VSEPR Theory • Lewis structures are valuable because?? • But what about the 3-D shape of molecules? • They have 3-D shapes, but is it important? • What about enzymes: is their 3-D shape important? • What about DNA: is their 3-D shape important?
VSEPR Theory • Just as the 3-D shape of proteins and enzymes are crucial for their proper function, many chemical reactions are very shape specific. • Many rxns will only occur if a molecule of the “right shape” comes along. • So how do we predict the 3-D shape of molecules? • VSEPR Theory: valence shell electron pair repulsion theory.
VSEPR Theory • The underlying principle behind VSEPR is that electron pairs (whether bonding pairs or nonbonding pairs) repel each other. • So molecules want to maximize the distance between electron pairs. • But this means that there are “bond angles” between electron pairs.
VSEPR Theory • But bond angles means 3-D geometric shapes! • So if we can predict the bond angles between electron pairs, we can predict the geometry and thus the 3-D shape. • We still use Lewis structures, as to find the bond angles you must first draw the Lewis structure.
VSEPR Terminology • Electron pairs: pairs of electrons which are within the electron cloud of an atom (they may be bonding or nonbonding). • Nonbonding pairs (lone pairs): pairs of electrons which are not used in bonding; they are not “shared” between 2 atoms. • Nonbonding electrons: electrons which are not used in bonding.
VSEPR Terminology • Electron Domain (electron charge clouds): the electron domain is a number between 2 and 6. It is: • ED = (#atoms bonded to central atom) + (#lone pairs on the central atom)* • *Although we start with finding the electron domain of central atoms, we can use the same technique to find the electron domain of any atom in a molecule.
VSEPR Terminology • VSEPR Geometry: based on the ED, the molecule is assigned a 3-D shape around the central atom (or atom of interest). There are 5 basic VSEPR geometries. • Molecular Geometry (molecular shape): Based on the ED and the number of lone pairs on an atom, the actual 3-D shape of the molecule around the atom is assigned. There are 13 different molecular shapes. • You need to know them all!
VSEPR Prediction To predict the 3-D shape: Draw the Lewis structure Determine how many lone pairs are on the atom of interest. Determine the ED Determine the VSEPR geometry and the molecular geometry around the atom of interest.
Covalent Bonding and Orbital Overlap • Lewis structures and VSEPR do not explain why a bond forms. • How do we account for shape in terms of quantum mechanics? • What are the orbitals that are involved in bonding? • We use Valence Bond Theory: • Bonds form when orbitals on atoms overlap. • There are two electrons of opposite spin in the orbital overlap. • The more the orbitals overlap, the stronger the bond. • When p or d orbitals are involved, then the resulting bond has direction (somewhere in x, y, z axis).
Covalent Bonding and Orbital Overlap This works great for a few simple molecules, but it fails for most! Example: CH4. If you look at the orbital energy diagram for C, C only has 2 unpaired electrons. (the 2p2 orbital electrons) How does C form 4 bonds?? So chemists had to develop a theory to explain how orbitals overlap to form covalent bonds.
Hybrid Orbitals • sp Hybrid Orbitals • Consider the BeF2 molecule (experimentally known to exist): • Be has a 1s22s2 electron configuration. • There is no unpaired electron available for bonding. • We conclude that the atomic orbitals are not adequate to describe orbitals in molecules. • We know that the F-Be-F bond angle is 180˚(VSEPR theory). • We also know that one electron from Be is shared with each one of the unpaired electrons from F.
Hybrid Orbitals • sp Hybrid Orbitals • We assume that the Be orbitals in the BeF bond are • 180°apart. • We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding. • BUT the geometry is still not explained. Why? • We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital (process called hybridization). • The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital.
Hybrid Orbitals sp Hybrid Orbitals The two lobes of an sp hybrid orbital are 180°apart.
Hybrid Orbitals • sp Hybrid Orbitals • Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be. • Note: If you start with x number of orbitals, you end up with x orbitals! • Once the 2 sp hybrid orbitals are formed, they overlap with 2 F atoms to form the 2 covalent bonds. • Also, if the Electron Domain of an atom is 2, then it is sp hybridized!
Hybrid Orbitals • sp2 Hybrid Orbitals • Important: when we mix n atomic orbitals we must get n hybrid orbitals. • sp2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.) • The large lobes of sp2 hybrids lie in a trigonal plane. • All molecules with Electron Domain of 3 have sp2 orbitals on the central atom.
Hybrid Orbitals • sp3 Hybrid Orbitals • sp3 Hybrid orbitals are formed from one s and three p orbitals. • Therefore, there are four large lobes. • Each lobe points towards the vertex of a tetrahedron. • The angle between the large lobes is 109.5° • All molecules with an Electron Domain of 4 are sp3 hybridized.
Hybrid Orbitals • Hybridization Involving d Orbitals • Since there are only three p-orbitals, trigonal bipyramidal and octahedral electron pair geometries must involve d-orbitals. • Trigonal bipyramidal geometries require sp3d hybridization. • Octahedral geometries require sp3d2 hybridization. • Note the Electron Domain from VSEPR theory determines the hybridization.
Hybrid Orbitals • Summary • To assign hybridization: • 1) draw a Lewis structure; 2) assign the Electron Domain using VSEPR theory; 3) from the Electron Domain, determine the hybridization; and 4) name the geometry by the positions of the atoms.
Multiple Bonds • σ-Bonds: electron density lies on the axis between the nuclei. (Head-on overlap.) • All single bonds areσ-bonds. • π-Bonds: electron density lies above and below the plane of the nuclei. (Side-by-side overlap.) • A double bond consists of oneσ-bond and oneπ-bond. • A triple bond has one σ-bond and two π-bonds. • Usually the p-orbitals involved inπ-bonding come from unhybridizedorbitals.
Multiple Bonds • Ethylene, C2H4, has: • oneσ- and oneπ-bond; both C atoms sp2 hybridized; • both C atoms with trigonal planar electron pair and molecular geometries.
Multiple Bonds • Consider acetylene, C2H2 • the electron pair geometry of each C is linear; • therefore, the C atoms are sp hybridized; • the sp hybrid orbitals form the C-C and C-Hσ-bonds; • there are two unhybridizedp-orbitals; • both unhybridizedp-orbitals form the twoπ-bonds; • oneπ-bond is above and below the plane of the nuclei; • oneπ-bond is in front and behind the plane of the nuclei. • When triple bonds form (e.g. N2) oneπ-bond is always above and below and the other is in front and behind the plane of the nuclei.
Multiple Bonds • DelocalizedπBonding So far all the bonds we have encountered are localized (or fixed) between two nuclei. In the case of benzene, C6H6, there are 6 C-Cσbonds, 6 C-Hσbonds, each C atom is sp2 hybridized, • there are 6 unhybridizedporbitals on each C atom.
Multiple Bonds Delocalized p Bonding
Multiple Bonds • DelocalizedπBonding In benzene there are two options for the 3πbonds localized between C atoms or delocalized over the entire ring (i.e. theπelectrons are shared by all 6 C atoms). You can draw 2 resonance structures for benzene. Experimentally, all C-C bonds are the same length in benzene. • Therefore, all C-C bonds are of the same type (recall single bonds are longer than double bonds).
Multiple Bonds Delocalized p Bonding
Multiple Bonds • General Conclusions • Every two atoms share at least 2 electrons. • Two electrons between atoms on the same axis as the nuclei areσbonds. • σ-Bonds are always localized. • If two atoms share more than one pair of electrons, the second and third pair formπ-bonds. • When resonance structures are possible, delocalization is also possible.
Molecular Orbitals • Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization. (E.g. why does O2 interact with a magnetic field?; Why are some molecules colored?) • For these molecules, we use Molecular Orbital (MO) Theory.
Molecular Orbitals • In Valence Bond Theory, atomic orbitals were hybridized, but they were still atomic orbitals: that is, these orbitals belong to the atom. • Molecular Orbital Theory takes atomic orbitals and creates molecular orbitals, that is, orbitals which belong to the entire molecule.
Molecular Orbitals • Molecular orbitals: • each contain a maximum of two electrons; • have definite energies; • can be visualized with contour diagrams; • a molecule has as many molecular orbitals as it originally had atomic orbitals; • come in 3 types: bonding, nonbonding, and antibonding; • are mathematical combinations of the atomic orbitals; • are associated with an entire molecule.
Molecular Orbitals • The Hydrogen Molecule When two AOs overlap two MOs form. Therefore, 1s (H) + 1s (H) must result in two MOs for H2: • one has electron density between nuclei (bonding MO); • one has little electron density between nuclei (antibonding MO). MOs resulting from sorbitals areσMOs. The σBonding MO is lower energy thanσ*(antibonding) MO.
Molecular Orbitals The Hydrogen Molecule
Molecular Orbitals • The Hydrogen Molecule • Energy level diagram or MO diagram shows the energies and orbitals in an orbital. • The total number of electrons in all atoms are placed in the MOs starting from lowest energy (σ1s) and ending when you run out of electrons. • Note that electrons in MOs have opposite spins. • H2 has two bonding electrons. • He2 has two bonding electrons and two antibonding electrons.
Molecular Orbitals The Hydrogen Molecule
Molecular Orbitals • Bond Order Bond Order = ½(bonding e- - antibonding e-). Bond order = 1 for single bond. Bond order = 2 for double bond. Bond order = 3 for triple bond. Fractional bond orders are possible. Bond order for H2 = ½(bonding e- - antibonding e-) = ½(2 -0) = 1. Therefore, H2 has a single bond. • Bond order for He2 = ½(bonding e- - antibonding e-) = ½(2 - 2) = 0. • Therefore He2 is not a stable molecule.