radii I Anions are always bigger than the corresponding neutral atom I Cations are always smaller than the corresponding neutral atom I For isoelecfranic ions the larger the nuclear charge the smaller the ion Ioniz rion Encrgy I IE amount of energy needed to remove an electron from a tree neutral atom Z energy gt Z e measure experiment similar to photoelectric effect I Tell us the relative stabilities of different orbitals Ionization Energy V 0 down a column at the Periodic Table I as you go across a row I Why I How would you expect IE to vary as ou 9 Ionization Energies EXDIaIn the general il and the variations kJmol Ionization Energies I IE mam going down a column of the periodic table I IE increases going across a row of the periodic table I Some exceptions IFilled shells or subshells are especially stable nHalffilled subshells are also fairly stable Higher Ionization Energies I Can also define and measure higher IE39s 2 quot 1E Z 1 energy a PM e I Can again predict and understand the based on electron configuration d m m of the ions Involve t ionization Encrgies Ionization nei gles dquot twwi m39 Tl39ulhllltin complications Electron Af nity I Measures atom39s tendency to form anions I EA energy released upon adding an electron to a neutral atom 29 e39 a 219 AHEA EA I ll energy is released upon adding an electron EA gt O In the definition above we refer to the Electron A ac ment Enthalpy AHEA which is opposite in sign E iro I Note EA can be positive or negative Electron Affinity I If EA is positive the atom wants to d electron and form an anion ie the process is exot ermic I If EA is negative the atom does not want to add an electron to form an anion ie anion is uns a le I Which elements will have the most positive EA39s hy Periodic Table Mcluls amp Nonmetals I What makes an element a metal or a non mekal Pro enies Electron con guration I How are metais amp nonmetals grouped in the periodic table Comments on the Periodic Tahlc e Mmu E i l i l i i Eu E l l l m MEEEM EYE Nnnmcuh
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