Unit ne Parts 3 & 4: molecular bonding hanging the order we approach this material... gjr- - 1 Locating electrons and drawing Lewis structures Assigning formal charges Describe bonds in terms of orbital overlap and the formation of σ & π bonds Look at the effects of different bonds on the shape & properties of molecules Look at resonance and the movement of π electrons Meet the organic chemists best friend...the 'curly arrow' dr gareth rowlands; g.j.rowlands@massey.ac.nz; science tower a4.12 http://www.massey.ac.nz/~gjrowlan
Types of bonds gjr- - 2 3 Br 3 Br rganic chemists view the world as bonds being made and broken But what does this mean? And what is this bond malarkey anyway? Ionic bonds ne atom takes an electron from another atom Ionic bond formed by attraction between cation (+) & anion ( ) a l a l a + l a l a + l al ovalent bonds Two electrons are shared by two atoms So the bond (and our line ) represents two electrons
Electron configuration gjr- - 3 1 18 1 2 13 14 15 16 17 e Li Be B F e a Mg 3 4 5 6 7 8 9 10 11 12 Al Si P S l Ar K a Sc Ti V r Mn Fe o i u Zn Ga Ge As Se Br Kr energy 2s 2p x 2p y 2p z hydrogen 1s 1 1s 1 18 2 13 14 15 16 17 2 e e 3 6 Li Be Li B F e a Mg Al Si P S l Ar 3 4 5 6 7 8 9 10 11 12 K a Sc Ti V r Mn Fe o i u Zn Ga Ge As Se Br Kr energy 2p x 2p y 2p z 2s 2p x 2p y 2p z 2s carbon helium lithium 1s 2 2s 1s 2 1s 2px 2 2s 2 1 2py 1 1s 2 2s 2 2p 2 1s Aufbau Principle lowest energy orbital Pauli Exclusion Principle spin +½ or -½ und's rule electrons as far apart as possible (degenerate orbitals)
Valence electrons gjr- - 4 1s 2 2s 2 2p 2 2s 2 2p 2 1s 2 2s 2 2p 3 2s 2 2p 3 valence electrons group 1 2 13 3 14 4 15 5 16 6 17 7 18 8 e Li Be B F e Valence electrons highest energy shell Gives us a good idea of how molecules will bond 8 is the magic number (but not a catchy song title...)
Lewis structures The basis of simple Lewis structures is the ctet Rule Most atoms want 8 valence electrons to fill the outer shell 2s 2 2p 6 (please remember this is a simplification) ydrofluoric acid F gjr- - 5 + F F F Methanol 3 + + 4 Ethene 22 + + 4
Lewis structures II gjr- - 6 Acetone 33 3 + + 6 3 3 The atoms are not charged even though they are sharing 4 other electrons Lone pairs of electrons are pairs of valence electrons that are not shared In ions a valence electron is either added ( ) or lost (+) gained valence electron B + 3 + lost valence electron + 3 + Borohydride anion B4 B B Ammonium cation 4 +
Formal charges gjr- - 7 So far we know if a molecule is charged or not and in reality that charge is spread over the entire molecule It is useful to "localise" the charge on one atom - this is the formal charge fc fc = number of valence electrons number of unshared electrons ½ number of shared electrons + 3 + nitrogen; fc = 5-0-½(8)=+1 cation + + 3 ozone neutral left-hand oxygen; fc = 6-4-½(4)=0 central oxygen; fc = 6-2-½(6)=+1 right-hand oxygen; fc = 6-6-½(2)=-1 atom's formal charges
Atomic orbitals: its a quantum world out there So we have got an idea where the electrons (& bonds) are... But how are they formed? A number of different models...we will look at a very simplified version Atomic electrons reside in atomic orbitals (areas of space with a 90% probability of finding the electron) nly two electrons per orbital (one of spin +½ & one of spin ½) gjr- - 8 s rbitals + not charge so less confusing to draw as two s orbitals with different phases two s orbitals with different phases Spherical in shape Larger s orbitals (>1) have 1 or more nodes (0% probability of finding electron) hange in phase either side of node Phase is just a result of maths...so won't worry about it yet...
Atomic orbitals: p orbitals gjr- - 9 p x z p y z p z z y y y x x x The three p orbitals are degenerate (same energy) Just differ by the direction they are orientated
Atomic orbits gjr- - 10 hydrogen carbon 2p x 2p y 2p z energy 2s 2p x 2p y 2p z energy 2s 1s 1s 1s 1s 2s 2p x 2p y umber of electrons per atom given by atomic number Add electrons to lowest energy orbital first If orbitals have the same energy electrons will not pair up
Single bonds gjr- - 11 + σ* energy 1s σ 1s A bond is formed by the overlap of two atomic orbitals (and electrons) The overlap forms two molecular orbitals - the bonding σ & anti-bonding σ* verlap with an s orbital always gives a σ bond (symmetrical) More information can be obtained at: www.massey.ac.nz/~gjrowlan/teaching.html
gjr- - 12 Single σ bond in 2 bonding σ molecular orbital (M) anti-bonding σ* molecular orbital Similar M obtained form the overlap of s & p So bonds are σ bonds similar to above In saturated hydrocarbons σ bonds are made from 1s and sp 3 The wonderful topic of hybridisation will be covered elsewhere...
energy Single bonds from p orbitals gjr- - 13 x z overlap end-on y overlap side-on What happens when we combine the p orbitals of two atoms? nly two can approach head-to-head these give a new σ bond Again, if carbon joined to 4 other atoms then it is sp 3 σ* 2p y σ 2p y
Double bonds: side-to-side overlap of p orbitals gjr- - 14 = π* energy = π carbon 2p z carbon 2p z Side-to-side overlap gives π bond & π* anti-bonding Differ from σ bonds as they have no axis of symmetry (phase change) It appears to be two orbitals (above & below) but is only one Double bond prevents rotation resulting in cis trans isomerism
Double bond molecular orbitals gjr- - 15 Bonding π M Anti-bonding π* M In an alkene the two are joined by one σ bond & one π bond So a carbon joined to three groups has 3 x σ bonds & 1 x π bond Such a carbon is said to be sp 2 as 1 x s & 2 x p used in σ bond and remaining p used in π bond
is-trans isomerism gjr- - 16 3 3 3 3 3 multistep enzymecatalysed reverse process light isomerises complexed cis-retinal 3 3 3 3 3 cisretinal transretinal To change between cis and trans double bonds we have to break π bond and then reform it the bond simply will not twist around
ther bonds Triple bond gjr- - 17 σ π (2p z + 2p z ) σ σ π π π (2p y + 2p y ) An alkyne or triple bond comprises 1 x σ bond and 2 x π bonds A atom joined by 2 x σ bonds & 2 x π bonds is called an sp carbon as only 1 x s & 1 x p involved in σ bonds on-bonded electron pairs (lone pairs) Lone pairs occupy n (non-bonding) molecular orbitals
Example gjr- - 18 ( 3 ) 3 P 2 2 3 Br Draw the Lewis structure & identify the types of bond (σ or π) in above compound First we need to know the number of valence electrons... P Br ext attach the atoms & obey the octet rule (8 electrons)... P Br Finally, draw as normal and point out the bonds... π bond P 3 Br 3 σ bonds 3 P 3 Br
Geometries about carbon atoms The wonderfully titled Valence Shell Electron-Pair Repulsion (VSEPR) theory tells that pairs of electrons (bonds or lone-pairs) want to be as far apart as possible...(like charges repel) gjr- - 19 180 If a carbon has just two groups (e.g. ethyne) they will be directly opposite each other (in a straight line) and hence linear linear 120 trigonal planar If there are three groups (e.g. ethene), they will point to the corners of a triangle. All groups are arranged in the same plane (flat) and so we call this trigonal planar
Geometry II gjr- - 20 109 Br tetrahedral Br Four groups again want to be as far apart as possible and this gives the tetrahedral arrangement This is probably the most important shape in organic chemistry 3 2 3 sp 1 linear sp 2 trigonal planar dynemicin A sp 3 tetrahedral
Bond strengths gjr- - 21 bond energy > > Bond Strength - energy to break a bond Multiple bonds are stronger than single bonds BUT - σ bond is stronger than a π bond This is a result of orbital overlap bond energy >
Bond strengths and lengths gjr- - 22 F 120 138 F l bond energy 134pm bond energy 178pm l 154pm 193pm Br Br Bond lengths - average distance between two linked atoms Shorter the bond, the stronger it is Explain the following: 134pm 122pm 610kJmol 1 736kJmol 1 For a full list of bond strengths and lengths see the course Study Guide (Pg34)
Bond Polarisation Is l covalent or ionic? gjr- - 23 δ+ δ l l l Polar covalent bond - electrons in the bond are shared BUT polarised (attracted) towards one atom more So covalent bond has a degree of ionic character Electronegativity E - the more electronegative (higher number) atom attracts the electrons 2.1 Li 1.0 1.5 Be a 0.9 Mg 1.2 K 0.8 1.0 a Rb 0.8 Sr 1.0 B 2.0 Al 1.5 2.5 Si 1.8 3.0 P 2.1 3.5 S 2.5 F 4.0 l 3.0 Br 2.8 I 2.5 Bond Type E difference Examples alculation ionic > 1.7 al 3.0(l) - 0.9(a) = 2.1 polar covalent 0.5 1.7 covalent 0 0.4 3 l 3 3.5() - 2.1() = 1.4 3.0(l) - 2.1() = 0.9 2.5() - 2.1() = 0.4 2.1() - 2.1() = 0.4
Bond Polarisation II An understanding of the polarisation of bonds in organic molecules helps us to understand the chemistry of these compounds gjr- - 24 δ+ δ δ 3 I 3 δ+ 3 δ δ+ δ+ Me δ 3 Mg δ+ δ Br It also explains why = bond is stronger than = (yet more reactive) π* π* energy energy p A carbon p A carbon π p A carbon π p A oxygen δ+ δ
Resonance gjr- - 25 + 3 + + 2 3 top oxygen; fc = 6-4-½(4)=0 nitrogen; fc = 5-0-½(8)=+1 bottom oxygen; fc = 6-6-½(2)=-1 If we draw the Lewis structure for nitromethane 32 we get... So the two are different? But...(and isn't there always a 'but' in chemistry...) Both bond lengths are 122pm ( = 130pm & = = 116pm) Reason... Moving the electrons gives us another acceptable Lewis structure These structures are resonance structures The truth (or reality) is a resonance hybrid - somewhere in the middle
Resonance II gjr- - 26 Resonance structures - two or more acceptable Lewis structures nly differ by the position of the electrons They do not exist - they are extremes - reality is in the middle The 'curly arrow' is represents the movement of 2 electrons In many respects it is the key to organic chemistry Pushable electrons e.g. Receptors e.g. lone pairs, positively charged atoms π bond electrons atoms that can accept electrons atoms which have pushable electrons
Delocalisation gjr- - 27 Resonance is not the movement of electrons, our structures are just extremes The electrons are spread over the system or delocalised 3 3 3 3 or 3 1 / 2 1 / 2 Electrons spread over three atoms - bond lengths are equal (130pm) Ph 3 3 X Ph Ph 3 3 Ph 3 3 X Ph δ 3 δ 3 Ph 3 3
Delocalisation II eutral molecules also have resonance forms gjr- - 28 3 3 δ 3 δ+ Again this explains the physics observations bond of ethanoic acid is far shorter than in an alcohol 124pm 129pm 3 ethanoic acid 3 2 ethanol 146pm 122pm 3 3 propanone Most famous example of delocalisation is probably benzene
Delocalisation III gjr- - 29 We can use resonance and delocalisation to explain why phenol is acidic (and causes burns) whilst an alcohol such as ethanol is not δ δ δ δ
onjugation gjr- - 30 Any double bond separated from another double bond or charge or lone pair can delocalise Such systems are said to be in conjugation Systems with a high degree of conjugation are frequently coloured... 3 3 3 3 3 3 3 3 3 3 3 3 3 3 3 3 3 3 3 3
verview gjr- - 31 What have we learnt? Quite a bit... Last two lectures have provided information about bonding We have learnt how to find electrons We have started to learn how electrons move We have introduced resonance and delocalisation What's next? A look at the forces between molecules ow polarisation and shape effects intermolecular forces Start to recognise which molecules are polarised dr gareth rowlands; g.j.rowlands@massey.ac.nz; science tower a4.12 http://www.massey.ac.nz/~gjrowlan