4 Section A-Structure and bonding the combination of the 1s atomic orbitals.Since this is the more stable MO,the valence electrons (one from each hydrogen)enter this orbital and pair up.The antibonding MO is of higher energy and consists of two deformed spheres.This remains empty.Since the electron end up in a bonding MO which is more stable y is released and bond formation is e shall concentr ate solely on the bond- ing MOs to describe bonding and molecular shape,but it is important to realize that antibonding molecular rbitals also exist. Sigma bonds The bonding molecuar orital of gm a (o)bond:c This is a str of as rognlcaa rlap of t n ar a bonds a shall see of o bonds formed by the Hybridization Ato 11 ectron s such that t a carbon atom ectr ns and so we wo expec carbon to form two bond wever,carbon fo ms our bo nds!How does a carbon atom form four bonds th only two unpaired electrons So far,we have described the electronic configuration of an isolated carbon atom.However,when a carbon atom forms bonds and is part of a molecular struc ture,it can mix the s and p orbit als of its second shell(the valence shell).This is known as hybridization and it allows carbon to form the four bonds which we observe in reality. There are three ways in which this mixing process can take place. the2sorbitalismixed withallthree2porbitals.Thisisknownassphybridization the 2s orbital is mixed with two of the 2p orbitals.This is known as sp hybridization; the 2s orbital is mixed with one of the 2p orbitals.This is known as sp hybridization
4 Section A – Structure and bonding the combination of the 1s atomic orbitals. Since this is the more stable MO, the valence electrons (one from each hydrogen) enter this orbital and pair up. The antibonding MO is of higher energy and consists of two deformed spheres. This remains empty. Since the electrons end up in a bonding MO which is more stable than the original atomic orbitals, energy is released and bond formation is favored. In the subsequent discussions, we shall concentrate solely on the bonding MOs to describe bonding and molecular shape, but it is important to realize that antibonding molecular orbitals also exist. Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond: σ bonds have a circular cross-section and are formed by the head-on overlap of two atomic orbitals. This is a strong interaction and so sigma bonds are strong bonds. In future discussions, we shall see other examples of σ bonds formed by the interaction of atomic orbitals other than the 1s orbital. Hybridization Atoms can form bonds with each other by sharing unpaired electrons such that each bond contains two electrons. In Topic A1, we identified that a carbon atom has two unpaired electrons and so we would expect carbon to form two bonds. However, carbon forms four bonds! How does a carbon atom form four bonds with only two unpaired electrons? So far, we have described the electronic configuration of an isolated carbon atom. However, when a carbon atom forms bonds and is part of a molecular structure, it can ‘mix’ the s and p orbitals of its second shell (the valence shell). This is known as hybridization and it allows carbon to form the four bonds which we observe in reality. There are three ways in which this mixing process can take place. ● the 2sorbital is mixed with all three 2porbitals. This is known as sp3 hybridization; ● the 2s orbital is mixed with two of the 2p orbitals. This is known as sp2 hybridization; ● the 2s orbital is mixed with one of the 2p orbitals. This is known as sp hybridization
Section A-Structure and bonding A3 SP HYBRIDIZATION Key Notes Definition In sphybridization,thes and the p orbitals of the second shell are'mixed to form four hybridize ed sporbitals of equal energy. Electronic configuration Geometry 无心 ves apart from each o ther as possi ners of a tetrahedron.sp'Hybridization explains the tetrahedral carbon in saturated hydrocarbon structures. Sigma bonds Sigma(c)bonds are strong bonds formed between two sphybridized car bons or between an sp hybridized carbon and a hydrogen atom.A o bond formed between two sphybridized carbon atoms involves the overlap of half filled sphybridized orbitals from each carbon atom.Ao bond formed between an sp hybridized carbon and a hydrogen atom involves a half- filled sp'orbital from carbon and a half-filled 1s orbital from hydrogen. Nitrogen,oxygen,and chlorine atoms can also be sp'hybridized in organic molecules.This means that nitrogen has three half-filled sp'orbitals and can form three bonds which are pyramidal in shape.Oxygen has two half-filled sp'orbitals and can form two bonds which are angled with respect to each other.Chlorine has a single half-filled sp'orbital and can only form a single bond.All the bonds which are formed are o bonds. Related topics Covalent bonding and Bonds and hybridized centers(A6) hybridization(A2) Definition In sphybridization,the 2sorbital is mixed with all three of the2porbitals to givea set of four sp"hybrid orbitals.(The number of hybrid orbitals must equal the number of original atomic orbitals used for mixing.)The hybrid orbitals will each have the same energy but will be different in energy from the original atomic orbitals.That energy difference will reflect the mixing of the respect- ive atomic orbitals.The energy of each hybrid orbital is greater than the originals orbital but less than the original porbitals(Fig.1). Electronic The valence electrons for carbon can now be fitted into the sp'hybridized orbitals configuration (Fig.1).There was a total of four electrons in the original 2s and 2p orbitals.Thes orbital was filled and two of the p orbitals were half filled.After hybridization, there is a total of four hybridized sp'orbitals all of equal energy.By Hund's rule
Section A – Structure and bonding A3 SP 3 HYBRIDIZATION Definition In sp3 hybridization, the 2s orbital is mixed with all three of the 2p orbitals to give a set of four sp3 hybrid orbitals. (The number of hybrid orbitals must equal the number of original atomic orbitals used for mixing.) The hybrid orbitals will each have the same energy but will be different in energy from the original atomic orbitals. That energy difference will reflect the mixing of the respective atomic orbitals. The energy of each hybrid orbital is greater than the original s orbital but less than the original p orbitals (Fig. 1). Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitals configuration (Fig. 1). There was a total of four electrons in the original 2s and 2p orbitals. The s orbital was filled and two of the p orbitals were half filled. After hybridization, there is a total of four hybridized sp3 orbitals all of equal energy. By Hund’s rule, Key Notes In sp3 hybridization, the s and the p orbitals of the second shell are ‘mixed’ to form four hybridized sp3 orbitals of equal energy. Each hybridized orbital contains a single unpaired electron and so four bonds are possible. Each sp3 orbital is shaped like a deformed dumbbell with one lobe much larger than the other. The hybridized orbitals arrange themselves as far apart from each other as possible such that the major lobes point to the corners of a tetrahedron. sp3 Hybridization explains the tetrahedral carbon in saturated hydrocarbon structures. Sigma (σ) bonds are strong bonds formed between two sp3 hybridized carbons or between an sp3 hybridized carbon and a hydrogen atom. A σ bond formed between two sp3 hybridized carbon atoms involves the overlap of half filled sp3 hybridized orbitals from each carbon atom. A σ bond formed between an sp3 hybridized carbon and a hydrogen atom involves a half- filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen. Nitrogen, oxygen, and chlorine atoms can also be sp3 hybridized in organic molecules. This means that nitrogen has three half-filled sp3 orbitals and can form three bonds which are pyramidal in shape. Oxygen has two half-filled sp3 orbitals and can form two bonds which are angled with respect to each other. Chlorine has a single half-filled sp3 orbital and can only form a single bond. All the bonds which are formed are σ bonds. Related topics Covalent bonding and hybridization (A2) Bonds and hybridized centers (A6) Electronic configuration Nitrogen, oxygen and chlorine Sigma bonds Definition Geometry
6 Section A-Structure and bonding they are all half filled with electrons which means that there are four unpaired electrons.Four bonds are now possible. 小Energy Original atomic orbitals Fig.1.sp'Hybridization Geometry Each of the sphybridized orbitals has the same shape-a rather deformed looking dumbbell (Fig.2).This deformed dumbbell looks more like a p orbital than an s orbital since more p orbitals were involved in the mixing process. Minor lobe Major lobe Fig.2.sp'Hybridized orbital. Each sp'orbital will occupy a space as far apart from each other as possible by chof these lobes is 109.5.This is what is meant by the exp ssion tetrahedral carbon.The three nsional sha of the tetrahedral ca mal fo bonds in the plane of the ge.Bondepre nted by behind the wing a no pag a hatched wedge,and bo geaenteplane Tetrahedral shape 8npagegod Fig.3.Tetrahedral shape of an sphybridized carbon Sigma bonds A half-filled sp'hybridized orbital from one carbon atom can be used to form a bond with a half-filled sphybridized orbital from another carbon atom.In Fig.4a the maior lobes of the two st orbitals overlap directly leading to a strong o bond. It is the ability of hybridized orbitals to form c bonds that explains why hybridization takes place in the first place.The deformed dumbbell shapes allow a m ald be obtained froma ora re p and a hydro volves the of its ha-filledporbitals and the
they are all half filled with electrons which means that there are four unpaired electrons. Four bonds are now possible. Geometry Each of the sp3 hybridized orbitals has the same shape – a rather deformed looking dumbbell (Fig. 2). This deformed dumbbell looks more like a p orbital than an s orbital since more p orbitals were involved in the mixing process. Each sp3 orbital will occupy a space as far apart from each other as possible by pointing to the corners of a tetrahedron (Fig. 3). Here, only the major lobe of each hybridized orbital has been shown and the angle between each of these lobes is 109.5. This is what is meant by the expression tetrahedral carbon. The threedimensional shape of the tetrahedral carbon can be represented by drawing a normal line for bonds in the plane of the page. Bonds going behind the page are represented by a hatched wedge, and bonds coming out the page are represented by a solid wedge. Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form a bond with a half-filled sp3 hybridized orbital from another carbon atom. In Fig. 4a, the major lobes of the two sp3 orbitals overlap directly leading to a strong σ bond. It is the ability of hybridized orbitals to form strong σ bonds that explains why hybridization takes place in the first place. The deformed dumbbell shapes allow a much better orbital overlap than would be obtained from a pure s orbital or a pure p orbital. A σ bond between an sp3 hybridized carbon atom and a hydrogen atom involves the carbon atom using one of its half-filled sp3 orbitals and the hydrogen atom using its half-filled 1s orbital (Fig. 4b). 6 Section A – Structure and bonding 2py 2pz 2s 2px Original atomic orbitals Energy sp3 hybridized orbitals Fig. 1. sp3 Hybridization. Minor lobe Major lobe Fig. 2. sp3 Hybridized orbital. H H H C C H Bond going behind the page Bond coming out Tetrahedral shape of the page Bond in the plane of the page 109 o .5 Fig. 3. Tetrahedral shape of an sp3 hybridized carbon
A3-sp'Hybridization a) ①·①%c① 。3 sigma bond b) sigma bond 5限&nga6om2meao9包obard6 c nsp structures trogen valence ctrons in its second hybridization,it will have three half-fillec sp orbitals and can form three nds.Oxygen has six valence electrons.After hybridization,it will have two half-filled sp'orbitals and will form two bonds.Chlorine has seven valence electrons.After hybridization,it will have one half-filled sp"orbital and will form one bond. The four sp'orbitals for these three atoms form a tetrahedral arrangement with one or more of the orbitals occupied by a lone pair of electrons.Considering the atoms alone,nitrogen forms a pyramidal shape where the bond angles are slightly less than 109.5(c.107)(Fig.5a).This compression of the bond angles is due to the orbital containing the lone pair of electrons,which demands a slightly greater amount of space than a bond.Oxygen forms an angled or bent shape where two lone pairs of electrons compress the bond angle from 109.5 to c.104(Fig.5b). Alcohols,amines,alkyl halides,and ethers all contain sigma bonds involving nitrogen,oxygen,or chlorine.Bonds between these atoms and carbon are formed by the overlap of half-filled sp'hybridized orbitals from each atom.Bonds involv. ing hydrogen atoms(e.g.O-H and N-H)are formed by the overlap of the half- filled 1s orbital from hydrogen and a half-filled sp'orbital from oxygen or nitrogen. a CH Pyramidal H CH 1040 Angled molecule Fg.5.(a)Geometry of sphybridized nitrogen:(b)geometry of sphybridized oxygen
A3 – sp3 Hybridization 7 Nitrogen, oxygen, Nitrogen, oxygen and chlorine atoms can also be sp3 hybridized in organic and chlorine structures. Nitrogen has five valence electrons in its second shell. After hybridization, it will have three half-filled sp3 orbitals and can form three bonds. Oxygen has six valence electrons. After hybridization, it will have two half-filled sp3 orbitals and will form two bonds. Chlorine has seven valence electrons. After hybridization, it will have one half-filled sp3 orbital and will form one bond. The four sp3 orbitals for these three atoms form a tetrahedral arrangement with one or more of the orbitals occupied by a lone pair of electrons. Considering the atoms alone, nitrogen forms a pyramidal shape where the bond angles are slightly less than 109.5 (c. 107) (Fig. 5a). This compression of the bond angles is due to the orbital containing the lone pair of electrons, which demands a slightly greater amount of space than a bond. Oxygen forms an angled or bent shape where two lone pairs of electrons compress the bond angle from 109.5 to c. 104 (Fig. 5b). Alcohols, amines, alkyl halides, and ethers all contain sigma bonds involving nitrogen, oxygen, or chlorine. Bonds between these atoms and carbon are formed by the overlap of half-filled sp3 hybridized orbitals from each atom. Bonds involving hydrogen atoms (e.g. O–H and N–H) are formed by the overlap of the half- filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen or nitrogen. C + sp3 sp sigma bond 3 C C a) C Fig. 4. (a) σ Bond between two sp3 hybridized carbons; (b) σ bond between an sp3 hybridized carbon and hydrogen + sp3 C H 1s sigma bond H C b) N CH3 H H N CH3 H H 107o = Pyramidal a) Fig. 5. (a) Geometry of sp3 hybridized nitrogen; (b) geometry of sp3 hybridized oxygen. O H3C H H O CH3 104o Angled molecule = b)
Section A-Structure and bonding A4 SP2 HYBRIDIZATION Key Notes Definition nhybridization,ismixedwith two of thep orbitas to itals ofequ Geometry isshaped likea deformed dumbbell withone obe much larger than the other.The remaining 2p orbital is a symmetrical dumbbell The major lobes of the three sphybridized orbitals point to the corners of a triangle,with the 2p orbital perpendicular to the plane. Alkenes Each sphybridized carbon forms three o bonds using three sphybridized orbitals.The remaining 2p orbital overlaps'side on'with a neighboring 2p orbital to form a pi(n)bond.The t bond is weaker than the o bond,but is strong enough to prevent rotation of the C=C bond.Therefore,alkenes are planar,with each carbon being trigonal planar. Carbonyl groups The oxygen and carbon atoms are both sp'hybridized.The carbon has three sp'hybridized orbitals and can form three o bonds,one of which is to the oxygen.The oxygen has one sp'orbital which is used in the o bond with carbon.The porbitals on carbon and oxygen are used to form a r bond. Aromatic rings Aromatic rings are made up of six sp2 hybridized carbons.Each carbon forms three g bonds which re ults in a planar ring.The remaining 2p orbital on each carbon is perpendicular to the plane and can overlap with a neigh- boring 2p orbital on either side.This means that a molecular orbital is formed round the whole ring such that the six electrons are delocalized around the ring.This results in increased stability such that aromatic rings are less reactiv e than alkenes. Conjugated systems Conjugated systems such as cojuated alkenes 元bond are nds In such with the p lobes of a all lev el of do .This partia at ter on gives increas od s tability to Related topics Properties (2) ted aldehydes and ated dienes(H11) Structureandproperties(K1)
Section A – Structure and bonding A4 SP 2 HYBRIDIZATION Key Notes In sp2 hybridization, a 2s orbital is ‘mixed’ with two of the 2p orbitals to form three hybridized sp2 orbitals of equal energy. A single 2p orbital is left over which has a slightly higher energy than the hybridized orbitals. For carbon, each sp2 hybridized orbital contains a single unpaired electron. There is also a half-filled 2p orbital. Therefore, four bonds are possible. Each sp2 orbital is shaped like a deformed dumbbell with one lobe much larger than the other. The remaining 2p orbital is a symmetrical dumbbell. The major lobes of the three sp2 hybridized orbitals point to the corners of a triangle, with the 2p orbital perpendicular to the plane. Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridized orbitals. The remaining 2p orbital overlaps ‘side on’ with a neighboring 2p orbital to form a pi (π) bond. The π bond is weaker than the σ bond, but is strong enough to prevent rotation of the CC bond. Therefore, alkenes are planar, with each carbon being trigonal planar. The oxygen and carbon atoms are both sp2 hybridized. The carbon has three sp2 hybridized orbitals and can form three σ bonds, one of which is to the oxygen. The oxygen has one sp2 orbital which is used in the σ bond with carbon. The p orbitals on carbon and oxygen are used to form a π bond. Aromatic rings are made up of six sp2 hybridized carbons. Each carbon forms three σ bonds which results in a planar ring. The remaining 2p orbital on each carbon is perpendicular to the plane and can overlap with a neighboring 2p orbital on either side. This means that a molecular orbital is formed round the whole ring such that the six π electrons are delocalized around the ring. This results in increased stability such that aromatic rings are less reactive than alkenes. Conjugated systems such as conjugated alkenes and α,β-unsaturated carbonyl compounds involve alternating single and double bonds. In such systems, the p lobes of one π bond are able to overlap with the p lobes of a neighboring π bond, and thus give a small level of double bond character to the connecting bond. This partial delocalization gives increased stability to the conjugated system. Related topics Properties of alkenes and alkynes (H2) Conjugated dienes (H11) Aromaticity (I1) Properties (J2) α,β-Unsaturated aldehydes and ketones (J11) Structure and properties (K1) Electronic configuration Geometry Definition Alkenes Carbonyl groups Aromatic rings Conjugated systems