Presentation on theme: "Molecular Orbitals and Hybridisation Organic Chemistry."— Presentation transcript:
Molecular Orbitals and Hybridisation Organic Chemistry
Molecular orbitals 1s atomic orbitals of hydrogen H Orbitals can be used to explain bonding between atoms. Atomic orbitals are the volume of space that the electrons of an atom are likely to be found in. H The atomic orbitals containing the valence electrons (outer electrons) are the ones that are important to us.
This molecular orbital is more stable than each of the separate atomic orbitals and contributes to the shape of the molecule. H H H In the case of hydrogen, the overlap of two 1s atomic orbitals results in the formation of a σ (sigma) molecular orbital. When atomic orbitals overlap, they combine to form molecular orbitals. H 1s atomic orbitals of hydrogen σ molecular orbital σ bonds are covalent bonds formed between atoms when end-on overlap of orbitals occurs.
The molecular orbital formed is a lower energy arrangement than the separate atomic orbitals. Increasing energy H H H 1s σ H Molecular orbitals encompass the whole molecule and are not simply found between atoms inside a molecule.
Bonding continuum The shape of the molecular orbital formed from overlapping atomic orbitals will govern the type of intermolecular bonding that is observed.
Non-polar (pure) covalent bonds Non-polar covalent bonds or pure covalent bonds are formed between two atoms of the same element, or two atoms with a very low difference in electronegativity. The molecular orbital formed from overlapping atomic orbitals is symmetrical around a mid-point where the bonding electrons are most likely to be found.
Non-polar (pure) covalent bonds Example: fluorine The overlap of two 2p orbitals results in the formation of a σ orbital. F F+ 2p atomic orbital σ molecular orbital
Non-polar (pure) covalent bonds Example: fluorine The overlap of two 2p orbitals results in the formation of a σ orbital. F F+ 2p atomic orbital σ molecular orbital FF In a fluorine molecule, or any non-polar covalent bond, the σ bonding orbital is symmetrical.
Polar covalent bonds When there is a large difference between the electronegativities of the two elements involved in the bond, the bonding molecular orbital will be asymmetrical. Example: water O H H δ–δ– δ+δ+ δ+δ+ Since oxygen is more electronegative than hydrogen, the molecular orbital formed will be asymmetrical, with the bonding electrons more likely to be found around the δ– oxygen atom.
Ionic bonds When ionic bonds form, there is extreme asymmetry and the bonding molecular orbital is almost entirely around one atom.
Hybridisation In its ground state, an isolated atom of carbon has the electron arrangement 1s 2 2s 2 2p 2. C H H H H Why then, if there are only two unpaired electrons, do carbon atoms form four covalent bonds? 1s2s2p
Hybridisation 1s2s2p The shapes of the atomic orbitals involved cannot explain the bonding observed in compounds such as alkanes. x z y x z y x z y x z y x z y 1s orbital 2s orbital 2p x orbital2p y orbital2p z orbital
Hybrid theory assumes that the 2s and 2p orbitals of carbon atoms combine (or mix) to form four degenerate orbitals (i.e. orbitals of equal energy) Increasing energy 2s 2p hybridised orbitals The hybrid orbitals formed from one s orbital and three p orbitals are called sp 3 orbitals.
Since electrons repel each other, the four sp 3 hybridised orbitals surrounding a central carbon atom result in a familiar tetrahedral shape, with a maximum possible angle between each orbital of 109.5°. an sp 3 hybridised orbital The sp 3 orbitals formed are all half-filled, with the electron far more likely to be found in the larger lobe.
In methane, all four hybrid orbitals are used to form σ bonds between the central carbon atom and hydrogen atoms. Alkanes C H H H H
Carbon-to-carbon single bonds in alkanes result from overlapping sp 3 orbitals forming σ bonds. Alkanes H H H C C H H H σ bond σ bonds are covalent bonds formed by end-on overlap of two atomic orbitals and since σ bonds must lie along the line joining both atoms, there will be free rotation around these orbitals.
How can we explain the existence of double bonds as observed in alkenes? Alkenes As with alkanes, bonding in alkenes is due to hybridisation. C H H H H H H H H H H
As with alkanes, an electron from the 2s shell is promoted to the empty 2p orbital. This results in the formation of three hybrid orbitals, with one remaining unhybridised 2p orbital. Increasing energy 2s 2p hybridised orbitals The hybrid orbitals formed from one s orbital and two p orbitals are called sp 2 orbitals. single unhybridised 2p orbital
The three sp 2 orbitals repel each other, resulting in a bond angle of 120° between them. sp 2 orbitals The hybrid orbitals are responsible for overlapping to form σ bonds joining their central carbon atoms to both carbon and hydrogen.
The unhybridised p orbitals are perpendicular to the plane of the molecule. The p orbitals of the carbon atoms are parallel and close enough to overlap sideways. σ bonds σ bond unhybridised 2p orbitals
This sideways overlap between the 2p orbitals produces a new molecular orbital between the two carbon atoms. This new orbital is called a pi (π) orbital or more commonly a π bond. A π bond is a covalent bond formed by the sideways overlap of two parallel atomic orbitals.
Looking at information comparing σ and π bonds, we can see that double bonds are stronger than single bonds, but not twice as strong. This is because the sideways overlap (π bond) is weaker than the end-on overlap (σ bond). σ and π bonds Bond type Bonding orbitals present Bond length Mean bond enthalpy 1 σ154 pm370 kJ mol –1 1 σ + 1 π134 pm602 kJ mol –1 1 σ + 2 π121 pm835 kJ mol –1 C
Aromatic compounds differ to other hydrocarbons as they contain delocalised electrons. Aromatic hydrocarbons Example: benzene (C 6 H 6 ) Chemists initially represented a molecule of benzene as shown here. C C C C C C H H H H H H However, contrary to what might be expected from this structure, benzene is a very stable, saturated structure that does not undergo addition reactions.
This model does not explain why all the bonds in benzene can be observed to be the same length, not three longer single bonds and three shorter double bonds. In benzene, each carbon atom has used three of its four valence electrons to form σ bonds. The fourth electron of each carbon atom is delocalised over the entire ring, not involved in π bonding. C C C C C C H H H H H H
The σ bonding can be described as existing between six sp 2 hybridised orbitals. There are six C–C σ bonds in the molecule and so each carbon atom has two σ bonds to adjacent carbon atoms. Every carbon atom also has a σ bond to a hydrogen.
This results in a planar molecule with the unfilled 2p orbital of each carbon atom above and below the plane of the molecule.
These 2p orbitals all combine to form a set of delocalised π molecular orbitals above and below the plane of the molecule. The structure of benzene is drawn as shown to represent the delocalised electron clouds.
A substituted benzene ring is called a phenyl group (C 6 H 5 ) and can be represented: Many medicines, antiseptics, drugs and other useful products contain aromatic rings. R Trinitrotoluene (TNT) CH 3 NO 2 O2NO2N Trichlorophenol (TCP) OH Cl