What we have learned so far toward molecular structure and properties.

Name: What we have learned so far toward molecular structure and properties.
Uploaded: 2017-07-12T14:07:01+00:00
Duration: PTM36S5
Channel: Theresa Sword
Description: What we have learned so far toward molecular structure and properties.

What we have learned so far toward molecular structure and properties.
Model of the atom The nuclear model of atom Thompson’s model Rutherford’s model Atomic spcetra -- Bohr’s model Quantum theory Periodicity of atomic properties Wave-particle duality Uncertainty principle Wave function – particle in a box Schroedinger equation Atomic radius Ionic radius Ionization energy Electron affinity Periodic table Hydrogen atom (one electron atoms) 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 Many electron atoms orbital energy split shielding effect effective nuclear charge Pauli exclusion principle Hund’s rule valence shell (electrons)

MOLECULAR SHAPE AND STRUCTURE
What we have learned so far toward molecular structure and properties. Chemical bond Interaction between two electrons Ionic bond Lewis structure Octet rule Exceptions to octet rule Resonance Formal charge Oxidation number Covalent bond electronegativity Ionic v.s. covalent Dipole moment Polar bond Nonpolar bond Bond strength Bond length IR(infrared) spectroscopy MOLECULAR SHAPE AND STRUCTURE Stability, reactivity, color, size, polarity, solubility, function etc…

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 3) Hints from Nature --- Natural Products 4) Computer-aided design of molecules with structures fitting into the active site

MOLECULAR SHAPE AND STRUCTURE
Chapter 3. MOLECULAR SHAPE AND STRUCTURE THE VSEPR MODEL (전자쌍 반발 모델) 3.1 The Basic VSEPR Model 3.2 Molecules with Lone Pairs on the Central Atom 3.3 Polar Molecules VALENCE-BOND THEORY (원자가 결합 이론) 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 ClF3 CH4 H2O BF3 NH3 BeCl2 SF4 XeF4 PCl5 IF5 SF6

Electron pairs (lone pairs & bonding pairs) repel each other.
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. 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

BF3 CH4 BeCl2

PCl5 SF6

Rule 2: (Almost) No distinction between single and multiple bonds.
BeCl2 CO2 CO32- BF3

3.2 Molecules with Lone Pairs on the Central Atom
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. NH3 CH4

Rule 4 The strength of repulsions are in the order
99 Rule 4 The strength of repulsions are in the order lone pair-lone pair > lone pair-atom > atom-atom H2O NH3

ClF3 SF4 PCl5 99 axial equatorial T-shaped seesaw shaped more stable

Predicting a molecular shape of XeF4
101 Predicting a molecular shape of XeF4 Step 1 Draw the Lewis structure. Step 2 Assign the electron arrangement around the central atom. Step 3 Identify the molecular shape. AX4E. Step 4 Allow for distortions. Square planar

AXE method A; central atom X; outside atom E; lone pair

102 3.3 Polar Molecules Polar molecule: a molecule with a nonzero dipole moment i.e. HCl with a dipole moment of 1.1 D HCl, H2O, CHCl3, cis-dichloroethane, ··· - 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. polar polar

Homonuclear diatomic molecules
102 3.3 Polar Molecules Nonpolar molecule: a molecule with a net zero dipole moment Homonuclear diatomic molecules Polyatomic molecules with symmetry; CO2, BF3, CH4, CCl4, trans-dichloroethane, ··· nonpolar nonpolar

VALENCE-BOND THEORY (Sections 3.4-3.7)
105 VALENCE-BOND THEORY (Sections ) 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. 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 ……..

3.4 s(Sigma) and p(Pi) Bonds: description of covalent bond
105 3.4 s(Sigma) and p(Pi) Bonds: description of covalent bond H2 1) 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

overlap with 1:1 match of orbitals
106 overlap with 1:1 match of orbitals

nodal plane containing the interatomic (bond) axis
106 p-bond nodal plane containing the interatomic (bond) axis - two cylindrical shapes (lobes), one above and the other below the nodal plane N2 one s-bond with two perpendicular p-bonds - multiple bonds: single bond (one s-bond), double bond (one s- and one p-bond) triple bond (one s- and two p-bonds)

3.5 Electron Promotion and the Hybridization of Orbitals
107 Polyatomic molecules Linus Pauling (1931) BeH2 3.5 Electron Promotion and the Hybridization of Orbitals Why do we have to make hybrid orbitals?

Linear combination of orbitals
not real. localization problem Linear combination of orbitals sp hybrid

Linear molecule

107 BH3

sp2 hybrid orbitals

107 CH4 promotion hybridization

sp3 hybrids 107 h1 = s + px + py + pz h2 = s - px - py + pz
similar ideas to VSEPR

NH3 H2O 107 hybrid orbitals can be determined by the steric number
based on the VSEPR model. steric number = # of atoms bonded to the central atom + # of lone pairs H2O

109 - sp3d hybrid orbitals in PCl5 - sp3d2 hybrid orbitals in SF6 and XeF4

3.7 Characteristics of Multiple Bonds
111 CO2 Carbon Steric # = 2 Oxygen Steric # = 3

111 CO2

111 - ethene, CH2=CH2 C-C s bond, s(C2sp2, C2sp2) C-C p bond, p(C2p, C2p) each C-H bond formed as s(C2sp2, H1s) restricted rotation

- ethyne (acetylene), C2H2
112 - ethyne (acetylene), C2H2 free rotation

Now, delocalization has the meaning !
112 - benzene, C6H6 Now, delocalization has the meaning ! Still, there are many properties that can not be explained by the current model.

111 CO32- resonance hybrid

3.8 The Limitations of Lewis’s Theory
& Valence Bond Theory 113 Paramagnetic O2; unpaired electron(s) Lewis's theory; Valence-bond theory; bond and bond 2 lone pairs on each O occupying the sp2 hybrid orbitals 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.

Shortcomings of the Valence Bond Model
3.8 The Limitations of Lewis’s Theory & Valence Bond Theory Shortcomings of the Valence Bond Model Inadequate treatment of odd-electron molecules and resonances O2 N2 2) Magnetism of molecules • Paramagnetic: molecules with unpaired electrons • Diamagnetic: weakly repelled by a magnetic field both are expected to be diamagnetic!!

3.8 The Limitations of Lewis’s Theory
& Valence Bond Theory 113 Electron deficient diborane, B2H6 First 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

MOLECULAR ORBITAL THEORY (Sections 3.8-3.12)
113 MOLECULAR ORBITAL THEORY (Sections ) 3.9 Molecular Orbitals 3.10 Electron Configurations of Diatomic Molecules 3.11 Bonding in Heteronuclear Diatomic Molecules 3.12 Orbitals in Polyatomic Molecules 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

H2+: 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 H2+ ~ 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 H2+

3.9 Molecular Orbitals 115 Quantum mechanics : the ideal solution to the problem, but……. Even for the smallest molecule, H2 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)

3.9 Molecular Orbitals 115 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 orbitals cij and Ej are determined by solving the Schrödinger equation

Trial wavefunctions for H2 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

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 H2, two 1s-orbitals merge to form the bonding orbital s1s and the antibonding orbital s1s*

3.10 Electron Configurations of Diatomic Molecules
116 Building-up principle for MO 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. H2: The energy of H2 is lower than that of the separate H atoms. Even the energy of H2+ is lower than that of the separate H atoms.

H2 He2 He2+ H2+ Bond order = 0 Bond order = 1
Bond order = ½( # of bonding electrons - # of antibonding electrons) He2+ H2+ Bond order = 1/2 Bond order = 1/2

For other homonuclear diatomic molecules of Period 2 elements,
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 - 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

antibonding s2p* orbitals
116 four 2px, 2py orbitals to form two p2p and two p2p* orbitals two 2pz orbitals to form bonding s2p and antibonding s2p* orbitals one bonding (s2s-orbital) and the other antibonding (s2s*-orbital)

- From Li2 to N2, the energy levels of 2s and 2p are close, and thus
the 2s orbital also participates in forming s2p orbitals.

- For O2 and F2, the energy levels of 2s and 2p are separated well.

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 bonds b = 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) = 2 accounts for paramagnetism of O2

118 Bond order (b) = 1 1 2 3 2 1 paramagnetic does not exist

Second-Row MO Diagrams

3.11 Bonding in Heteronuclear Diatomic Molecules
120 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, 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.

less electronegative atom
Fig 3.33 more electronegative atom

HF No net overlap between H1s and (F2px or F2py) ⇒ 2 "nonbonding" orbitals

orbital mainly of F2pz (energy level close to F2pz)
orbital mainly of H1s (energy level close to H1s) ⇒

CO and NO CO NO most stable diatomic molecule from 2p-2p mixing only

3.12 Orbitals in Polyatomic Molecules
121 - 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) 1b1; nonbonding, mainly O2py, lone pair effect 2a1; almost nonbonding ⇒

121 antibonding orbitals nonbonding orbital O2px H1s-O2py-H1s bonding orbitals H1s-(O2s,2pz)-H1s

MO and energy levels of a linear triatomic dihydride HXH
LUMO Reversed the energy levels σ2s* -σ2p (cf. Fig. 7.7) energy – no. of nodes relationship

• 8 valence electrons of water → (σ2s)2(σ2p)2(nπ2p)4 • By bending ;
σ2s→ favorable since of constructive overlap between the end hydrogens. σ2p→ destabilizes since 1s orbitals have opposite signs n2p→ depends on their orientation e.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

MO of water 121s LUMO HOMO Energy decrease; Energy increase

CH4 ; 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; (1a1)2(1t1)6 ⇒ Lower energy for the 1a1 electron pair

122 benzene, 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.

6 p-orbital shapes from fully bonding to fully antibonding
great stability: the p-electrons occupy only orbitals with a net bonding effect

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. 123 Hypervalent compounds - From VB, SF6 with sp3d2 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.

Colors of vegetation 124 for benzene p-electrons Unoccupied MO hn LUMO
excitation Occupied MO HOMO - lowest unoccupied molecular orbital (LUMO) - highest occupied molecular orbital (HOMO)

Colors of vegetation 124 Particle in a box (one dimensional)
b-carotene Particle in a box (one dimensional) lycopene Retinal (vitamin A)

ULTRAVIOLET AND VISIBLE SPECTROSCOPY
130 ULTRAVIOLET AND VISIBLE SPECTROSCOPY The 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 energy levels of molecules. i.e. Chlorophyll absorbs red and blue light, leaving the green light present in white light to be reflected.

131 Chromophores - Characteristic groups of atoms in the molecules absorbing certain bands 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 migrate from the ligands to the metal atom or vice versa i.e. deep purple color of MnO4-

The nuclear model of atom Thompson’s model Rutherford’s model
Model of the atom The nuclear model of atom Thompson’s model Rutherford’s model Atomic spectra -- Bohr’s model Quantum theory Quantization: M. Plank Wave-particle duality: de Broglie Uncertainty principle: Heisenberg Wave function – Schroedinger equation E = hn particle in a box Solution:

Hydrogen atom (one electron atoms)
particle in a box Solution: Hydrogen atom (one electron atoms) n= principal quantum number

Toward molecules…... Hydrogen atom (one electron atoms)
Principal quantum number Atomic orbitals: orbital angular momentum --- shape magnetic quantum number -- orientation spin magnetic quantum number – spin direction Periodicity of atomic properties Many-electron atoms orbital energy shielding effect effective nuclear charge Pauli exclusion principle valence shell (electrons) Hund’s rule Atomic radius Ionic radius Ionization energy Electron affinity Periodic table Toward molecules…... Chemical bond is the link between atoms. Ionic bond; electron transfer + electrostatic attraction Covalent bond; sharing electrons Lewis structure Octet rule coordinate covalent bond Resonance Formal charge Oxidation number

toward molecular structure and properties.
115 toward 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 orbitals cij and Ej are determined by solving the Schrödinger equation

116 Building-up principle for MO 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: 1. No mixing between AO's of the same atom 2. Significant mixing only between AO's of similar energies and substantial overlap

3.11 Bonding in Heteronuclear Diatomic Molecules
120 3.11 Bonding in Heteronuclear Diatomic Molecules CO and NO

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