lectures Handout 1 - Bonding Dr Jonathan Burton

Transcription

1 Introduction to rganic hemistr 1 Introduction to rganic hemistr lectures andout 1 - Bonding Dr Jonathan Burton rganic hemistr J. laden,. Greeves, S. Warren Stereochemistr at a Glance J. Eames & J. M. Peach The majorit of organic chemistr tet books have good chapters on the topics covered b these lectures Eliel Stereochemistr of rganic ompounds (advanced reference tet)

2 Introduction to rganic hemistr 2 ourse utline These seven lectures will focus on bonding, stereochemistr, and an introduction to mechanism (curl arrows). All of these areas are crucial to being able to rationalise, and more importantl predict, what will happen when organic molecules are treated with an of a plethora of reagents used b organic chemists and chemists in general. Aspects of bonding will receive much more in-depth treatments in both the inorganic and phsical and theoretical lecture courses throughout the degree and here we will meet simple, but useful views of bonding, that are fit for our purposes. url arrow mechanisms will be met in virtuall all of the organic chemistr lecture courses and this course will introduce the ke basic principles for drawing reasonable curl arrow mechanisms. Uniquel, this course will provide virtuall all of the content ou require regarding stereochemistr. Stereochemistr is chemistr in three-dimensions and developing skills to be able to manipulate molecules in 3-D in one s own head will prove important. Practising organic chemists frequentl build ball-and-stick models of molecules to better aid visualisation and a cheap, plastic molecular model kit can be ver useful when working though stereochemical problems and will be useful throughout our degree a decent one can be purchased here: Snopsis Bonding Lewis structures and bonding, formal charge, basic VSEPR, hbridisation, molecular orbitals, structures of some functional groups, drawing organic molecules, skeletal representations Stereochemistr conformation, stereochemical definitions, optical activit, configuration, chiralit, enantiomers and diastereomers, IP rules, compounds with more than one stereocentre, meso compounds, racemisation, planar and helical chiralit, resolution, relative and absolute configuration Mechanism acid base reactions, curl arrows, delocalisation

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5 Introduction to rganic hemistr 5 Atoms Bohr atomic model - electrons moving in specific orbits, s, p, d, f. Quantum number ma. no. of electrons in shell 1 2e 1 1s 1 2 8e 3 Li 2s 1 atomic number = number of protons 4 Be 2s 2 5 B 2s 2,2p 1 6 2s 2,2p 2 7 2s 2,2p 3 8 2s 2,2p 4 9 F 2s 2,2p 5 3 8e 11 a 12 Mg 13 Al 14 Si 15 P 16 S 17 l 18 Ar 2 e 1s 2 10 e 2s 2,2p 6 electrons in outer shell (valence shell) Lewis Bonding Theor (1916) historic and useful Bonding = achievement of completel filled (or empt) valence shells b sharing (covalent bonds) or transfer of electrons (ionic bonds) between atoms Lewis bonding forms a simple basis for the qualitative understand of the structures and properties of organic molecules ou will meet more sophisticated theories of bonding particularl molecular orbital theor as the course progresses Gilbert.. Lewis

6 Introduction to rganic hemistr 6 Atoms Electronegativit (c chi) measure of the relative abilit of an atom to attract valence electrons to itself (Pauling Scale 1932, empirical) 2.2 Li 1.0 a 0.9 Be 1.6 Mg 1.3 increasing electronegativit B 2.0 Al Si 1.9 the following order of electronegativities can be useful to remember: F > > l > Ionic Bonds usuall occur between atoms at opposite ends of the periodic table i.e. between atoms that have ver different electronegativities 3.0 P 2.2 the atoms involved in ionic bonding achieve filled shells b the transfer of electrons (usuall 1 or 2 electrons) 3.4 S 2.6 F 4.0 l 3.2 Br 2.9 I 2.6 increasing electronegativit Linus Pauling eamples 3 Li + (2s 1 ) Li F a l 9 F (2s 2,2p 5 ) 3 Li + 9 F (2s 0 2s 2,2p 6 )

7 Introduction to rganic hemistr 7 ovalent bonds generall between atoms of similar electronegativit sharing of electrons + (1s 1 ) (1s 1 ) covalent bond each atom has 2 electrons around it i.e. 1 st quantum number shell is full 9 F F 1 + (1s 1 ) (2s 2, 2p 5 ) both atoms have filled shells 2 electrons around 8 electrons around F in F the atoms have a large difference in nuclear charge and hence electronegativit the shared pair is therefore not equall distributed but displaced towards F this is bond polarisation b the inductive effect leading to a permanent dipole d + d - F the Greek letter d (delta) is used to indicate a partial charge organic molecules are held together b covalent bonds carbon is in the middle of the period, with no great difference in electonegativit with other elements ver difficult to form stable closed shell b ionisation i.e. 4+ and 4 Energ 2s 2 1s 2 (core) 2p 2

8 Introduction to rganic hemistr methane bonds formed b the combination of 4 atoms (1s 1 ) and 1 atom (2s 2, 2p 2 ) 4 + (1s 1 ) (2s 2,2p 2 ) d + d - 8 electrons 2 electrons electronegativit of carbon and hdrogen are similar c c = 2.6, c = 2.2, bonds slightl polarised 3 - ammonia lone pair not involved in covalent bond 3 + (1s 1 ) (2s 2,2p 3 ) 8 electrons 2 electrons 2 - water 2 + (1s 1 ) (2s 2,2p 4 ) 8 electrons 2 electrons 2 lone pairs Rules for Lewis structures line = shared pair = covalent bond no more drawing dots for electrons in bonds = unshared pair or lone pair formal charge on atom = number of valence electrons (number of lone pair electrons + ½ bonding pair electrons) for 2 nd row elements Li F the sum of the number of bonds plus the number of lone pairs cannot eceed 4 i.e. no more than 8 electrons around the atom Lewis octet rule

9 Introduction to rganic hemistr 9 formal charge charge assigned to each atom in a molecule assuming the electrons in the bonds are shared equall between atoms regardless of electronegativit bookkeeping tool to denote relationship between the bonding electrons and those formall belonging to each atom a useful formalism but must be interpreted with care formal charge on atom = number of valence electrons (number of lone pair electrons + number of bonds) (0 + 8/2) = +1 for 5 (0 + 8/2) = +1 for 6 (6 + 2/2) = -1 valence electrons lone pair electrons bonding electrons + + for 6 (2 + 6/2) = +1 (organic) chemists do not work out the formal charge on atoms, the just remember that whenever ogen has 8 electrons around it, and has 1, 2, or 3 bonds, it is negative, neutral or positive respectivel more generall, whenever an atom has 8 electrons around it, and has one more bond than its neutral state, it will be positivel charged, and one fewer bonds it will be negative hdronium ion in the hdronium ion the formal positive charge is drawn on ogen in realit the ogen, being more electronegative than hdrogen, does not have a positive charge and the positive charge is spread out over the 3 hdrogen atoms

10 Introduction to rganic hemistr 10 phsical origin of bonding forces (simplified) if two atoms form a bond, the bonded state must be more stable then the separated atoms -1 2 in covalent bonds the repulsion between the attraction each electron sees 2 electrons and between the nuclei is balanced b the positive nuclei ( ) attraction of the electrons to the positive nuclei at repulsion electron-electron and the stable internuclear distance the bond length -1 nucleus-nucleus ( ) lone pairs, a.k.a. non-bonded pairs, formall interact with onl one positive nucleus, hence their interelectron repulsion is greater than for a bond pair the are higher in energ and more available for bonding electron-electron repulsion in lone pair reactions which reduce this electron-electron repulsion are favoured + + in the product the lone pair electrons are now a bond pair and see 2 positive nuclei this is the Lewis basicit of lone pairs - frequentl the reactive site of a molecule towards electrophiles Lewis structures are useful in predicting the bonding in molecules but the tell us nothing of the shape of the molecules methane ethene (ethlene) ethne (acetlene) 2 2 hdrogen canide allene formaldehde (methanal)

11 Introduction to rganic hemistr 11 the shapes of the vast majorit of organic compounds are determined b repulsion between electrons in the valance shell (VSEPR - valence shell electron pair repulsion) i.e. bond pair bond pair repulsions bond pair lone pair repulsions bond pair bond pair repulsions geometr of molecules is determined b minimising the above interactions for methane, 4, there are 4 bonding pairs of electrons and the repel each other to be as far apart as possible leading to the tetrahedral structure of methane with ammonia there are still four groups to place around the central nitrogen atom 3 bond pairs and 1 lone pair 3 is therefore based on a tetrahedron its shape is 1.09 Å described as pramidal 1.01 Å (1 Å = m) tetrahedral geometr, bond angle bond angle bond length = 1.01 Å similarl with water there are still four groups to place around the central ogen atom 2 bond pairs and 2 lone pairs water 2 is therefore also based on a tetrahedron its shape is described as bent Å bond angles decrease in the order 4, 3, 2 because: lone pair / lone pair repulsions are greater than lone pair / bond pair repulsions which are greater than bond pair / bond pair repulsions

12 Introduction to rganic hemistr 12 some eamples BF 3 3F + (2s 2,2p 5 ) B (2s 2,2p 1 ) F B F F F B F F planar 120 bond angle B 6 electrons (no lone pair) F 8 electrons 3 + planar 120 bond angle, 6 electrons (no lone pair), 2 electrons Quantum umbers and Atomic rbitals atomic orbitals mathematical functions (wavefunctions) related to the probabilit of finding an electron in a particular region of space for organic chemistr we will generall be concerned with s and p-orbitals atomic orbitals come in sets associated with the principal quantum number n = 1, 2, 3, s p an atom s highest principal quantum number determines the valence shell of the atom - onl these orbitals are generall involved in bonding each row of the periodic table indicates a different principal quantum number for hdrogen the principal quantum number is 1 giving rise to the familiar 1s orbital for carbon the principal quantum number is 2 giving rise to the familiar 2s and three 2p orbitals

13 Introduction to rganic hemistr 13 s and p orbitals nucleus s p p p s-orbitals - spherical electron densit decreases in moving awa from the nucleus there are three p-orbitals, p, p and p, orthogonal to each other p-orbitals have phase (sign) and a node at the nucleus (where the wavefunction changes sign) the electron densit is ero at a node the order of occupanc of atomic orbitals for a ground state configuration follows the Aufbau principle 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f ground state electronic configuration i) orbitals filled lowest energ to highest energ following the arrows from top to bottom i.e. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc. ii) maimum two electrons in each orbital which must have opposite spins (Pauli eclusion principle) iii) empt orbitals of equal energ (e.g. 3 2p orbitals) take one electron each before pairing occurs (und s rule of maimum multiplicit)

14 Introduction to rganic hemistr 14 Molecular rbitals (simplified an organic chemistr view) molecular orbitals are formed from combining atomic orbitals we will look at 2 overlap of two hdrogen 1s atomic orbitals leads to two molecular orbitals, one lower in energ, and one higher in energ, than the original atomic orbitals electrons are fed into the lowest energ molecular orbital first + + destructive overlap of atomic orbitals leads to anti-bonding molecular orbital (internuclear node) constructive overlap of atomic orbitals leads to bonding molecular orbital 1s 1s (s*) antibonding molecular orbital (empt) molecular orbital diagram for 2 energ released in combining two hdrogen atoms to give 2 is the bond strength = 432 kjmol -1 s* bonding molecular orbital in 2 has clindrical smmetr and is termed a σ- orbital (sigma orbital) c.f. s-atomic orbital in 2 this orbital can be viewed as the (Lewis) bond and is called a σ-bond although the electron densit is distributed over the whole molecule Energ 1s 1s (s) bonding molecular orbital (full) s

15 Introduction to rganic hemistr 15 combinations of p-orbitals (with all orbitals, the strongest bonding occurs when maimum overlap of (atomic) orbitals occurs) in a similar manner we can combine p-orbitals to make bonding and antibonding s-molecular orbitals p-orbital p-orbital + + p-orbital p-orbital s bond s* s destructive overlap of atomic orbitals leads to anti-bonding s* molecular orbital internuclear node constructive overlap of atomic orbitals leads to a bonding s molecular orbital - often drawn as so the provenance of the molecular orbital is clear p-bonding p-orbitals can also overlap edge on which leads to p-bonds weaker than the corresponding s-bonds as the orbital overlap is not as good + p* destructive overlap of atomic orbitals leads to an anti-bonding p* molecular orbital internuclear node + p p bond as in ethene constructive overlap of atomic orbitals leads to a bonding p molecular orbital - often drawn as of the molecular orbital is clear so the provenance

16 Introduction to rganic hemistr 16 returning to 4 (methane) 2 centre, 2 electron bond organic chemists generall draw Lewis structures where the line connecting the atoms represents a localised 2 centre, 2-electron bond. ow can we rationalise this with a molecular orbital view of bonding involving overlap of orbitals? methane 4 Energ 2s 2 2p 2 carbon ground state 1s 2, 2s 2, 2p 2 the 1s electrons are core leaving the 2s and 2p valence electrons for bonding the 2s orbital is spherical (non-directional) and the three 2p orbitals point along their respective aes (i.e. at 90 to one another) 2p 1s 2 (core) 2s 2p 2p 2p how can s and p atomic orbitals accommodate the tetrahedral structure of methane (and the structures of man organic molecules)? one approach, initiall proposed b Pauling, is the idea of hbridisation a method of creating orbitals that reproduces the observed shapes of molecules atomic orbitals are just mathematical functions, so we can add and subtract them (as we have done to form simple molecular orbitals) full molecular orbital analsis (2 nd ear lecture course) is another approach

17 Introduction to rganic hemistr 17 hbridisation miing the 2s and three 2p orbitals gives four sp 3 hbrids that point towards the corners of a tetrahedron ± ± ± sp 3 hbrid orbitals 2s 2p 2p 2p four sp 3 hbrid orbitals these hbrid orbitals can overlap with the hdrogen 1s orbitals to form s-bonding orbitals (and the corresponding s* antibonding orbitals) constructive overlap of hbrid orbital with 1s atomic orbital + + leads to a bonding s molecular orbital (s bond) - often drawn as so the provenance of the molecular orbital is clear sp 3 hbrid orbital on hdrogen 1s orbital s bonding orbital a s bond the four sp 3 hbrid orbitals on carbon overlap with the 1s orbitals of the four hdrogen atoms to give methane the carbon atom is said to be sp 3 hbridised hdrogen 1s orbitals bonding hbrid molecular orbitals in methane corresponding to the Lewis structure bonds each bond = 435 kjmol -1

18 Introduction to rganic hemistr 18 as hbrid orbitals are directional the provide much better overlap with the 1s orbitals of the hdrogen atoms ethane 2 6 sp 3 hbrid orbitals can overlap with each other as in the bond in ethane to give a bonding s molecular orbital + sp 3 hbrid orbitals s orbital (bond) constructive overlap of hbrid orbitals leads to bonding s molecular orbital - often drawn as below so the provenance of the molecular orbital is clear + rotation around the ais does not alter the overlap of the sp 3 orbitals (the s orbital) and occurs readil to give a variet of conformers of ethane more of this later staggered conformation eclipsed conformation the end-on views are termed ewman projections

19 Introduction to rganic hemistr 19 hbridisation sp 2 miing the 2s and two 2p orbitals gives three sp 2 hbrids that lead to trigonal geometr (120 bond angles) sp 2 hbrid orbitals ± ± p forming three sp 2 hbrids leaves one p-orbital perpendicular to the plane containing the hbrid orbitals p p 2s three sp 2 hbrid orbitals (drawn in plane) the bonding in ethene/ethlene ( 2 4 ) is a result of overlap of sp 2 hbridised carbon atoms with each other to form the s-bond, and with four hdrogen 1s orbitals to form the four s-bonds the remaining two p-orbitals overlap to form a p-bond three sp 2 hbrid orbitals and perpendicular p-orbital around central (carbon) atom

20 Introduction to rganic hemistr 20 overlap of the sp 2 hbrid orbitals leads to three s-bonds two bonds and one bond overlap of the remaining perpendicular p -orbitals gives a p-bond overlap of p orbitals gives p-bond s-bond the two lobes of the p-bond overlap of sp 3 hbrid orbitals gives s-bond the p-bond leads to electron densit above and below the molecular plane rotation about the ais is impossible as this would reduce overlap of the p-orbitals and would break the p-bond as before the p-molecular orbital (p-bond) is generall drawn as two separate p-atomic orbitals for clarit X rotation about ais would break the p-bond and does not occur hence alkenes are configurationall stable (i.e. cis and trans)

21 Introduction to rganic hemistr 21 hbridisation sp miing the 2s and one 2p orbitals gives two sp hbrids that lead to linear geometr (180 bond angles) sp hbrid orbitals ± two sp hbrid orbitals p p p 2s p-orbitals p-bond formed b p-orbital overlap forming two sp hbrids leaves two p-orbitals which can form p-bonds the two sp hbrid orbitals are arranged 180 apart this is the bonding arrangement found in ethne (acetlene) with the sp hbrids overlapping with the hdrogen 1s orbitals (not shown) the remaining p orbitals overlapping to form the two π-bonds sp hbrid orbitals p-bond formed b p-orbital overlap the triple bond in acetlene is made up from one s-bond and two orthogonal p-bonds, the carbons are said to be sp hbridised sp-hbrids overlap to form s-bond view without s-orbitals

22 Introduction to rganic hemistr 22 a quick method to work out the hbridisation of an atom is to count the number of substituents on that atom (including lone pairs of electrons), remembering that in the bonded environment first row elements have 8 electrons around them (there are eceptions to this method which we will meet later) 4 substituents = sp 3 hbridised, 3 substituents = sp 2 hbridised, 2 substituents = sp hbridised allene 2 == 2 what is its molecular shape? terminal carbons bonded to 2 hdrogen atoms (s-bonds) and double bonded to central carbon (s-bond and p-bond) terminal carbon atoms are sp 2 hbridised central carbon bonded to 2 substituents has linear geometr sp hbridised central carbon forms a single s-bond with each terminal carbon (sp-sp 2 hbrid orbital overlap) sphbridised carbon (c.f. acetlene) no free rotation sp 2 hbridised carbon (c.f. ethene) 2 groups are perpendicular alternative perspective

23 Introduction to rganic hemistr 23 hbridisation can be used with atoms other than carbon the nitrogen atom in ammonia can be viewed as sp 3 hbridised as can the ogen atom in water although the -X- bond angle is slightl less than 109 due to lone pair bond pair repulsion -atom is sp 3 hbridised similarl for amines, alcohols, ethers etc. lone pair in sp 3 orbital -atom is sp 3 hbridised 3 methlamine 3 ethanol diethlether BF 3 three substituents around boron no lone pair planar, sp 2 hbridised, planar with empt p-orbital on boron F B F F 3+ - three substituents around carbon no lone pair planar, sp 2 hbridised, planar with empt p-orbital on carbon other carbocations have similar planar structure

24 Introduction to rganic hemistr 24 hbridisation in functional groups (sites of chemical reactivit) formaldehde (c.f. bonding in ethene) sp 2 sp 2 p-bond formed b overlap of adjacent p-orbitals lone pairs in sp 2 orbital bonding and reactivit is similar in aldehdes and ketones acetaldehde (ethanal) Me acetone (propanone) Me Me electronegativit difference results in polarisation of bond nucleophiles react on carbon and break the weaker p-bond u d + d - u conversel, lone pairs are sites of electron densit electrophiles react on ogen imines planar, lone pair in sp 2 orbital Me Me Me Me oimes planar, lone pair in sp 2 orbital nitriles e.g. 3 (c.f. bonding in acetlene) d + d - 3 lone pair in sp-orbital 3