Lecture 1: Chemistry of the Carbonyl Group bjectives: By the end of this lecture you will be able to: 1. identify and name all major carbonyl functional groups; 2. use a molecular orbital approach to describe and account for the structure of a carbonyl group; 3. account for the electrophilicity of a carbonyl group. Important Carbonyl Compounds There are many types of carbonyl group. They are named according to their substitution patterns either side of the C= group. As we shall see later, they can be formed by oxidising alcohols. Secondary alcohols are oxidised to ketones; further oxidation - without breaking a C C bond - is not possible. Primary alcohols can be oxidised to aldehydes, which are at the same oxidation level as a ketone. However the presence of a hydrogen substituent means that further oxidation to a carboxylic acid is also possible. oxidation H H oxidation H ' H oxidation '
Carboxylic Acid Derivatives H N ' amidation ' NH 2 cat. H ' Cl ' H esterification ' H ' carboxylic acid cat. H Cl S Cl Cl Cyclic Esters and Amides N H (cyclic amide) (cyclic ester) Treating a Carboxylic Acid with Base H Base Na carboxylic acid e.g. NaH C 2 H NaH C 2 Na H 2
Structure of the Carbonyl Group using a Molecular rbital Approach Why does carbon always form four bonds and not two? Consider the ground state electron configuration of a carbon atom: To form a covalent bond requires the overlap of two orbitals each containing an unpaired electron. In its ground state, carbon has two unpaired electrons in two 2p atomic orbitals (As); thus in this state, it can form two covalent bonds. We can improve matters by taking the 2s and two of the 2p As and forming three sp 2 -hybridised atomic orbitals (HAs) which all have the same energy. ne 2p A is remaining. This so-called hybridisation requires energy. However, as orbitals of the same energy are filled singly first before pairing, even in this state we can still only form two covalent bonds as we still only have two unpaired electrons. If we provide additional energy that promotes one of the electrons out of an sp 2 -HA and into the 2p A then we now have a valence state in which we have four orbitals (three sp 2 HAs and one 2p A) that are all singly occupied. We can now form four covalent bonds. The energy required to form this valence state is more than compensated for by forming these two extra covalent bonds.
We can rationalise the structure of a carbonyl group by assuming the carbon and oxygen atoms making up the carbonyl group are both sp 2 -hybridised. Using this orbital approach, the carbon atom has three sp 2 -hybridised atomic orbitals (HAs) that are used to form three σ-bonds, and one p atomic orbital (A) to form a π-bond; likewise for the oxygen atom, only two of the sp 2 HAs are already filled with two electrons (remember oxygen has six electrons in its valence shell) and therefore do not form bonds, these are lone pairs. The p A on both atoms is orthogonal (i.e. at 90 ) to the plane containing the three sp 2 -HAs. The bonding structure of a carbonyl group can therefore be divided into its two constituent parts: the σ -framework (made from overlapping sp 2 -HAs) and the π-framework (made from overlapping p As). We shall consider each in turn and take acetaldehyde as an example: σ-framework: Formation of a σ-bond requires the direct overlap of two atomic or hybridised atomic orbitals. The C σ-bond is therefore formed by the direct overlap of one sp 2 -HA from carbon and one sp 2 - HA from oxygen. When these two HAs overlap (mix) they generate two new molecular orbitals (Ms) (Conservation of rbitals). This can be represented pictorially in a molecular orbital energy diagram:
There is one electron in each overlapping HA so the two new Ms are filled with the two electrons from these overlapping HAs. The lowest energy orbital is filled first (Hund's rule). Two electrons are required to fill an orbital (A, HA or M). Thus with only two electrons the lower lying σ-m (known as the bonding M) is completely filled, whilst the higher energy σ*-m (known as the antibonding M) remains empty. The effect of two singly occupied orbitals overlapping is to generate a system that is overall lower in energy - this is why a bond forms. Shape of σ- and σ*-ms:
The two remaining sp 2 -HAs on the carbon form similar σ-bonds to the atoms either side (i.e. the adjacent carbon and hydrogen atoms in the case of acetaldehyde). The two remaining sp 2 -HAs on the oxygen are each filled with two electrons (lone pairs). All the sp 2 -HAs and associated Ms are coplanar. The bond angle between any three atoms is approximately 120 π-framework: The p As on the carbon and oxygen are parallel to one another and orthogonal to the plane containing the σ-bonds. They too can overlap to form a bond. This is lateral overlap and results in the formation of a (weaker) π-bond. As before, two overlapping p As must form two Ms. The lower energy π-m is bonding, the higher energy π*-m is antibonding.
Shape of π- and π*-ms: NTE: there is zero probability of the electrons in the π-m existing in the plane containing the σ-bonding framework (this is known as a nodal plane). The electron density in the π-bond therefore lies above and below the plane containing the σ-bonds. The p As that form the C π-bond are higher in energy than the sp 2 -HAs that form the C σ-bond. Furthermore since π-ms are formed by lateral overlap of orbitals whereas σ-ms are formed by direct overlap, the net energy gain in bonding is lower when a π-bond is formed than when a σ-bond is formed. Consequently the resultant π-m is higher in energy that the σ-m. Just as there is a greater loss of energy in forming a σ-bond, the accompanying σ*-m is more 'antibonding' (higher in energy) than a π*-m and is therefore shifted to higher energy. This is best seen by overlaying the two molecular orbital energy diagrams:
E σ* π* C p A p A C sp 2 HA π sp 2 HA σ The Highest ccupied Molecular rbital (HM) is therefore the π (C=) -Molecular rbital. The Lowest Unoccupied Molecular rbital (LUM) is the π* (C=) -Molecular rbital. Polarisation of the C= Bond. Electronegativity of oxygen on the Pauling Scale = 3.44 Electronegativity of carbon on the Pauling Scale = 2.55 Consequently in the π-m (and the σ-m for that matter), there is more electron density on the oxygen than on the carbon. The bond is said to be polarised. Polarisation of the bond means that there is an uneven distribution of electron density between the two atoms leading to the build up of positive charge on the carbon end of the functional group with concomitant build up of negative charge on the oxygen end of the functional group.
Consequences of a Polarised Bond. 1. The carbonyl group exhibits a strong dipole moment. This has the effect of making the C= stretching band in I spectra very intense (see next lecture). 2. Although the C= bond is primarily a covalent bond, it also has some ionic character i.e. the atoms are also held together by electrostatic interactions. The net result is to make the C= group a thermodynamically more stable functional group than a C=C double bond. The formation of a C= group is a common thermodynamic driving force in chemical reactions. The C bond distance in a carbonyl group is 1.22 Å and shorter than the C C bond distance in a C=C double bond (1.32 Å).
3. CABNYL GUPS AE ELECTPHILIC i.e. they are susceptible to attack by NUCLEPHILES (species that contain high electron density). Since we know that the carbon atom has a δ+ charge on it and the oxygen a δ- charge, electron-rich nucleophiles will be attracted to the carbon atom and repelled by the oxygen atom on simple electrostatic grounds. NUCLEPHILES EACT WITH CABNYL CMPUNDS AT THE CABN CENTE Summary In this introductory lecture we have seen that there are a wide range of different functional groups containing a C= bond. There are many more! They all differ in the substituents either side of the C= group. These substituents have an important effect on the electrophilicity of the C= group (see later). We have also seen how a simple molecular orbital approach can be used to describe the planar structure of carbonyl compounds. The high electronegativity of oxygen compared to carbon means that the molecular orbitals that we derived by overlapping sp 2 -HAs (σ-ms) and p As (π-ms) are distorted such that more electron density lies on the oxygen end of the bond. Whilst both σ (C ) and π (C=) Ms are affected by this polarisation effect, the ramifications of this on the reactivity of carbonyl groups are most felt in the π-m system since this is where the HM and LUM are found.