Topic 2 2-1 Structure and Bonding Models of Covalent Compounds of p-block Elements Bonding 2-2 Many different approaches to describe bonding: Ionic Bonding: Elements with large electronegativity differences; bonding is due to electrostatic interaction between ions (formed by the transfer of electrons between the two bonded atoms) Covalent Bonding: Electronegativity difference smaller; covalent interaction through sharing of electrons between bonded atoms)
Bond Distance and Atom Radius 2-3 The length of a covalent bond is given by the internuclear distance > determined by microwave spectroscopy or diffraction methods (X-ray, neutron or electron diffraction) Covalent radius of an atom X = half of the bond length of a homonuclear single bond X X Note: The van der Waals radius of an atom X is half of the distance of closest approach of two non-bonded atoms (larger than the covalent radius) Covalent Bonding 2-4 Modern models describe chemical bond based on quantum mechanical methods using molecular orbitals Earlier models and concepts are much simpler, but can still be very useful for qualitative description of molecule structures and geometries: Lewis theory Resonance structures and formal charges Valence Shell Electron Pair Repulsion Theory (VSEPR)
Lewis Theory: The Octet Rule 2-5 Lewis theory of bonding (1916) was one of the earliest models to have success Basic concept: OCTET RULE: Main group elements are surrounded by 8 electrons when forming covalent compounds Note: works well with second period elements, but runs into problems with all others Bond Order 2-6 Often, Lewis structures can be used to estimate bond orders that correlate well with experimentally measured bond strengths and lengths: Single bond: bond order = O O Double bond: bond order = N N Triple bond: bond order =
ormal Charges 2-7 ormal charges are calculated by dividing shared electrons equally between the bonded atom pair ormal charge = number of electrons in valence shell number of assigned electrons Note: A molecule may have more than one plausible Lewis structure > The best Lewis structure is the one that has the least charge separation and puts the negative charge on the most electronegative elements N N O N N O N N O Oxidation Numbers 2-8 Oxidation numbers are calculated by assigning shared electrons to the more electronegative atom Oxidation number = number of electrons in valence shell number of assigned electrons General rules for assigning oxidation numbers in molecules: Group 1 metals: Group 2 metals: Al is only Al 3+ is always O is always O 2
Resonance Structures 2-9 Some molecules have more than one distinct Lewis structures: Acetic acid: Note: The real structure is an average of the two drawn structures = resonance hybrid > the bond order of the the C O bonds is therefore ailures of the Lewis Model 2-10 Molecules with odd numbers of electrons exist than 8 electrons in the valence shell of their atoms exist: NO BCl 3 > both are stable molecules A central atom may have more than 8 electrons (for n > 3 = third period and higher period main group elements) S 6 O 2
VSEPR Theory 2-11 Valence Shell Electron Pair Repulsion Theory is useful for predicting the geometry of main group compounds Theory is based on idea, that molecules adopt the geometry for which the repulsion between electron pairs (bonding or non-bonding) are as small as possible Geometries with Minimum Repulsion 2-12
Rules 2-13 1) Count all single bond electron pairs (BP) 2) Count all non bonding electron pairs = lone pairs (LP) 3) Calculate sum of BP + LP 4) Determine geometry according to previous rules Note: our double and triple bonds only ONE BP is counted, since the second (or third) does not require much more space Linear Geometry 2-14 180 180 Cl Be Cl O C O
Trigonal Geometry 2-15 B O N O O N O 120 134 115 Note: Lone pair require more space than bonding pairs One non-bonding electron requires less space than a bonding pair Tetrahedral Geometry 2-16 H H C H H 109.5 H N H H O H H 107 104 The arrangement of LP and BP is tetrahedral, but the molecular shape is named after the arrangements of the ATOMS!
Trigonal Bipyramidal Geometry (TBP) 2-17 l Cl P Cl Cl Cl Br 86 Xe S S Note: Lone pair are placed equatorial to minimize repulsion Octahedral Geometry 2-18 S I Xe Note: Lone pair are placed such to minimize repulsion
Exceptions to VSEPR 2-19 Transition metal compounds do not follow VSEPR rules Species that are sterically crowded often do not obey VSEPR Xe 6 TeCl 6 2 Molecular Orbital Theory 2-20 In principle, the electronic structure of molecules can be worked out in the same way as for atoms: > solve the Schrödinger equation This gives molecular orbitals rather than atomic orbitals But: It is difficult to solve the Schrödinger equation for molecular species (only through approximation!)
LCAO approximation 2-21 LCAO = Linear Combination of Atomic Orbitals The wavefunctions of molecular orbitals can be approximated by taking linear combinations of atomic orbitals Ψ = 1 σ Ψ + 1s( Ha) Ψ1s( Hb) 2 [ ] linear combination (addition) of the wavefunction from two 1s orbitals LCAO approximation 2-22 A second MO (molecular orbital) can be obtained via subtraction of two AOs 1 Ψσ * = Ψ1s( Ha) Ψ1s( Hb) 2 [ ] > the resulting wavefunction has a nodal plane perpendicular to the H H bond axis (electron density = zero); the of an electron in this orbital is higher compared to the additive linear combination = antibonding orbital nodal plane linear combination (subtraction) of the wavefunction from two 1s orbitals
irst Period Diatomic Molecules 2-23 Linear Combinations of p z Orbitals 2-24 Addition of two p z AOs results a bonding σ p MO, subtraction will give an antibonding σ p * MO with a nodal plane perpendicular to the bond axis [ ] 1 Ψσ = Ψ2 ( H ) + Ψ2 ( H ) 2 p z a p z b [ ] 1 Ψσ * = Ψ2 ( H ) Ψ2 ( H ) 2 p z a p z b
Linear Combinations of p x and p y Orbitals 2-25 Addition of two p x (or p y ) AOs results a bonding π p MO containing a nodal plane along the bond axis: Subtraction results an antibonding π p * MO with two nodal planes: one plane perpendicular and one parallel to the bond axis Energy Level Diagram 2-26 Electrons are filled according to the same guidelines as for multielectron elements (Aufbau principle)
2-27 Energy Level Diagram Rules for the Use of MOs 2-28 When two AOs to give MOs, two MOs will be produced or mixing AOs must have similar energies Each orbital can have a total of two electrons (Pauli principle) Lowest orbitals are filled first (Aufbau principle) Unpaired electrons have parallel spin (Hund s rule) Bond order = 1/2 (bonding electrons antibonding electrons)
Molecular Oxygen 2-29 AOs AOs Bond order = unpaired electrons Molecular luorine 2-30 AOs AOs Bond order =
Neon 2-31 AOs AOs Bond order = Orbital Mixing 2-32 Orbitals with similar interact, if they have the appropriate symmetries The σ 2p and σ 2s orbitals are symmetry related and give rise to two new orbitals, one with higher and one with lower
2-33 Note: With mixing the σ g orbital is higher in than the π 2p orbitals 2-34 Energy Levels
Boron Molecule 2-35 Bond order = unpaired electrons Carbon Molecule 2-36 Bond order = unpaired electrons
Photoelectron Spectroscopy 2-37 UV-photoelectron spectroscopy can be used to verify the MO level diagram: Molecules are ionized with monochromatic light: N 2(g) + hν N 2(g) + + e the kinetic of the resulting photoelectrons is measured Photoelectron Spectrum of Nitrogen 2-38 σ 2p π 2p σ 2s Note: the orbital energies may shift when an electron is removed
Molecular Nitrogen 2-39 According to calculations the σ g orbital is higher in than the two π 2p orbitals: Bond order = unpaired electrons Beryllium Molecule 2-40 Bond order =
Lithium Molecule 2-41 Bond order = unpaired electrons Bond Order vs. Bond Length & Energy 2-42