1.4 Sigma(o-)and pi(-)bonds 3 Atom(s) N,P 0,s F,Cl,Br,I Group numb 14 15 16 17 4 3 2 1 to atom 6- -10 Example:Nitric acid (HNO Nitrogen with 4 covalent bonds has a formal charge of+1 benzene (Section 7.2.2) Formal charge:15-4-0-10=+1 The nitrogen atom donates a pair of electrons to make this bond Carbon forms four covalent bonds.When only three covalent bonds are ave ei ither a formal nega tive cha arge or a formal The Section 4.3 .Carbanions-three covalent bonds to carbon and a formal negative charge Formal charge on C: outer electr ----3 14-3-2-10=-1 .Dep ation of a R ion called me negative charge is used to show the 2 non-bonding electrons an enolate ion (Section 84) Carbocarions-three covalent bonds to carbon and a formal positive charge. Formal charge on C: rond Carbocations are intermediates ina 14-3-0-10=+1 R SI reactions (Section 5.3.1.2) The positive charge is used to show the absence of 2 electrons 1.4 Sigma (o-)and pi(-)bonds The electrons shared in a covalent bond result from overlap of atomic orbitals to Molecular orbitals and chemical give a new molecular orbital.Electrons in Is and 2s orbitals combine to give sigma(-)bonds. When two Is orbitals combine in-phase,this produces a bonding molecular orbital
Formal charge = group number in periodic table number of bonds to atom number of unshared electrons – – Group number Normal number of 2 electron bonds Atom(s) 14 4 N, P 15 3 O, S 16 2 F, Cl, Br, I 17 1 C – 10 Example: Nitric acid (HNO3) H O N O O Nitrogen with 4 covalent bonds has a formal charge of +1 Formal charge: 15 – 4 – 0 – 10 = +1 The nitrogen atom donates a pair of electrons to make this bond Carbon forms four covalent bonds. When only three covalent bonds are present, the carbon atom can have either a formal negative charge or a formal positive charge. Carbanions–three covalent bonds to carbon and a formal negative charge. C R R R 8 outer electrons: 3 two-electron bonds and 2 non-bonding electrons Formal charge on C: The negative charge is used to show the 2 non-bonding electrons 14 – 3 – 2 – 10 = –1 Carbocations–three covalent bonds to carbon and a formal positive charge. C R R R 6 outer electrons: 3 two-electron bonds Formal charge on C: The positive charge is used to show the absence of 2 electrons 14 – 3 – 0 – 10 = +1 1.4 Sigma (s) and pi (p) bonds The electrons shared in a covalent bond result from overlap of atomic orbitals to give a new molecular orbital. Electrons in 1s and 2s orbitals combine to give sigma (s) bonds. When two 1s orbitals combine in-phase, this produces a bonding molecular orbital. Nitric acid is used in synthesis to nitrate aromatic compounds such as benzene (Section 7.2.2) The stability of carbocations and carbanions is discussed in Section 4.3 Carbanions are formed on deprotonation of organic compounds. Deprotonation of a carbonyl compound, at the a-position, forms a carbanion called an enolate ion (Section 8.4.3) Carbocations are intermediates in a number of reactions, including SN1 reactions (Section 5.3.1.2) Molecular orbitals and chemical reactions are discussed in Section 4.10 1.4 Sigma (s) and pi (p) bonds 3
4 Structure and bonding ⊙+© ， s-orbital s-orbital bonding molecular orbital When orbitals combine u-phase,this produces an molecular orbital. ⊙+ ● s-orbital s-orbital antibonding molecular orbital Electrons in p orbitals can combine to give sigma ()or pi ()bonds. .Sigma(o-)bonds are strong bonds formed by head-on overlap of two atomic orbitals. ● +● p-orbital p-orbital bonding p-p a-orbital ● +● c● p-orbital p-orbital antibonding p-p -orbital Alkenes have ac -C bone .Pi(-)bonds are weaker bonds formed by side-on overlap of two p-orbitals. nd and one weaker -bond (Section 6.1) p-orbital p-orbital bonding p-p -orbital p-orbital p-orbital antibonding p-porbital Only -or at-bonds are present in organic compounds.All single bonds are a-bonds while all multiple (double or triple)bonds are composed of one o-bond and one or two at-bonds. 1.5 Hybridisation Hund's mle stat that when filling up a set of orbitals of the same The ground-state electronic configuration of carbon is 1s22s22p2p,. The six electrons fill up lower energy orbitals before entering higher energy orbitals (Aufbau principle). .Each orbital is allowed a maximum of two electrons(Pauli exclusion principle) one orbital .The two 2p electrons occupy separate orbitals before pairing up (Hund's rule)
+ s-orbital s-orbital bonding molecular orbital When two 1s orbitals combine out-of-phase, this produces an antibonding molecular orbital. + s-orbital s-orbital antibonding molecular orbital Electrons in p orbitals can combine to give sigma (s) or pi (p) bonds. Sigma (s) bonds are strong bonds formed by head-on overlap of two atomic orbitals. + p-orbital p-orbital bonding p-p σ-orbital + p-orbital p-orbital antibonding p-p σ∗-orbital Pi (p) bonds are weaker bonds formed by side-on overlap of two p-orbitals. + p-orbital p-orbital bonding p-p π-orbital + p-orbital p-orbital antibonding p-p π∗-orbital Only s- or p-bonds are present in organic compounds. All single bonds are s-bonds while all multiple (double or triple) bonds are composed of one s-bond and one or two p-bonds. 1.5 Hybridisation The ground-state electronic configuration of carbon is 1s2 2s2 2px 1 2py 1 . The six electrons fill up lower energy orbitals before entering higher energy orbitals (Aufbau principle). Each orbital is allowed a maximum of two electrons (Pauli exclusion principle). The two 2p electrons occupy separate orbitals before pairing up (Hund’s rule). Alkenes have a CC bond containing one strong s-bond and one weaker p-bond (Section 6.1) All carbonyl compounds have a CO bond, which contains one strong s-bond and one weaker p-bond (Section 8.1) Hund’s rule states that when filling up a set of orbitals of the same energy, electrons are added with parallel spins to different orbitals rather than pairing two electrons in one orbital 4 Structure and bonding
1.5 Hybridisation 5 12px12py—2p2 2s Energy s hybrid orbit spHybridisation.For four single a-bonds-carbon is sphybridised (e.g.in methane,CH4).The the comersof a tetrahedron (15bodpoo itals move H ⊙ methane:4xC-H o-bonds .sp2Hybridisation.For three single a-bonds and one -bond the -bond one weaker a-bond (Section 6.1) H2C=CH2).The three sp'-orbitals point to the corners of a triangle(120 bond angle),and the remaining p-orbital is perpendicular to the sp plane. bond H ①六⊙ C-C a-bond ethene:4 x C-H a-bonds.1x C-C a-bond.1 x C-C x-bond sp Hybridisation.For two single g-bonds and two -bonds -the two n-bonds Alkynes havea C=C bond equire two p-orbitals,and her nce the carbon is sp hybridised (eg.in ethyne HC=CH).The two sp-orbitals point in the opposite directions (180 bond angle).and the two p-orbitals are perpendicular to the sp plane 2 -bonds HRCO0 CH sp hybridisation C-C a-bond ethyne:2x C-H o-bonds.1x C-C a-bond.2x C-C -bonds
Energy 1s 2s 2px 2py 2pz The carbon atom can mix the 2s and 2p atomic orbitals to form four new hybrid orbitals in a process known as hybridisation. sp3 Hybridisation. For four single s-bonds – carbon is sp3 hybridised (e.g. in methane, CH4). The orbitals move as far apart as possible, and the lobes point to the corners of a tetrahedron (109.5 bond angle). sp3 hybridisation 109.5° o ox o ox methane: 4 × C–H σ-bonds H H H H H x C x H H H sp2 Hybridisation. For three single s-bonds and one p-bond – the p-bond requires one p-orbital, and hence the carbon is sp2 hybridised (e.g. in ethene, H2CCH2). The three sp2 -orbitals point to the corners of a triangle (120 bond angle), and the remaining p-orbital is perpendicular to the sp2 plane. sp2 hybridisation p orbital 120° π-bond C–C σ-bond ethene: 4 × C–H σ-bonds, 1 × C–C σ-bond, 1 × C–C π-bond o o o o o o H H H H C X XX X H H C H H X X sp Hybridisation. For two single s-bonds and two p-bonds – the two p-bonds require two p-orbitals, and hence the carbon is sp hybridised (e.g. in ethyne, HCCH). The two sp-orbitals point in the opposite directions (180 bond angle), and the two p-orbitals are perpendicular to the sp plane. sp hybridisation p orbitals 180° C–C σ-bond 2 π-bonds o ooo o ethyne: 2 × C–H σ-bonds, 1 × C–C σ-bond, 2 x C–C π-bonds H X C XXX C X H H H Alkenes have a CC bond containing one strong s-bond and one weaker p-bond (Section 6.1) All carbonyl compounds have a CO bond, which contains one strong s-bond and one weaker p-bond (Section 8.1) Alkynes have a CC bond containing one strong s-bond and two weaker p-bonds (Section 6.1) 1.5 Hybridisation 5
6 Structure and bonding .For a single C-C or C-O bond,the atoms are sphybridised and the carbon atom(s)is tetrahedral For a double C=C or C=O bond,the atoms are sp2 hybridised and the carbon atom(s)is trigonal nlanar For a triple C=C or C=N bond,the atoms are sp hybridised and the carbon atom(s)is linear. This compound contains four 2H lgroups. H introduced in Section 2 C2 120° 3=5p32=5p21=sp The shape of organic molecules is therefore determined by the hybridisation hond in an all Bond Mean bond enthalpies(kJ mol-1)Mean bond lengths(pm) +347 153 c=c +612 734 +838 120 The shorter the bond length,the st the'scha the electrons are held closer to the nucleus expla ed bythe changein SD3 s02 HaC-CH2TH HC=CH-H HC=CTH (Section 1.74) longest shortest A single C-C o-bond can undergo free rotation at room temperature,but a t-bond prevents free rotation around a C=C bond.For maximum orbital overlap in a it-bond,the two p-orbitals need to be parallel to one another.Any rotation around the C=C bond will break the at-bond. 1.6 Inductive effects,hyperconjugation and mesomeric effects 1.6.1 Inductive effects Inacovalent bond between two,the in the -bond are no sharedequally.The electrons are attracted towards the most electronegative atom.An
For a single CC or CO bond, the atoms are sp3 hybridised and the carbon atom(s) is tetrahedral. For a double CC or CO bond, the atoms are sp2 hybridised and the carbon atom(s) is trigonal planar. For a triple CC or CN bond, the atoms are sp hybridised and the carbon atom(s) is linear. C O C C C C C C C C N C C H O H H H H 3 3 H H H H 3 = sp3 2 2 2 2 2 2 2 2 2 = sp2 1 1 1 = sp C O C C C C C C C C N C C H O H H H H H H H H 109.5° 120° 120° 180° 2 120° 120° 120° 109.5° 120° 2 The shape of organic molecules is therefore determined by the hybridisation of the atoms. Functional groups (Section 2.1) that contain p-bonds are generally more reactive as a p-bond is weaker than a s-bond. The p-bond in an alkene or alkyne is around þ250 kJ mol1 , while the s-bond is around þ350 kJ mol1 . Mean bond enthalpies (kJ mol–1) Mean bond lengths (pm) C C +347 153 C C +612 134 C C +838 120 Bond The shorter the bond length, the stronger the bond. For CH bonds, the greater the ‘s’ character of the carbon orbitals, the shorter the bond length. This is because the electrons are held closer to the nucleus. H3C CH2 H H2C CH H HC C H sp3 sp2 sp longest shortest A single CC s-bond can undergo free rotation at room temperature, but a p-bond prevents free rotation around a CC bond. For maximum orbital overlap in a p-bond, the two p-orbitals need to be parallel to one another. Any rotation around the CC bond will break the p-bond. 1.6 Inductive effects, hyperconjugation and mesomeric effects 1.6.1 Inductive effects In a covalent bond between two different atoms, the electrons in the s-bond are not shared equally. The electrons are attracted towards the most electronegative atom. An This compound contains four functional groups, including a phenol. Functional groups are introduced in Section 2.1 A hydrogen atom attached to a CC bond is more acidic than a hydrogen atom attached to a CC bond or a CC bond; this is explained by the change in hybridisation of the carbon atom that is bonded to the hydrogen atom (Section 1.7.4) Rotation about CC bonds is discussed in Section 3.2 6 Structure and bonding
1.6 Inductive effects,hyperconjugation and mesomeric effects arrow drawn above the line representing the covalent bond can show this (Sometimes an arrow is drawn on the line.Electrons are pulled in the direction of the arrow. polarisation of electrons through -bonds When the atomism eee09e,e8 electrons attracted toX electrons attracted toc remo en atom from an alkane(Section 2.2). pos -l groups +l groups X=Br,CI,NO2.OH.OR,SH, Z=R (alkyl or aryl). metals (e.a.Lior Ma】 al6heciRe5hoea23et9ea atomstrongeect Pauling electronegativity scale he nduoaid g-6--& The overall polarity of a molecule is determined by the individual bond polarities,formal charges and lone pair contributions and this can be measured by the dipole moment().The larger the dipole moment(often measured in debyes, D).the more polar the compound. 1.6.2 Hyperconjugation A o-bond can stabilise a neighbouring carbocation(or positively charged carbon. e.g.R3C)by donating electrons to the vacant p-orbital.The positive charge is delocalised or'spread out'and this stabilising effect is called resonance. a-bond The stability of carbocations is The electrons in the C-H discussed in Section 4.3. time in the mpty p-orb empty p-orbital 1.6.3 Mesomeric effects Whilst inductive effects pull electrons through the a-bond framework,electrons can also move through the t-bond network.A -bond can stabilise a negative charge,a
arrow drawn above the line representing the covalent bond can show this. (Sometimes an arrow is drawn on the line.) Electrons are pulled in the direction of the arrow. C X C Z positive inductive effect. +I δ+ δ– δ– δ+ electrons attracted to X electrons attracted to C When the atom (X) is more electronegative than carbon When the atom (Z) is less electronegative than carbon negative inductive effect. –I –I groups +I groups The more electronegative the atom (X), the stronger the –I effect The more electropositive the atom (Z), the stronger the +I effect X = Br, Cl, NO2, OH, OR, SH, Z = R (alkyl or aryl), SR, NH2, NHR, NR2, CN, CO2H, metals (e.g. Li or Mg) CHO, C(O)R The inductive effect of the atom rapidly diminishes as the chain length increases H3C CH2 CH2 CH2 Cl δδδ+ δδ+ δ+ δ– experiences a strong –I effect experiences a negligible –I effect K = 0.82 C = 2.55 N = 3.04 O = 3.44 I = 2.66 Br = 2.96 Cl = 3.16 F = 3.98 The higher the value the more electronegative the atom Pauling electronegativity scale The overall polarity of a molecule is determined by the individual bond polarities, formal charges and lone pair contributions and this can be measured by the dipole moment (m). The larger the dipole moment (often measured in debyes, D), the more polar the compound. 1.6.2 Hyperconjugation A s-bond can stabilise a neighbouring carbocation (or positively charged carbon, e.g. R3Cþ) by donating electrons to the vacant p-orbital. The positive charge is delocalised or ‘spread out’ and this stabilising effect is called resonance. C empty p-orbital C C–H σ-bond The electrons in the C–H σ-bond spend some of the time in the empty p-orbital H 1.6.3 Mesomeric effects Whilst inductive effects pull electrons through thes-bond framework, electrons can also move through the p-bond network. A p-bond can stabilise a negative charge, a An inductive effect is the polarisation of electrons through s-bonds An alkyl group (R) is formed by removing a hydrogen atom from an alkane (Section 2.2). An aryl group (Ar) is benzene (typically called phenyl, Ph) or a substituted benzene group (Section 2.2) Hyperconjugation is the donation of electrons from nearby CH or CC s-bonds The stability of carbocations is discussed in Section 4.3.1 1.6 Inductive effects, hyperconjugation and mesomeric effects 7