Presentation on theme: "What we have learned so far toward molecular structure and properties."— Presentation transcript:
1 What we have learned so far toward molecular structure and properties. Model of the atomThe nuclear model of atomThompson’s modelRutherford’s modelAtomic spcetra -- Bohr’s modelQuantum theoryPeriodicity of atomic propertiesWave-particle dualityUncertainty principleWave function – particle in a boxSchroedinger equationAtomic radiusIonic radiusIonization energyElectron affinityPeriodic tableHydrogen atom(one electron atoms)Principle quantum numberAtomic orbitals: radial wave functionangular wave functionorbital angular momentummagnetic quantum numbersradial distribution functionshape of atomic orbitalselectron spinMany electron atomsorbital energy splitshielding effecteffective nuclear chargePauli exclusion principleHund’s rulevalence shell (electrons)
2 MOLECULAR SHAPE AND STRUCTURE What we have learned so far toward molecular structure and properties.Chemical bondInteraction between two electronsIonic bondLewis structureOctet ruleExceptions to octet ruleResonanceFormal chargeOxidation numberCovalent bondelectronegativityIonic v.s. covalentDipole momentPolar bondNonpolar bondBond strengthBond lengthIR(infrared) spectroscopyMOLECULAR SHAPE AND STRUCTUREStability, reactivity, color, size, polarity, solubility, function etc…
3 3D structure of a molecule is crucial for its property. Sophisticated quantum mechanical calculations are neededto predict the structure.⇒ Drugs by Design and DiscoveryBox 3.11) Identification of key enzymes2) Molecular structure determination3) Hints from Nature --- Natural Products4) Computer-aided design of moleculeswith structures fitting into the active site
4 MOLECULAR SHAPE AND STRUCTURE Chapter 3.MOLECULAR SHAPE AND STRUCTURETHE VSEPR MODEL (전자쌍 반발 모델)3.1 The Basic VSEPR Model3.2 Molecules with Lone Pairs on the Central Atom3.3 Polar MoleculesVALENCE-BOND THEORY (원자가 결합 이론)3.4 Sigma and Pi Bonds3.5 Electron Promotion and the Hybridization of Orbitals(혼성궤도 함수)3.6 Other Common Types of Hybridization3.7 Characteristics of Multiple Bonds2012 General Chemistry I
5 THE VSEPR MODEL (Sections 3.1-3.3) 95Lewis structure:showing the linkages between atoms and the presence of lone pairs,but not the 3D arrangement of atomsClF3CH4H2OBF3NH3BeCl2SF4XeF4PCl5IF5SF6
6 Electron pairs (lone pairs & bonding pairs) repel each other. Estimating the 3D structure: THE VSEPR MODEL3.1 The Basic VSEPR ModelValence Shell Electron-Pair Repulsion theoryElectron pairs (lone pairs & bonding pairs) repel each other.Proposed by R. J. Gillespie in 1959.Rule 1: Electron pairs move as far apart as possible.VSEPR structures for AXn with no lone pair
9 Rule 2: (Almost) No distinction between single and multiple bonds. BeCl2CO2CO32-BF3
10 3.2 Molecules with Lone Pairs on the Central Atom 98Rule 3 All regions of high electron density, lone pairs and bonds, are included in a description of the electronic arrangement,But only the positions of atoms are considered when identifying the shape of a molecule.NH3CH4
11 Rule 4 The strength of repulsions are in the order 99Rule 4 The strength of repulsions are in the orderlone pair-lone pair > lone pair-atom > atom-atomH2ONH3
13 Predicting a molecular shape of XeF4 101Predicting a molecular shape of XeF4Step 1 Draw the Lewis structure.Step 2 Assign the electron arrangementaround the central atom.Step 3 Identify the molecular shape. AX4E.Step 4 Allow for distortions.Square planar
15 AXE methodA; central atomX; outside atomE; lone pair
16 1023.3 Polar MoleculesPolar molecule: a molecule with a nonzero dipole momenti.e. HCl with a dipole moment of 1.1 DHCl, H2O, CHCl3, cis-dichloroethane, ···- A polyatomic molecule is polar if it has polar bonds arranged inspace in such a way that the dipole moments associated withthe bonds do not cancel.polarpolar
17 Homonuclear diatomic molecules 1023.3 Polar MoleculesNonpolar molecule: a molecule with a net zero dipole momentHomonuclear diatomic moleculesPolyatomic molecules with symmetry;CO2, BF3, CH4, CCl4, trans-dichloroethane, ···nonpolarnonpolar
19 VALENCE-BOND THEORY (Sections 3.4-3.7) 105VALENCE-BOND THEORY (Sections )Lewis model of the chemical bond; localized electron modelValence-bond theory; Walter Heitler, Fritz London (1927)Linus Pauling (1931)Quantum mechanical description of the distribution of electrons in bondsValence electrons are localized either between pairs of atoms or on atoms as lone pairs.Hybridization of atomic valence orbitals with proper symmetrythat are localized between pairs of atoms.2) Placing valence electrons in the hybridized orbitals as pairs (↑↓) or leaving them localized in lone-pair orbitals on individual atoms in the molecule.VSEPR theory is a simplified one : powerful way of predicting the shape of simplemolecules. --- does not explain many things including multiple bond, bond angles ……..
20 3.4 s(Sigma) and p(Pi) Bonds: description of covalent bond 1053.4 s(Sigma) and p(Pi) Bonds: description of covalent bondH21) Two hydrogen 1s-orbitals merge (overlap) to form a s-orbital between the two hydrogen atoms.Walter Heitler, Fritz London (1927)2) A s-bond is formed as two electrons (↑↓) fill the s-orbital.s-bond ;cylindrically symmetrical with no nodal planes containing the intermolecular axis.i.e. H2, 1s-1s HF, 1s-2pz N2, 2pz-2pz
21 overlap with 1:1 match of orbitals 106overlap with 1:1 match of orbitals
22 overlap with 1:1 match of orbitals 106overlap with 1:1 match of orbitals
24 nodal plane containing the interatomic (bond) axis 106p-bondnodal plane containing the interatomic (bond) axis- two cylindrical shapes (lobes), one above and the other below the nodal planeN2one s-bondwith twoperpendicularp-bonds- multiple bonds: single bond (one s-bond),double bond (one s- and one p-bond)triple bond (one s- and two p-bonds)
25 3.5 Electron Promotion and the Hybridization of Orbitals 107Polyatomic moleculesLinus Pauling (1931)BeH23.5 Electron Promotion and the Hybridization of OrbitalsWhy do we have to make hybrid orbitals?
26 Linear combination of orbitals not real.localization problemLinear combination of orbitalssp hybrid
41 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory113Paramagnetic O2; unpaired electron(s)Lewis's theory;Valence-bond theory;bond and bond2 lone pairs on each O occupying the sp2 hybrid orbitalsParamagnetic: tendency to move into the magnetic field.When there are unpaired electrons in the molecule.Diamagnetic: tendency to move out of the magnetic field.When all the electrons in the molecules are paired.
42 Shortcomings of the Valence Bond Model 3.8 The Limitations of Lewis’s Theory& Valence Bond TheoryShortcomings of the Valence Bond ModelInadequate treatment of odd-electron molecules and resonancesO2N22) Magnetism of molecules• Paramagnetic: molecules withunpaired electrons• Diamagnetic: weakly repelledby a magnetic fieldboth are expected to be diamagnetic!!
43 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory113Electron deficient diborane, B2H6First published by H. C. Lunguet-Higgins,a 2nd year undergraduate student !does not have enough electrons !At least seven bonds (= 14 electrons) are required,but only 12 valence electrons.- No simple explanation for spectroscopic properties of compounds
44 MOLECULAR ORBITAL THEORY (Sections 3.8-3.12) 113MOLECULAR ORBITAL THEORY (Sections )3.9 Molecular Orbitals3.10 Electron Configurations of Diatomic Molecules3.11 Bonding in Heteronuclear Diatomic Molecules3.12 Orbitals in Polyatomic MoleculesMolecular Orbital (MO) theory advantages- Addresses all of the above shortcomings of VB theory- Provides a deeper understanding of electron-pair bonds- Accounts for the structure and properties of metals and semiconductors- Universally used in calculations of molecular structures
45 H2+: Prototype Molecular Orbital System • Atomic orbital(AO) theory → successful for orbital structures of all atomswith both even and odd numbers of electrons• Assume that molecule H2+ ~ as an united atom with a fragmentednucleus if the nuclei in molecule were fused together• construct the one-electron orbital corresponding to the arrangement ofnuclear charges presented by the moleculeCoulomb interactions in H2+
46 3.9 Molecular Orbitals115Quantum mechanics : the ideal solution to the problem, but…….Even for the smallest molecule, H2Schroedinger equation will look like…..and way too complex and complicated… So, need simplification !Simplification 1 (Born - Oppenheimer Approximation) Simplification 2 (Orbital Approximation) Simplification 3 (LCAO Approximation)
47 3.9 Molecular Orbitals115The valence-bond (VB) and molecular orbital (MO) theories are both procedures for constructing approximate wavefunctions of electrons.- In VB theory, bonding electrons are localized on atoms or between pairs of atoms.Molecular orbitals (MOs)The MO theory can account for electron-deficient compounds, paramagnetic O2, and many other properties by focusing on electrons delocalized over the whole molecule.MOs formed by linear combination of atomic orbitals (LCAO-MO)Approximate molecular wavefunctions by superimposing (mixing) of N atomic orbitalscij and Ej are determinedby solving the Schrödinger equation
48 Trial wavefunctions for H2 using two 1s atomic orbitals of H Increased amplitude in the internuclear region bondingLarger volume for electrons lower kinetic energy(particle-in-a-box)Decreased amplitude in the internuclear region & nodal plane antibonding
49 115Molecular orbital energy-level diagram- relative energies of original AOs and resulting MOs- arrows to show electron spin and location of the electrons- In H2, two 1s-orbitals merge to formthe bonding orbital s1s and the antibonding orbital s1s*
50 3.10 Electron Configurations of Diatomic Molecules 116Building-up principle for MOValence electrons in molecular orbitals1. Electrons are accommodated in the lowest-energy MO, then inorbitals of increasingly higher energy.2. Pauli exclusion principle:each MO can accommodate up to two electrons.If two electrons are present in one orbital, they must be paired.3. Hund’s rule:If more than one MO of the same energy is available,the electrons enter them singly and adopt parallel spins.H2:The energy of H2 is lower thanthat of the separate H atoms.Even the energy of H2+ is lower thanthat of the separate H atoms.
51 H2 He2 He2+ H2+ Bond order = 0 Bond order = 1 Bond order = ½( # of bonding electrons - # of antibonding electrons)He2+H2+Bond order = 1/2Bond order = 1/2
52 For other homonuclear diatomic molecules of Period 2 elements, Linear combination of 10 atomic orbitals;1. No mixing between AO's of the same atom2. Significant mixing only between AO's of similar energies and substantial overlap⇒ Negligible mixing between the core 1s and the valence 2s and 2p orbitals⇒ No MO from 2s–2p mixing due to symmetry- two 2s orbitals (one on each atom) overlap to form two s orbitals,one bonding (s2s-orbital) and the other antibonding (s2s*-orbital)- six 2p orbitals (three on each atom) overlap to form six MOs,two 2pz orbitals to form bonding and antibonding (s2p, s2p*)four 2px, 2py orbitals to form two p2p and two p2p* orbitals
55 antibonding s2p* orbitals 116four 2px, 2py orbitals toform two p2p andtwo p2p* orbitalstwo 2pz orbitals to formbonding s2p andantibonding s2p* orbitalsone bonding (s2s-orbital)and the otherantibonding (s2s*-orbital)
56 - From Li2 to N2, the energy levels of 2s and 2p are close, and thus the 2s orbital also participates in forming s2p orbitals.
57 - For O2 and F2, the energy levels of 2s and 2p are separated well.
58 118- In N2, each atom supplies five valence electrons.A total of ten electrons fill the MOs.The ground configuration is,Bond order (b): net number of bondsb = ½(8-2) = 3- In O2, each atom supplies six valence electrons.A total of twelve electrons fill the MOs.The ground configuration is,b = ½(8-4) = 2accounts for paramagnetism of O2
59 118Bond order (b) =112321paramagneticdoes not exist
61 3.11 Bonding in Heteronuclear Diatomic Molecules 1203.11 Bonding in Heteronuclear Diatomic MoleculesA diatomic molecule built from atoms of two different elementsin polar, with the electrons shared unequally by the two atoms.- In a nonpolar covalent bond, cA2 = cB2- In an ionic bond, the coefficient belonging to one ion is zero.In a polar covalent bond,the AO belonging to the more electronegative atom has the lower energy, and so it makes the larger contribution to the lowest energy (bonding) MO.Conversely, the contribution to the highest-energy (most antibonding) orbital is greater for the higher-energy AO, which belongs to the less electronegative atom.
62 less electronegative atom Fig 3.33more electronegative atom
63 HFNo net overlap between H1s and (F2px or F2py) ⇒ 2 "nonbonding" orbitals
64 orbital mainly of F2pz (energy level close to F2pz) orbital mainly of H1s (energy level close to H1s)⇒
65 CO and NOCONOmost stablediatomic moleculefrom 2p-2p mixing only
66 3.12 Orbitals in Polyatomic Molecules 121- The MOs spread over all atoms in the molecule.experimentally studied by using ultraviolet and visible spectroscopytoo complex --- qualitative assessment- A water molecule with six atomic orbitals(one O2s, three O2p, and two H1s)1b1; nonbonding, mainly O2py, lone pair effect2a1; almost nonbonding⇒
68 MO and energy levels of a linear triatomic dihydride HXH LUMOReversed the energy levels σ2s* -σ2p (cf. Fig. 7.7)energy – no. of nodes relationship
69 • 8 valence electrons of water → (σ2s)2(σ2p)2(nπ2p)4 • By bending ; σ2s→ favorable since of constructive overlapbetween the end hydrogens.σ2p→ destabilizes since 1s orbitals have opposite signsn2p→ depends on their orientatione.g. bending occurs in the xz plane py – little change;px – can overlap with H 1s orbital: lowing its energy 2 orbitals go down in energy(σ2s, n2px),1 goes up (σ2p), and 1 remains same (n2py)→ water is favorable to be bent geometry
70 MO of water121sLUMOHOMOEnergy decrease;Energy increase
71 CH4 ; 1 of the 4 electron pairs is slightly lower in energy. from photoelectron spectroscopyVB-theory; all eight electrons have the same energy.MO-theory; (1a1)2(1t1)6⇒ Lower energy for the 1a1 electron pair
72 122benzene, C6H6- All thirty C2s-, C2p-, and H1s-orbitals contribute MOs.- The orbitals in the ring plane:C2s-, C2px, C2py, and six H1s-orbitals → delocalized s-orbitals for C-C and C-H- six C2pz-orbitals perpendicular to the ring → delocalized p-orbitals spreading the ring- consider the two separately !!From VB, each C atom with sp2 hybrid orbitals forming s-bonds and 120° angles.From MO, the six C2pz-orbitals form six delocalized p-orbitals.
73 6 p-orbital shapes from fully bonding to fully antibonding great stability: the p-electrons occupyonly orbitals with a net bonding effect
74 MO does not require electron pair for each bonding or octet rule for a particular atom as all electrons are spread over all the atoms in the molecule.123Hypervalent compounds- From VB, SF6 with sp3d2 hybridization- From MO, four orbitals of S and six of F,a total of 10 AOs → 10 MOs12 electrons occupy bonding and nonbondingorbitals.- Average bond order of each S-F is 2/3.
75 Colors of vegetation 124 for benzene p-electrons Unoccupied MO hn LUMO excitationOccupiedMOHOMO- lowest unoccupied molecular orbital (LUMO)- highest occupied molecular orbital (HOMO)
76 Colors of vegetation 124 Particle in a box (one dimensional) b-caroteneParticle in a box (one dimensional)lycopeneRetinal (vitamin A)
77 ULTRAVIOLET AND VISIBLE SPECTROSCOPY 130ULTRAVIOLET AND VISIBLE SPECTROSCOPYThe Technique- The electrons in the molecule can be excited to a higher energy state,by electromagnetic radiation.Bohr frequency condition, DE = hn- UV-vis absorption gives us information about the electronic energylevels of molecules.i.e. Chlorophyll absorbs red and bluelight, leaving the green light presentin white light to be reflected.
78 131Chromophores- Characteristic groups of atoms in the molecules absorbing certainbands in visible and ultraviolet spectra- p-to-p* transition in conjugated double bonds ~ 160 nm- n-to-p* transition in the carbonyl group ~ 280 nm- d-to-d transition in d-metal complexes in visible ranges- charge transfer transition in d-metal complexes electrons migratefrom the ligands to the metal atom or vice versai.e. deep purple color of MnO4-
79 The nuclear model of atom Thompson’s model Rutherford’s model Model of the atomThe nuclear model of atomThompson’s modelRutherford’s modelAtomic spectra -- Bohr’s modelQuantum theoryQuantization: M. PlankWave-particle duality: de BroglieUncertainty principle: HeisenbergWave function – Schroedinger equationE = hnparticle in a boxSolution:
80 Hydrogen atom (one electron atoms) particle in a boxSolution:Hydrogen atom (one electron atoms)n= principal quantum number
81 Toward molecules…... Hydrogen atom (one electron atoms) Principal quantum numberAtomic orbitals: orbital angular momentum --- shapemagnetic quantum number -- orientationspin magnetic quantum number – spin directionPeriodicity of atomic propertiesMany-electron atomsorbital energyshielding effecteffective nuclear chargePauli exclusion principlevalence shell (electrons)Hund’s ruleAtomic radiusIonic radiusIonization energyElectron affinityPeriodic tableToward molecules…...Chemical bond is the link between atoms.Ionic bond; electron transfer + electrostatic attractionCovalent bond; sharing electronsLewis structureOctet rule coordinate covalent bondResonanceFormal chargeOxidation number
82 toward molecular structure and properties. 115toward molecular structure and properties.The valence-bond (VB) and molecular orbital (MO) theories are both procedures for constructing approximate wavefunctions of electrons.- In VB theory, bonding electrons are localized on atoms or between pairs of atoms.Molecular orbitals (MOs)The MO theory can account for electron-deficient compounds, paramagnetic O2, and many other properties by focusing on electrons delocalized over the whole molecule.MOs formed by linear combination of atomic orbitals (LCAO-MO)Approximate molecular wavefunctions by superimposing (mixing) of N atomic orbitalscij and Ej are determinedby solving the Schrödinger equation
83 116Building-up principle for MOValence electrons in molecular orbitals1. the lowest-energy MO first then in orbitals of increasingly higher energy.2. Pauli exclusion principle:3. Hund’s rule:1. No mixing between AO's of the same atom2. Significant mixing only between AO's of similar energies and substantial overlap
84 3.11 Bonding in Heteronuclear Diatomic Molecules 1203.11 Bonding in Heteronuclear Diatomic MoleculesCO and NO