Presentation on theme: "2012 General Chemistry I What we have learned so far toward molecular structure and properties. Quantum theory Principle quantum number Atomic orbitals:"— Presentation transcript:
2012 General Chemistry I What we have learned so far toward molecular structure and properties. Quantum theory Principle quantum number Atomic orbitals: radial wave function angular wave function orbital angular momentum magnetic quantum numbers radial distribution function shape of atomic orbitals electron spin orbital energy split shielding effect effective nuclear charge Pauli exclusion principle Hund’s rule valence shell (electrons) Model of the atom The nuclear model of atom Thompson’s model Rutherford’s model Atomic spcetra -- Bohr’s model Hydrogen atom (one electron atoms) Wave-particle duality Uncertainty principle Wave function – particle in a box Schroedinger equation Many electron atoms Periodicity of atomic properties Atomic radius Ionic radius Ionization energy Electron affinity Periodic table
2012 General Chemistry I What we have learned so far toward molecular structure and properties. Ionic bond Lewis structure Octet rule Exceptions to octet rule Resonance Formal charge Oxidation number Ionic v.s. covalent Dipole moment Polar bond Nonpolar bond Bond strength Bond length Chemical bond Interaction between two electrons Covalent bond electronegativity MOLECULAR SHAPE AND STRUCTURE Stability, reactivity, color, size, polarity, solubility, function etc… IR(infrared) spectroscopy
2012 General Chemistry I 3D structure of a molecule is crucial for its property. Sophisticated quantum mechanical calculations are needed to predict the structure. ⇒ Drugs by Design and Discovery Box 3.1 1) Identification of key enzymes 2) Molecular structure determination 4) Computer-aided design of molecules with structures fitting into the active site 3) Hints from Nature --- Natural Products
2012 General Chemistry I Chapter 3. MOLECULAR SHAPE AND STRUCTURE 2012 General Chemistry I THE VSEPR MODEL ( 전자쌍 반발 모델 ) VALENCE-BOND THEORY ( 원자가 결합 이론 ) 3.1 The Basic VSEPR Model 3.2 Molecules with Lone Pairs on the Central Atom 3.3 Polar Molecules 3.4 Sigma and Pi Bonds 3.5 Electron Promotion and the Hybridization of Orbitals ( 혼성궤도 함수 ) 3.6 Other Common Types of Hybridization 3.7 Characteristics of Multiple Bonds
2012 General Chemistry I THE VSEPR MODEL (Sections 3.1-3.3) 95 Lewis structure: showing the linkages between atoms and the presence of lone pairs, but not the 3D arrangement of atoms BF 3 BeCl 2 H2OH2O NH 3 CH 4 ClF 3 SF 4 XeF 4 PCl 5 IF 5 SF 6
2012 General Chemistry I Estimating the 3D structure: THE VSEPR MODEL 3.1 The Basic VSEPR Model Valence Shell Electron-Pair Repulsion theory Electron pairs (lone pairs & bonding pairs) repel each other. Rule 1: Electron pairs move as far apart as possible. Proposed by R. J. Gillespie in 1959. VSEPR structures for AX n with no lone pair
2012 General Chemistry I Rule 2: (Almost) No distinction between single and multiple bonds. BF 3 BeCl 2 CO 2 CO 3 2-
2012 General Chemistry I 98 Rule 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. 3.2 Molecules with Lone Pairs on the Central Atom NH 3 CH 4
2012 General Chemistry I 99 Rule 4 The strength of repulsions are in the order lone pair-lone pair > lone pair-atom > atom-atom NH 3 H2OH2O
2012 General Chemistry I 99 axial equatorial more stable seesaw shaped T-shaped ClF 3 PCl 5 SF 4 axialequatorial
2012 General Chemistry I 101 Predicting a molecular shape of XeF 4 Step 1 Draw the Lewis structure. Step 2 Assign the electron arrangement around the central atom. Step 3 Identify the molecular shape. AX 4 E. Step 4 Allow for distortions. Square planar
2012 General Chemistry I AXE method A; central atom X; outside atom E; lone pair
2012 General Chemistry I 3.3 Polar Molecules 102 Polar molecule: a molecule with a nonzero dipole moment i.e. HCl with a dipole moment of 1.1 D polar - A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that the dipole moments associated with the bonds do not cancel. HCl, H 2 O, CHCl 3, cis-dichloroethane, ···
2012 General Chemistry I 3.3 Polar Molecules 102 Nonpolar molecule: a molecule with a net zero dipole moment nonpolar Homonuclear diatomic molecules Polyatomic molecules with symmetry; CO 2, BF 3, CH 4, CCl 4, trans-dichloroethane, ···
2012 General Chemistry I VALENCE-BOND THEORY (Sections 3.4-3.7) 105 Lewis model of the chemical bond; localized electron model Valence-bond theory; Walter Heitler, Fritz London (1927) Linus Pauling (1931) Quantum mechanical description of the distribution of electrons in bonds Valence electrons are localized either between pairs of atoms or on atoms as lone pairs. 1)Hybridization of atomic valence orbitals with proper symmetry that 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 simple molecules. --- does not explain many things including multiple bond, bond angles ……..
2012 General Chemistry I 3.4 Sigma) and Pi) Bonds: description of covalent bond 105 H 2 1) Two hydrogen 1s-orbitals merge (overlap) to form a -orbital between the two hydrogen atoms. 2) A -bond is formed as two electrons (↑↓) fill the -orbital. -bond ; cylindrically symmetrical with no nodal planes containing the intermolecular axis. i.e. H 2, 1s-1s HF, 1s-2p z N 2, 2p z -2p z Walter Heitler, Fritz London (1927)
2012 General Chemistry I 106 overlap with 1:1 match of orbitals
2012 General Chemistry I 106 overlap with 1:1 match of orbitals
106 -bond nodal plane containing the interatomic (bond) axis - two cylindrical shapes (lobes), one above and the other below the nodal plane one -bond with two perpendicular -bonds N2N2 - multiple bonds: single bond (one -bond), double bond (one - and one -bond) triple bond (one - and two -bonds)
2012 General Chemistry I Polyatomic molecules 107 Linus Pauling (1931) BeH 2 3.5 Electron Promotion and the Hybridization of Orbitals Why do we have to make hybrid orbitals?
2012 General Chemistry I not real. localization problem Linear combination of orbitals sp hybrid
3.7 Characteristics of Multiple Bonds 111 CO 2 Carbon Oxygen Steric # = 2 Steric # = 3
2012 General Chemistry I 3.7 Characteristics of Multiple Bonds 111 CO 2
2012 General Chemistry I 3.7 Characteristics of Multiple Bonds 111 - ethene, CH 2 =CH 2 C-C bond, (C2sp 2, C2sp 2 ) C-C bond, (C2p, C2p) each C-H bond formed as (C2sp 2, H1s) restricted rotation
2012 General Chemistry I - ethyne (acetylene), C 2 H 2 112 free rotation
2012 General Chemistry I 112 - benzene, C 6 H 6 Now, delocalization has the meaning ! Still, there are many properties that can not be explained by the current model.
2012 General Chemistry I 111 CO 3 2- resonance hybrid
2012 General Chemistry I 113 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory Paramagnetic: 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. Paramagnetic O 2 ; unpaired electron(s) Lewis's theory; Valence-bond theory; bond and bond 2 lone pairs on each O occupying the sp 2 hybrid orbitals
2012 General Chemistry I Shortcomings of the Valence Bond Model 1)Inadequate treatment of odd-electron molecules and resonances 2) Magnetism of molecules Paramagnetic: molecules with unpaired electrons Diamagnetic: weakly repelled by a magnetic field O2O2 N2N2 both are expected to be diamagnetic!! 113s 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory
2012 General Chemistry I 113 Electron deficient diborane, B 2 H 6 does not have enough electrons ! At least seven bonds (= 14 electrons) are required, but only 12 valence electrons. First published by H. C. Lunguet-Higgins, a 2 nd year undergraduate student ! - No simple explanation for spectroscopic properties of compounds 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory
2012 General Chemistry I MOLECULAR ORBITAL THEORY (Sections 3.8-3.12) 113 Molecular 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 3.9 Molecular Orbitals 3.10 Electron Configurations of Diatomic Molecules 3.11 Bonding in Heteronuclear Diatomic Molecules 3.12 Orbitals in Polyatomic Molecules
2012 General Chemistry I H 2 + : Prototype Molecular Orbital System Atomic orbital(AO) theory → successful for orbital structures of all atoms with both even and odd numbers of electrons Assume that molecule H 2 + ~ as an united atom with a fragmented nucleus if the nuclei in molecule were fused together construct the one-electron orbital corresponding to the arrangement of nuclear charges presented by the molecule Coulomb interactions in H 2 + 115s
2012 General Chemistry I 3.9 Molecular Orbitals 115 Even for the smallest molecule, H 2 Quantum mechanics : the ideal solution to the problem, but……. Schroedinger 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)
2012 General Chemistry I 3.9 Molecular Orbitals 115 Molecular orbitals (MOs) - In VB theory, bonding electrons are localized on atoms or between pairs of atoms. The valence-bond (VB) and molecular orbital (MO) theories are both procedures for constructing approximate wavefunctions of electrons. The MO theory can account for electron-deficient compounds, paramagnetic O 2, 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 orbitals c ij and E j are determined by solving the Schrödinger equation
2012 General Chemistry I Trial wavefunctions for H 2 using two 1s atomic orbitals of H Increased amplitude in the internuclear region bonding Larger volume for electrons lower kinetic energy (particle-in-a-box) Decreased amplitude in the internuclear region & nodal plane antibonding
2012 General Chemistry I 115 Molecular orbital energy-level diagram - relative energies of original AOs and resulting MOs - arrows to show electron spin and location of the electrons - In H 2, two 1s-orbitals merge to form the bonding orbital 1s and the antibonding orbital 1s *
2012 General Chemistry I 3.10 Electron Configurations of Diatomic Molecules 116 Valence electrons in molecular orbitals 1. Electrons are accommodated in the lowest-energy MO, then in orbitals 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. Building-up principle for MO H 2 : The energy of H 2 is lower than that of the separate H atoms. Even the energy of H 2 + is lower than that of the separate H atoms.
2012 General Chemistry I H2H2 He 2 Bond order = ½( # of bonding electrons - # of antibonding electrons) Bond order = 1 Bond order = 0 H2+H2+ He 2 + Bond order = 1/2
2012 General Chemistry I For other homonuclear diatomic molecules of Period 2 elements, - two 2s orbitals (one on each atom) overlap to form two s orbitals, one bonding (s 2s -orbital) and the other antibonding (s 2s *-orbital) - six 2p orbitals (three on each atom) overlap to form six MOs, two 2p z orbitals to form bonding and antibonding (s 2p, s 2p *) four 2p x, 2p y orbitals to form two p 2p and two p 2p * orbitals Linear combination of 10 atomic orbitals; 1. No mixing between AO's of the same atom 2. 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
116 two 2p z orbitals to form bonding 2p and antibonding 2p * orbitals four 2p x, 2p y orbitals to form two 2p and two 2p * orbitals one bonding ( 2s -orbital) and the other antibonding ( 2s *-orbital)
2012 General Chemistry I - From Li 2 to N 2, the energy levels of 2s and 2p are close, and thus the 2s orbital also participates in forming s 2p orbitals.
2012 General Chemistry I - For O 2 and F 2, the energy levels of 2s and 2p are separated well.
2012 General Chemistry I - In N 2, each atom supplies five valence electrons. A total of ten electrons fill the MOs. The ground configuration is, Bond order (b): net number of bonds b = ½(8-2) = 3 - In O 2, each atom supplies six valence electrons. A total of twelve electrons fill the MOs. The ground configuration is, b = ½(8-4) = 2 accounts for paramagnetism of O 2 118
2012 General Chemistry I 118 paramagnetic Bond order (b) =10 1 2321 does not exist
2012 General Chemistry I Second-Row MO Diagrams
2012 General Chemistry I 3.11 Bonding in Heteronuclear Diatomic Molecules A diatomic molecule built from atoms of two different elements in polar, with the electrons shared unequally by the two atoms. - In a nonpolar covalent bond, c A 2 = c B 2 - 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. 120
2012 General Chemistry I more electronegative atom less electronegative atom
2012 General Chemistry I HF No net overlap between H1s and (F2p x or F2p y ) ⇒ 2 "nonbonding" orbitals
2012 General Chemistry I orbital mainly of F2p z (energy level close to F2p z ) orbital mainly of H1s (energy level close to H1s) ⇒
2012 General Chemistry I CO and NO from 2p-2p mixing only CONO most stable diatomic molecule
2012 General Chemistry I 3.12 Orbitals in Polyatomic Molecules - The MOs spread over all atoms in the molecule. experimentally studied by using ultraviolet and visible spectroscopy too complex --- qualitative assessment - A water molecule with six atomic orbitals (one O2s, three O2p, and two H1s) 121 1b 1 ; nonbonding, mainly O2p y, lone pair effect 2a 1 ; almost nonbonding ⇒
2012 General Chemistry I H1s-O2p y -H1s O2p x nonbonding orbital antibonding orbitals bonding orbitals H1s-(O2s,2p z )-H1s 121
2012 General Chemistry I MO and energy levels of a linear triatomic dihydride HXH Reversed the energy levels σ 2s * -σ 2p (cf. Fig. 7.7) energy – no. of nodes relationship LUMO 121s
2012 General Chemistry I 8 valence electrons of water → (σ 2s ) 2 (σ 2p ) 2 (nπ 2p ) 4 By bending ; σ 2s → favorable since of constructive overlap between the end hydrogens. σ 2p → destabilizes since 1s orbitals have opposite signs n 2p → depends on their orientation e.g. bending occurs in the xz plane p y – little change; p x – can overlap with H 1s orbital: lowing its energy 2 orbitals go down in energy(σ 2s, n 2px ), 1 goes up (σ 2p ), and 1 remains same (n 2py ) → water is favorable to be bent geometry 121s
2012 General Chemistry I MO of water Energy decrease;Energy increase LUMO HOMO 121s
2012 General Chemistry I CH 4 ; 1 of the 4 electron pairs is slightly lower in energy. from photoelectron spectroscopy VB-theory; all eight electrons have the same energy. MO-theory; (1a 1 ) 2 (1t 1 ) 6 ⇒ Lower energy for the 1a 1 electron pair
2012 General Chemistry I benzene, C 6 H 6 - All thirty C2s-, C2p-, and H1s-orbitals contribute MOs. - The orbitals in the ring plane: C2s-, C2p x, C2p y, and six H1s-orbitals → delocalized -orbitals for C-C and C-H - six C2p z -orbitals perpendicular to the ring → delocalized -orbitals spreading the ring From VB, each C atom with sp 2 hybrid orbitals forming -bonds and 120° angles. From MO, the six C2p z -orbitals form six delocalized -orbitals. 122 - consider the two separately !!
2012 General Chemistry I 6 -orbital shapes from fully bonding to fully antibonding bonding anti- bonding great stability: the -electrons occupy only orbitals with a net bonding effect
2012 General Chemistry I Hypervalent compounds - From VB, SF 6 with sp 3 d 2 hybridization - From MO, four orbitals of S and six of F, a total of 10 AOs → 10 MOs 12 electrons occupy bonding and nonbonding orbitals. - Average bond order of each S-F is 2/3. 123 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.
2012 General Chemistry I Colors of vegetation - highest occupied molecular orbital (HOMO) 124 for benzene -electrons Unoccupied MO Occupied MO h HOMO - lowest unoccupied molecular orbital (LUMO) LUMO excitation
2012 General Chemistry I Colors of vegetation 124 -carotene lycopene Retinal (vitamin A) Particle in a box (one dimensional)
2012 General Chemistry I ULTRAVIOLET AND VISIBLE SPECTROSCOPY 130 The Technique - The electrons in the molecule can be excited to a higher energy state, by electromagnetic radiation. Bohr frequency condition, E = h - UV-vis absorption gives us information about the electronic energy levels of molecules. i.e. Chlorophyll absorbs red and blue light, leaving the green light present in white light to be reflected.
2012 General Chemistry I 131 Chromophores - Characteristic groups of atoms in the molecules absorbing certain bands in visible and ultraviolet spectra - -to- * transition in conjugated double bonds ~ 160 nm - n-to- * 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 migrate from the ligands to the metal atom or vice versa i.e. deep purple color of MnO 4 -
2012 General Chemistry I Quantum theory Model of the atom The nuclear model of atom Thompson’s model Rutherford’s model Atomic spectra -- Bohr’s model Quantization: M. Plank Wave-particle duality: de Broglie Uncertainty principle: Heisenberg Wave function – Schroedinger equation E = h particle in a box Solution:
2012 General Chemistry I particle in a box Solution: Hydrogen atom (one electron atoms) n= principal quantum number
2012 General Chemistry I Principal quantum number Atomic orbitals: orbital angular momentum --- shape magnetic quantum number -- orientation spin magnetic quantum number – spin direction Hydrogen atom (one electron atoms) Many-electron atoms orbital energy shielding effect effective nuclear charge Pauli exclusion principle valence shell (electrons) Hund’s rule Periodicity of atomic properties Atomic radius Ionic radius Ionization energy Electron affinity Periodic table Chemical bond is the link between atoms. Toward molecules…... Ionic bond; electron transfer + electrostatic attraction Covalent bond; sharing electrons Lewis structure Octet rule ---- ---- coordinate covalent bond Resonance Formal charge Oxidation number
2012 General Chemistry I 115 Molecular orbitals (MOs) - In VB theory, bonding electrons are localized on atoms or between pairs of atoms. The valence-bond (VB) and molecular orbital (MO) theories are both procedures for constructing approximate wavefunctions of electrons. The MO theory can account for electron-deficient compounds, paramagnetic O 2, 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 orbitals c ij and E j are determined by solving the Schrödinger equation toward molecular structure and properties.
2012 General Chemistry I 116 Valence electrons in molecular orbitals 1. the lowest-energy MO first then in orbitals of increasingly higher energy. 2. Pauli exclusion principle: 3. Hund’s rule: Building-up principle for MO 1. No mixing between AO's of the same atom 2. Significant mixing only between AO's of similar energies and substantial overlap
2012 General Chemistry I 3.11 Bonding in Heteronuclear Diatomic Molecules 120 CO and NO