CHEM 1066 2016 Inorganic Module 2 Lectures & Notes PDF

CHEM 1066 INORGANIC MODULE 2 12 LECTURES BONDING THEORY Dr. Fairman Learning Language Examinations At the end of Part 1 you should be able to: a) Draw the shapes of atomic orbitals b) Describe the tenets of molecular orbital theory c) Utilize energy level d...

CHEM 1066 INORGANIC MODULE 2 12 LECTURES BONDING THEORY Dr. Fairman Learning Language Examinations At the end of Part 1 you should be able to: a) Draw the shapes of atomic orbitals b) Describe the tenets of molecular orbital theory c) Utilize energy level diagrams for molecules d) Interpret the molecular orbital overlap integral e) Differentiate bonding and antibonding orbitals f) Infer non-bonded and lone pairs of electrons g) Cite the assumptions of the VSEPR theory h) Predict the shapes of molecules using VSEPR [ Review Chemistry3 Section 2.4 -2.8 ]  Electron radius 9.1x10-17 m  speed length height velocity frequency amplitude power Nobel Prize : 1933 Diffraction Refraction waves bend and mix Superposition Interference : constructive & destructive In 1934, Erwin Schrödinger openly lived with two women (Anny and Hilde), one of whom was herself married to another man.  E=mc2 E=hn l=h/p Einstein Planck DeBroglie ET = Ek + Ep Ek= 1/2 mv2 E=p2/2m + U differential equation for any old wave  Probability of finding electron Max Born: 1954 Nobel prize in Physics Delbrück, Fermi, Heisenberg, Goeppert-Mayer, Herzberg, Pauli, Wigner Notes The wave picture of an electron matches with some experimental results. For example, diffraction or bending of electron beams is quite convincing that electrons are wavelike. Also by experiment (photoelectric effect), the electron transfers a precise amount of energy and momentum just like a tiny ball would. Some persons ascribe a “duality” to the electron but this description is only one of convenience - at the scale of an electron, the world (e-verse) is a just a lot different. For one thing, “matter” becomes indistinguishable from “energy”. Furthermore, moving packets of energy more accurately look like disturbances in some type of field between the source and the destination. In fact, all of modern physics now treats phenomena on the scale of the electron and smaller as field disturbances – the electron corresponding to a disturbance in an “electromagnetic” field. The equation of Schrodinger works well to describe all types of “field disturbances”, including those that carry a fundamental unit of negative charge - the electron.  Electron can be described as a wave We can approximate the function which describes the electron “wave” Parts of the wave-function can be zero (node) positive or negative We can estimate* the region where an electron is likely to be found for most atoms Atomic orbital *within limits (see next slide) Notes Common plots of orbitals are actually plots of wavefunction values that you are likely to measure at certain points around a nucleus. Our precise analytical scientific approach shifts into statistical analysis because of the abundance of evidence that the e-verse actually has MANY interacting, interrelated and fluctuating fields at any particular point in space. It turns out that even the act of measurement will unavoidable add to the melee; we must disturb the e-verse to get information out. We can only be confident, then, that after a given number of measurements we are likely to get a useful average value of the property being measured provided that the disturbance we are interested in is much larger than the noise in the background. It follows also that if even empty space is full of these fields, then there will be a lower limit on the precision with which we can do any measurement. This limit is of the order h/4π≈5x10-34 and gives us one interpretation of the Heisenberg Uncertainty Principle. Of course, such a ridiculously small number is only meaningful on the scale of the e-verse. dxz dyz dyx dzd2-xz22-y2 dx2-y2  Bonding is an interaction between the electrons in atoms A linear combination of electron-waves describes a bond The square of the linear combination gives the new electron density 2 extra electron density 2 reduced electron density Molecular orbital theory  SUMMARY An electron has wavelike properties actually, all particles An electron-wave can be described by a mathematical function wavefunction The square of the electron-wave’s function gives a probability of finding electron’s charge in a certain region of space When atoms get close together, the electron-waves overlap constructively or destructively The resulting waves (LCAO) describe a molecular orbital bonding and antibonding Notes Modelling electrons using wavefunctions (20th century science) gives us access to all the age-old mathematical tools for manipulating functions. The situation of atoms in close proximity translates to overlapping functions. In two dimensions, the area under a curve between some limiting values is its integral. For molecules, we compute a similar integral for the overlapping functions called …(drumroll) the overlap integral. Mixing equal and opposite functions (positive and negative areas) gives zero overlap integral. Mixing a large positive and a small negative gives … small positive.. Mixing a large negative and a small positive gives … small negative.. In three dimensions, the overlap integral is perhaps easier to visualize as the mixing of sets. If the function values which combine as the sets intersect are positive, then the resultant has larger positive values. The other possibilities follow naturally. The situation post overlap therefore has three outcomes: increased wavefunction values → increased probability → bonding interaction decreased wavefunction values → decreased probability → anti-bonding interaction equal increase and decrease → unchanged probability → non-bonding interaction  The overlap integral area under the curves when atomic wavefunctions mix depends on phase of each electron-wavefunction values are positive negative or zero positive values imply enhanced bonding negative values imply reduced bonding zero indicates no impact on bonding  Overlap YES Integral ZERO Bond Effect ZERO Non-bonding Molecular Orbital Non-bonding Molecular Orbital Non-bonding Molecular Orbital  1 1 1 1 1 1 1 1 1 1 1 1 1 1 2 2 2 2 2 2 1 1 1 1 2 2 2 2 2 2 1 1 1 1 2 2 3 3 3 3 2 2 1 1 1 1 2 2 3 3 3 3 2 2 1 1 1 2 2 3 4 4 4 4 3 2 2 1 1 2 2 3 4 4 4 4 3 2 2 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 2 2 3 5 7 11 11 7 5 3 2 2 1 1 2 2 3 5 7 11 11 7 5 3 2 2 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 2 2 3 4 4 4 4 3 2 2 1 1 2 2 3 4 4 4 4 3 2 2 1 1 1 2 2 3 3 3 3 2 2 1 1 1 1 2 2 3 3 3 3 2 2 1 1 1 1 2 2 2 2 2 2 1 1 1 1 2 2 2 2 2 2 1 1 1 1 1 1 1 1 1 1 1 1 1 1  1 1 1 1 1 1 1 1 1 2 2 2 1 1 1 1 1 1 2 2 2 2 2 2 2 2 2 1 1 1 2 2 2 3 3 3 3 3 3 3 2 2 2 1 1 2 2 3 3 4 4 4 4 4 4 4 3 3 2 2 1 1 1 2 3 4 5 5 6 6 6 6 6 5 5 4 3 2 1 1 1 2 2 3 5 6 7 8 8 8 8 8 7 6 5 3 2 2 1 1 2 3 4 5 8 10 11 10 9 10 11 10 8 5 4 3 2 1 1 2 3 4 6 9 12 13 11 10 11 13 12 9 6 4 3 2 1 1 2 3 4 5 8 10 11 10 9 10 11 10 8 5 4 3 2 1 1 2 2 3 5 6 7 8 8 8 8 8 7 6 5 3 2 2 1 1 1 2 3 4 5 5 6 6 6 6 6 5 5 4 3 2 1 1 1 2 2 3 3 4 4 4 4 4 4 4 3 3 2 2 1 1 2 2 2 3 3 3 3 3 3 3 2 2 2 1 1 1 2 2 2 2 2 2 2 2 2 1 1 1 1 1 1 2 2 2 1 1 1 1 1 1 1 1 1 overlap integral zone  1 1 1 1 1 1 1 1 1 1 1 1 1 1 2 2 2 2 2 2 1 1 1 1 2 2 2 2 2 2 1 1 1 1 2 2 3 3 3 3 2 2 1 1 1 1 2 2 3 3 3 3 2 2 1 1 1 2 2 3 4 4 4 4 3 2 2 1 1 2 2 3 4 4 4 4 3 2 2 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 2 2 3 5 7 11 11 7 5 3 2 2 1 1 2 2 3 5 7 11 11 7 5 3 2 2 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 2 2 3 5 7 9 9 7 5 3 2 2 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 1 2 3 4 5 6 6 5 4 3 2 1 1 1 2 2 3 4 4 4 4 3 2 2 1 1 2 2 3 4 4 4 4 3 2 2 1 1 1 2 2 3 3 3 3 2 2 1 1 1 1 2 2 3 3 3 3 2 2 1 1 1 1 2 2 2 2 2 2 1 1 1 1 2 2 2 2 2 2 1 1 1 1 1 1 1 1 1 1 1 1 1 1  1 1 1 1 1 1 1 1 2 2 2 2 2 2 2 2 1 1 1 1 2 2 3 3 2 1 1 2 3 3 2 2 1 1 1 2 2 3 4 5 4 2 2 4 5 4 3 2 2 1 1 2 3 4 6 7 6 3 3 6 7 6 4 3 2 1 1 2 3 4 6 9 8 4 4 8 9 6 4 3 2 1 1 2 3 4 6 7 6 3 3 6 7 6 4 3 2 1 1 2 2 3 4 5 4 2 2 4 5 4 3 2 2 1 1 1 2 2 3 3 2 1 1 2 3 3 2 2 1 1 1 1 2 2 2 2 2 2 2 2 1 1 1 1 1 1 1 1 overlap integral zone Notes Some fancy mathematical techniques allow for the calculation of electron wavefunctions around small nuclei with good agreement to experimentally determined electron energies. The wavefunctions are evaluated within small volumes at different locations around the nucleus and plotted, usually in false colours for positive and negative values, in a manner that gives a three dimensional feel for the resultant map (projection). There a couple of important features of such plots: a mimimum value of the function is usually specified (cut-off or threshold) and shown as the border of the plot. This border does not represent zero values of φ, just the smallest value that is calculated. orbitals computed in similar fashion are usually normalized. This is a system of rescaling all of the values so that their maxima and minima are the same. The resulting electron density distributions are then comparable to each and zones of high/low electron density stand out. Some calculated wavefunctions and their corresponding quantum numbers are shown below. Notes Calculations become much more complicated for systems with many electrons because the electrons are not really independent of each other – they interact (repel) there are restrictions on electrons in the same general area (state) at any one time However, with a number of important assumptions, modern software can give good approximations to wavefunction values around all the nuclei in sizeable molecules. Arnab Rit, A; Campos, J; Niu, H; Aldridge, S; Nature Chemistry, 8, 1022–1026 (2016)  Classifying Orbitals Atomic Molecular sg s-orbitals s – sigma type bonding molecular orbitals pu p-orbitals p – type bonding molecular orbitals d – delta dg d-orbitals  Molecular Orbital Energy Level Diagrams Energy (kJ/mol) 400 +436.4 Anti-bonding Molecular orbitals on H2 200 0 -200 Atomic orbitals on Hydrogen -436.4 -400 Bonding Molecular orbitals on H2 *bonding and antibonding are exactly opposite in energy  Molecular Orbital Energy Level Diagram of Fluorine Energy (kJ/mol) s* p* 2p 2p isoenergetic p degenerate s s* 2s 2s s s* 1s 1s s F F  Molecular Orbital Energy Level Diagram of Fluorine Energy (kJ/mol) 2p 2p Auf bau 2s 2s Hund Pauli 1s 1s F F  Molecular Bond Order = (bonding–antibonding) Energy (kJ/mol) 2 2p 2p Bonding = 10 AntiBonding = 8 2s 2s Bond Order=1 1s 1s F F  Molecular Orbital Energy Level Diagram of Dinitrogen Energy (kJ/mol) 2p 2p Auf bau 2s 2s Hund Pauli 1s 1s N N  Molecular Orbital Energy Level Diagram of Dinitrogen Energy (kJ/mol) 2p 2p Bonding = 10 AntiBonding = 4 2s 2s Bond Order = (10-4)/2=3 1s 1s N N  Molecular Orbital Energy Level Diagram of Hydrogen Fluoride Energy (kJ/mol) s* 1s Bond Order = 2p p (2-0)/2 = 1 s 2s s F H HOMO Highest Occupied Molecular Orbital LUMO Lowest Unoccupied Molecular Orbital  Molecular Orbital Energy Level Diagram of Dinitrogen Energy (kJ/mol) LUMO 2p 2p HOMO 2s 2s 1s 1s N N  Molecular Ions 1. Draw the energy level diagram for the neutral molecule 2. Add or remove electrons from HOMO/LUMO Molecular Orbital Energy Level Diagram of Dinitrogen Cation 2p 2p N2 N2 + 2p 2p HOMO Molecular Orbital Energy Level Diagram of Dinitrogen Anion 2p 2p N2 LUMO N2 - 2p 2p  The mixing of atomic orbitals to facilitate bonding Wavefunctions of different atomic orbitals share space The energy of orbital overlap between atoms is reduced Notes Something old, something new, something borrowed … We make sense of the observation that an atom can make bonds to other atoms in different geometrical arrangements with the idea that the central atom or ion senses the arrival of their bethroded on the big day (ignore the polygamous implications) and makes some preparations. Firstly, there needs to be an equal number of equivalent single electron orbitals as there are number of suitors ! Remember, each suitor brings along a single electron to form a bond. But how difficult is it to prepare such identical orbitals ? Our interpretation continues with the notion that mixing of functions is easy : 1. the functions are really related to probability maps or statistical distributions anyway 2. we only need consider the linear combinations (addition and subtraction) Energy (kJ/mol) Be H 2p 2p 2s 2s 1s 1s Energy (kJ/mol) Be Be H 2p isoenergetic 2p sp 2s 2s 1s 1s 1s  BeH2 Both hydrogen atoms bond to the same type of orbital a hybrid Orbital overlap is enhanced stronger bonding  VSEPR Valence Shell outer electrons only Electron Pair a molecular orbital can only have 2 electrons Repulsion Theory electrostatic repulsion pushes electron pairs as far apart as possible Common Geometries of Molecules Regular solids are defined by Vertices (Intersection) points Faces facet Edges surface contacts  4 5 6 7 Tetrahedron Trigonal bipyramid Octahedron Pentagonal bipyramid 12 8 8 Square antiprism Dodecahedron 9 9 Icosahedron Mono-capped square antiprism Tri-capped trigonal prism USING VSEPR The central atom is electropositive Count number of ligands Hybridize Arrange ligands so they are evenly spaced Determine observed shape of molecule How to Hybridize and Spatialize 1. Write configuration of central atom 2. Extract valence electrons 3. Promote electrons to equal number of ligands (Assume energy sequence for orbitals is s, p, d, f) 4. Name hybrid orbitals 5. Count all valence electron pairs that result 6. Determine the best arrangement of pairs 7. Describe overall shape of molecule  CH4 1. Write configuration of central atom 1s2 2s2 2p2 2. Extract valence electrons 2s2 2p2 2s 2p 3d sp3 s1p3 4. Name hybrid orbitals 3. Promote electrons to equal number of ligands 5. Four valence pairs 6,7. Tetrahedron  PCl5 P : 1s2 2s2 2p63s2 3p3 3s2 3p3 3s 3p 3d s1p3d1 sp3d Trigonal bipyramidal  NH3 N : 1s2 2s2 2p3 2s 2s2 2p3 2s 2p 2p ‘Sum’= sp Close to 2s Close to 2p s1p3 sp3 Four electron pairs Tetrahedral Trigonal pyramidal  SF6 S : 1s2 2s2 2p63s2 3p4 3s2 3p4 3s 3p 3d s1p3d2 sp3d2 Six electron pairs Octahedral Octahedral  SF4 S : 1s2 2s2 2p63s2 3p4 3s2 3p4 3s 3p 3d s1p3d1 sp3d equatorial Five electron pairs axial Trigonal bipyramidal major axis Saw-horse Lone pair electrons experience less nuclear pull Lone pair electrons in equatorial positions are slightly farther away from neighbours Lone pairs preferentially occupy equatorial positions  Lone pair push Bonded pair squeeze  ClF3 Cl : 1s2 2s2 2p63s2 3p5 3s2 3p5 3s 3p 3d s1p3d1 sp3d Five electron pairs Trigonal bipyramidal T-shaped  XeF2 Xe : [Kr]4d10 5s2 5p6 5s2 5p6 5s 5p 5d s1p3d1 sp3d Five electron pairs Trigonal bipyramidal Linear  Other ligand groups PF3(CH3) 2 Alkyl groups need one electron each Phosphorus must supply five in total Trigonal bipyramidal Bulky methyl groups in equatorial positions (more room)  XeF2 Xe : [Kr]4d10 5s2 5p6 5s2 5p6 5s 5p 5d s 1p3d1 sp3d Five electron pairs Trigonal bipyramidal Linear  Other ligand groups PF3(CH3) 2 Alkyl groups need one electron each Phosphorus must supply five in total Trigonal bipyramidal Bulky methyl groups in equatorial positions (more room) Magnetic Properties Number of unpaired electrons Paramagnetic - unpaired electrons Diamagnetic - no unpaired electrons i.e. all are paired Notes Although useful, VSEPRT has some severe limitations: 1. the energy values associated with pure atomic orbitals vary significantly according to electron configuration and quite dramatically for heavy atoms. This is due to electron screening effects, the additional stability observed for configurations with half-filled subshells and changes in effective nuclear charge that result from these combined. 2. there is no way to predict the extent of distortion in bond angles that is likely for systems with lone pairs of electrons. 3. only sigma bonded electron pairs (and lone pairs) are considered to arrive at the molecular geometry 4. Many molecules with sideways overlap cannot be modelled with VSEPRT  ABC : Cubic Closest Packing Staggered arrangement Octahedral hole Octahedral hole  THE UNIT CELL NOT UNIT CELL Smallest cell that will reproduce the structure in three dimensions Notes If we could dissect a crystal structure to obtain the single repeat unit which reproduces the entire structure, we might end up with shapes like those shown below. The fractions and stoichiometry are evident. However, with more complicated structures, these units rapidly become too crowded to be easily interpreted. To get around this, unit cells are usually shown as simple frames that connect the centres of the atoms/groups in the crystal. Tetrahedral Hole Octahedral Hole  STOICHIOMETR Y Cell Shared by # Per Cell Caesium 1 1 1 Chloride 8 8 1 Total Caesium = 1 x 1/1 =1 Total Chloride = 8 x 1/8 =1 STOICHIOMETRY 1:1 CsCl Fluorite Fluoride Calcium  STOICHIOMETRY 4:8 Ca4F8 CaF2 Cell Shared by # Per Cell Calcium(corners) 8 8 1 Calcium(faces) 6 2 3 4 Fluoride 8 1 8 8 Fluoride ions in all tetrahedral holes Sulfur Zn ZnS Anti-Fluorite Blende  Rutile TiO2 Cell Shared by # Per Cell Titanium(corners) 8 8 1 Titanium(centre) 1 1 1 2 Oxygen(inside diagonal) 2 1 2 Oxygen(faces) 4 2 2 4 Wurtzite Hexagonal closest packed Fajan’s Rules (1923) The amount of covalency in an ionic bond 1. Small cation, high charge 2. Large anion, small charge 3. Electronic configuration Be2+ θ=6 4 Charge density θ = Z+/r Al3+ θ=6 0 Effective nuclear charge  radius(pm) CN Siemens Tin Sn0 140.5 12 9x106 Tin(II)Chloride Sn2+ 112 6 22 Tin(IV)Chloride Sn4+ 83 6 0 High charge density gives covalent bonding ! 3. Effect of Configuration on Ionic /Covalent bonding Shielding of electrons varies with electron configuration Efficiency of shielding depends on shape and occupancy of orbitals and penetration of nuclear region Transition series elements a) smaller than expected b) dispersed d-orbitals (poor screening) c) increased effective nuclear charge Cation radius(pm) Covalency Configuration sHg2+ 102 ++++ [Xe]4f145d10 Ca2+ 100 + [Ar]4s2 Na+ 97 + 1s22s22p6 Cu+ 96 ++ [Ar]3p63d10 Anions Se2+ I- Te2- pm 170 206 207 Soft (easily polarized) Electronegativity χ The power of an atom in a molecule to attract electrons to itself a) Difficult to measure precisely (empirical approach used) b) Pauling - thermochemical data condensed c) Allred–Rochow - 3590 Z / r2 + 0.744 d) Sanderson -considers electron density distribution also  PAULING ELECTRONEGATIVITIES I II III IV V VI VII VIII 2.20 0.98 1.57 2.04 2.55 3.04 3.44 3.98 3.20 0.93 1.31 1.61 1.90 2.19 2.58 3.16 4.70 0.82 1.00 1.81 2.01 2.18 2.55 2.96 3.60 0.82 0.95 1.78 1.96 2.05 2.10 2.66 3.10 0.79 0.89 2.04 2.33 2.02 2.00 2.20 2.50 0.70 0.90 1.8 a) ionic compound : ∆ χ ≈2 b) AlBr3 2.96-1.61 =1.35 Intermediate; polar bond c) CH4 2.55-2.20 =0.35 Covalent bond 1.2 Limitations of χ a) Varies with oxidation state bp c Dc* (oC) Sn2+ Sn4+ Fe2+ Fe3+ Pb2+ Pb4+ SiF4 2.08 -99 1.80 1.96 1.83 1.96 1.87 2.33 BF3 1.94 -86 b) Bond sensitive i. Halides have lower values when bonded to non- metals More covalent ii. Exceptions: Lithium Iodide is ionic Enthalpy (H) Heat content a) Zero for elemental forms eg. solid sulfur b) Change in heat content is DH c) Under standard conditions, DHo d) Products – Reactants for chemical processes e) Negative DH, heat evolved to surroundings DHodiss Dissociation of reactant molecule to atoms DHoEA Addition of one mole of electrons to one mole of gaseous ions DHof Total heat change when one mole of product is formed DHofus Melts one mole of solid DHoIE Converts one mole of gaseous atoms to ions DHoIatt Converts one mole of gaseous ions to ionic solid DHosolv Converts one mole of gaseous ions to solvated form DHosub Converts one mole of solid directly to gas DHovap Converts one mole of solid to liquid then gas Lattice Energy A Madelung’s Constant A e electron charge NaCl 1.74756 Z highest common factor of CsCl2 1.76267 charges Blende 1.63806 ε Dielectric constant, vacuum Wurtzite 1.64132 B Repulsion constant Fluorite 5.03878 d Separation of ions n Born Exponent 1st Hydration Sphere 2nd Hydration Sphere  Dissolution of an ionic compound in a solvent requires energy to separate oppositely charged species : solvation energy or hydration energy(water) Heat of solution (L) is combined solvation and lattice energies Hess’s Law: Path independence [First Law] Conductors Semiconductors and Insulators Chem3 p234 Covalent Bonds Ionic Bonds Polar Bonds Metallic Bonds 4 AMD Opteron(16 Core) + 4 NVIDIA Tesla GPU (2688 Cores) Does extended conjugation cause a red shift or blue ?  http://hyperphysics.phy-astr.gsu.edu/hbase/solids/band.html Band Theory of Solids Valence Band – electrons held to individual atoms Conduction Band – electrons free to sample entire system Insulator – energy gap is very large  Fermi Level : the upper limit of filled valence electron shells at a particular temperature (usually half-way in gap) 1 𝑓 𝐸 = (𝐸−𝐸 ) 𝐹 𝑒 𝑘𝑇 +1 Electron Density of States 3 8 2𝜋𝑚2 𝐸 − 𝐸𝑢𝑝𝑝𝑒𝑟 𝜌 𝐸 = ℎ3 Population in conduction band is product Solar Cell P-type Doped Semiconductor N-type Doped Semiconductor B Al Ga In Tl P As Sb Bi Czochralski Process Doping Silicon  p n Sb+ B- Sb+ B- Depleted region  + - DIODE + - - + Reverse bias Forward bias  p n Sb+ B- Sb+ B- Depleted region  + - DIODE + - - + Reverse bias Forward bias  Bipolar transistor +7 +5 Base +10 -10 +10 -10 Emitter Collector +7 Base +5 Base +10 -10 +10 -10 Emitter Collector Emitter Collector + - p n p Base  Metal Oxide n n Semiconductor Field p Effect Transistor -5 +2 +5 Electric field effect Operating as switches, each of these gate components can sustain a blocking voltage of 120 volts in the OFF state, and source drain can conduct a continuous current of 30 amperes in the ON state, dissipating up to about 100 watts and controlling a load of over 2000 watts. Coordinate-Covalent Bonding M=Metal L=Ligand Donor Group Donor Atom Coordination Geometry Geometrical Isomerism Optical Isomerism Linkage Isomerism  Atoms with capacity for Dative Bonding Nitrogen Oxygen Phosphorus Sulfur Arsenic Selenium Halides We exclude carbon-metal bonding (sigma)  Lone pair from donor atom occupies LUMO of metal in a covalent–type bond  No oxidation or reduction occurs  Electron density on metal is increased Naming Conventions  ML6  M(L)6  M[L]6  [M]L6  [M][L]6  [M(Lx)]Ln  [M(Lx)][P(Ln)]y  Some common molecular ligands pyridine 2,2-bipyridine terpyridine N donor N donor N donor monodentate bidentate tridentate 1,10-phenanthroline N donor bidentate  1,2-diaminoethane 2-hydroxybenzaldehyde N,N'-Ethylenebis(salicylimine) + H2O SALEN + H2O Donor atoms: N,O Denticity: tetra [Cr(salen)(CN)2]2[Mn(phen)2] phen salen salen phen  trans cis facial meridonal  Λ ∆ levo dextro [Co(NH3)4(NO2)Cl]+ [Co(NH3)4(ONO)Cl]+ Nitro Nitrito Linkage Isomers Energy Free Ion Fe2+ d6 Low spin High spin diamagnetic paramagnetic  INORGANIC LINKAGE ISOMERISM OF THIOCYANATE ION Inorganic Chemistry Vol. 3, No. 11, November, 1964 p1587 Complex Colour M.P. ( C ) C-N (cm- C-S (cm-1) 1) [Pd(As(C6H5)3)2(SCN)2] Yellow- 195 (130) 2119 - orange [Pd(As(C6H5)3)2(NCS)2] Yellow 195 2089 854 [Pd(bipy)(SCN)2] Light orange 270 (120) 2108 700 [Pd(bipy)(NCS)2] Light Yellow 270 2100 849 Reciprocal Centimeters = number of wave cycles that cover 1 centimeter 2119 cm-1 = 63526021850200 Hz 2100 cm-1 = 62956416180000 Hz Johannes Diderik Van Der Waals Condensation of gases to liquids Attractive force between neutral molecules Probability of finding electron at a specific point 1837-1923 (85 years) very slight momentary positive charge  London (Dispersion) Force Instantaneous dipole due to uneven electron density around nucleus Vector addition Increases with size of atoms Increases with polarizability of molecules Attractive force varies as 1/r6 Induced-Dipole Induced-Dipole Non-polar to covalent compounds Dipole-Induced Dipole Polar to covalent compounds Dipole-Dipole Polar to polar compound Ion-Induced Dipole Ionic to covalent compound  Hydrogen Bonding Chem3 p1146 D-H---A D=donor A=acceptor 10% of covalent bond energy “Copyright. Blueberg taken by Captain W.F. Wood S.S. Etonian on 12/4/12 [April 12, 1912] in Lat 41° 50 N Long 49° 50 W. Titanic struck 14/4/12 [April 14, 2012] and sank in three hours.” Clathrate : water cage Thermodynamic Effects of Hydrogen Bonding Thermodynamic Effects of Hydrogen Bonding