Structure and Bonding 1 PDF

Structure & bonding_1 Structure & bonding – Bonding in carbon compounds: the C atom, bonds as MOs, Energy, Electrons and the covalent bond – Bonding in carbon compounds: hybridisation, sp3, sp2, sp, other atoms, polarization, MOs from ‘filled’ orbitals – Stereochemistry and conformation – Reactivity...

Structure & bonding_1 Structure & bonding – Bonding in carbon compounds: the C atom, bonds as MOs, Energy, Electrons and the covalent bond – Bonding in carbon compounds: hybridisation, sp3, sp2, sp, other atoms, polarization, MOs from ‘filled’ orbitals – Stereochemistry and conformation – Reactivity and mechanism: general types, reaction arrows, stabilization of ions and radicals  This section covers the basic structural aspects of organic chemistry providing general concepts applicable to most organic compounds 1 Hydrogen-like atomic levels Quantum mechanics describes electrons around a nucleus of an atom in terms of wave functions or orbitals The first few orbitals and sub-orbitals for a ‘hydrogen-like’ atom are listed 2 Atomic orbital shapes z x-axis y s orbital x,y,z-axes centred on nucleus s orbital shown with x,y,z-axes p orbital px orbital py orbital pz orbital +z -y -x +x +y -z half of pz orbital along + side of z-axis half of pz orbital along - side of z-axis What do these orbitals look like? x, y and z-axes can be set up at the nucleus Orbital shapes are shown for the s and p types The p sub-orbitals can also have 3 different orientations, called px, py and pz p orbitals are often shown with differently coloured ‘halves’ to reflect the part on the (+) side of the axis and on the (-) side. Each sub-orbital can ‘hold’ up to 2 electrons 3 The carbon atom 2s 2px 2py 2pz energy 1s Sometimes it is useful to consider the orbital as the showing the probability of an electron being in that space The H-like model can be applied to the smaller atoms, including carbon The 6 electrons of neutral C fill the s and p sub-orbitals as shown Unlike H, 2s and 2p sub-orbitals in C have slightly different energy 4 A simple orbital view of bonding H H 1s H + H 1s H H H2 molecule The simple Bohr model of a covalent bond is too …. Simple! H 1s + H 1s H2 molecule A covalent bond is formed by overlap of 1s orbitals (each with1 e) on different H and sharing of the electrons Overlap of the atomic orbitals gives a molecular orbital (MO) The 2 shared electrons fill this MO 5 Anti-bonding orbitals The quirks of quantum mechanics mean that an antibonding molecular orbital is also possible H 1s - H 1s H2 antibonding Since the 2 electrons are already used in the bonding molecular orbital, this one is usually empty But it can exist and be populated by an electron in some cases 6 A more advanced view of bonding H2 antibonding MO H 1s H 1s H2 bonding MO Quantum mechanics calculates the bonding MO energy to be 436 kJ/mol lower than the ‘free’ H 1s orbital energy (stabilisation energy) So molecular H2 with the 2 electrons in the new bonding MO is more stable (lower energy) than two free H atoms This is the principle for all covalent bonding The antibonding MO is higher in energy than the free atoms 7 Bonding using other orbitals  molecular anti-bonding orbital sp3 atomic (hybrid) orbital sp3 atomic (hybrid) orbital Orbitals other than s can overlap to form MOs Overlap ‘head-tohead’ gives a ‘sigma’ (s) MO Those shown here are sp3 hybrid orbitals – we will find out about hybrid orbitals later  molecular bonding orbital 8 Bonding using other orbitals In organic chemistry pz orbitals do not overlap ‘headto-head’ They prefer a ‘sideways’ (p) overlap NOTE: Each MO is a single orbital even if there is a gap (the antibonding one) in the middle! p antibonding MO C 2pz C 2pz p bonding MO 9 Bonding using orbitals - summary A covalent bond is a molecular orbital formed by overlap of two atomic orbitals on different atoms The molecular orbital (or more accurately the electrons in it) has a lower energy than the original atomic orbitals A rough idea of what a MO looks like can be obtained by simply squashing the two atomic orbital together An antibonding MO is also formed at higher energy, but is usually empty Lower energy mean more stable – good! 10 Bonding in carbon compounds The geometry of C compounds is not well explained by using the simple 2s, 2px and 2py orbitals Experimental geometries are found to be tetrahedral (Td), trigonal planar (Tg) or linear (Lr) 11 sp3 Hybridisation 2s 2px 2py 1s 2pz sp3 sp3 sp3 sp3 1s C-atom sp3 hybridised C-atom The correct C geometry can be obtained from the AOs by hybridisation, i.e. mixing of the 2s, 2px, 2py and 2pz orbitals This gives 4 new equivalent hybrid orbitals Each with 1 electron Pointing to the corners of a tetrahedron C a single sp3 hybrid orbital 2s FOUR sp3 hybrid orbitals 2px 2py 2pz replaced by... C 12 sp3 Hybrids and s bonds H 1s atomic orbital C sp3 atomic (hybrid) orbital H C C C H sigma (s) bond Overlap of a C sp3 hybrid with a H 1s gives a C-H s bond Each contributes 1 electron, so the s MO (or bond) has 2 electrons and is filled Repeating this for all hybrids gives 4 x C-H s bonds and the molecule methane Note: minor lobes (red part on previous slide) are often omitted for clarity 13 Csp3-H s bonds H C C-H  molecular anti-bonding orbital C H sp3 atomic (hybrid) orbital 1s atomic orbital H What does this C sp3 and H 1s interaction look like in MO terms? The overlap creates a CH s bonding MO (or more simply a C-H s bond) with 2 electrons And a C-H s* antibonding MO which is empty C C-H  molecular bonding orbital 14 Csp3-Csp3 s bonds C sp3 atomic (hybrid) orbital C sp3 atomic (hybrid) orbital C C C C sigma (s) bond C C An sp3 hybrid on one C can overlap with an sp3 on another C A C-C s MO (or s bond) is formed If the remaining sp3 orbitals form bonds to H, this gives ethane If further sp3 C are added, this gives propane, butane, etc. 15 Csp3-Csp3 s bonds C The overlap creates a C-C s bonding MO (or more simply a C-C s bond) with 2 electrons And a C-C s* anti-bonding MO which is empty C C-C  molecular anti-bonding orbital C C 3 sp atomic (hybrid) orbital sp3 atomic (hybrid) orbital C C Note: Do not worry if the shapes are not identical to other pictures, e.g. slide 8, or textbooks. It is the general concept (orbital density between atoms for bonding, ‘gap’ in orbital density for anti-bonding, etc.) that is important here C-C  molecular bonding orbital 16 Organic Structure, Reactivity and Functional Groups Deeper Study 17 Orbital shapes Ψ ( 1 𝑠 )=𝐶 𝑒 −𝑟 / 𝑎 s orbital shown with x,y,z-axes +z -y -x +x +y -z half of pz orbital along + side of z-axis half of pz orbital along - side of z-axis +z Q -z 𝑊 ( 2𝑝 𝑧 ) =𝐶 ′ 𝐶 𝑟 {𝑒} ^ {− 𝑟 /2 𝑎} 𝑐𝑜𝑠 Why is the s orbital all ‘blue’ (+), but the pz blue/red (+ and – parts)? The orbitals are mathematical solutions to the Schrödinger Equation H atom 1s is shown, C and a are constants, while r is the distance from the origin. – The term e-r/a can only be (+) H atom 2pz is also shown where C’, C” and a are constants. – r and the term e-r/2a can only be (+). – However, cos is (-) for  > 90o, i.e. along (-) part of z-axis 18 Antibonding molecular orbitals H2 antibonding MO s orbital shown with x,y,z-axes Ψ ( 1 𝑠 )=𝐶 𝑒 −𝑟 / 𝑎 H 1s H 1s H2 bonding MO + Ψ 𝐻 =Ψ 𝐴 ( 1 𝑠 )+ Ψ 𝐵 (1 𝑠) 2 YA(1s) YA(1s) Ψ 𝐻 =Ψ 𝐴 ( 1 𝑠 ) −Ψ 𝐵 ( 1 𝑠 ) 2 Ψ 𝐻 =Ψ 𝐴 ( 1 𝑠 )+ {− Ψ 𝐵 ( 1 𝑠 ) } YA(1s) 2 + YA(1s) Why has the H2 antibonding MO got blue and red parts and why is there a gap in the middle? Remember the 1s orbital is a mathematical function The bonding interaction adds two 1s from two different H-atoms But subtracting is also valid for the Schrödinger Equation Subtraction A – B is the same as A addition of (-B), i.e. A + (-B) So one of the H(1s) becomes red ((-) mathematical sign) Where the blue and red overlap (in the middle) the (+) and the (-) maths cancel each other out 19