UNIT TWO BOOKLET 1. Molecular Orbitals and Hybridisation

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1 DUNCANRIG SECONDARY ADVANCED HIGHER CHEMISTRY UNIT TWO BOOKLET 1 Molecular Orbitals and Hybridisation

2 In the inorganic unit we learned about atomic orbitals and how they could be used to write the electron configuration of atoms and ions. Molecular orbital theory is used to describe how atomic orbitals combine when atoms combine to produce molecules. It explains many issues with regard to chemical bonding such as molecular shape, the existence of double and triple bonds and why some organic molecules are coloured while others are not. Consider two hydrogen atoms approaching each other and attempting to bond. Molecular orbital theory states that when any two atomic orbitals meet each other and overlap they will form two new molecular orbitals (one called bonding and one called anti-bonding) in which a maximum of two electrons can be found. One of these molecular orbitals will be lower in energy than the two atomic orbitals from which it was made; the other molecular orbital lies at higher energy. The lower-energy MO of H 2 concentrates electron density between the two hydrogen nuclei and is called the bonding molecular orbital. This sausage shaped MO results from summing the two atomic orbitals so that the atomic orbitals combine in the region between the two nuclei. Because an electron in

3 this MO is attracted to both nuclei, the electron is more stable (it has lower energy) than it has in the 1s atomic orbital of an isolated hydrogen atom. Further, because this bonding MO concentrates electron density between the nuclei, it holds the atoms together in a covalent bond. The higher-energy MO has very little electron density between the nuclei and is called the antibonding molecular orbital. Instead of combining in the region between the nuclei, the atomic orbitals cancel each other in this region, leaving the greatest electron density on opposite sides of the nuclei. Thus, this MO excludes electrons from the very region in which a bond must be formed. An electron in this MO is repelled from the bonding region and is therefore less stable (it has higher energy) than it is in the 1s orbital of a hydrogen atom In both these types of molecular orbital the electron density is concentrated along the internuclear axis. This type of molecular orbital is called a sigma (σ) orbital. To distinguish between them an antibonding oribital is designated as sigma star (σ * ). We can draw energy diagrams showing how the atomic orbitals and molecular orbitals are related. The energy diagram for the molecular orbitals for hydrogen is shown below.

4 Similar to atomic orbitals, when placing electrons into molecular orbitals, the Aufbau principle, Hund s rule and the Pauli exclusion principle are obeyed. Obviously the picture gets more complicated when dealing with larger molecules with more and more electrons. The second-row atoms of the Periodic Table have valence 2s and 2p orbitals, and we need to consider how they interact to form MOs. Due to their shape and orientation, when p-orbitals are involved in forming bonds, they can overlap and therefore interact in two ways; END-ON and SIDE- ON. In general END-ON overlap affords a better contact between p-orbitals and so this leads to a lower energy (a more stable) molecular orbital than results from a side-on overlap. Orbital formed from side-on overlap are called pi(π) orbitals. In the diagram above the green p- atomic orbitals will form an end-on sigma molecular orbital while the others will form side-on pi molecular orbitals.

5 Molecular orbital diagrams for N 2, O 2, F 2 and Ne 2 N 2 Note there is one bonding orbital and one anti-bonding orbital for each atomic orbital. O 2 Bond Order this tells us how many bonds exist between the atoms. Bond order = ½(no of bonding electrons number of anti- bonding electrons ) F 2 Ne 2 The diagram for Ne2 conforms that this molecule is not allowed - Ne2 does not exist. Only the valence shell electrons are shown as these will be the electrons which are involved in bonding.

6 The picture on the right shows a representation of a methane molecule showing the familiar tetrahedral shape with the central carbon atom having a valency of four. Consider the electronic structure and the shape of the of the atomic orbitals in the valence shell of an isolated carbon atom. This picture is problematic. 1. How does carbon form four bonds with only 2 half- filled p-orbitals. 2. Why does methane have a tetrahedral geometry when none of the oribitals have this shape. In methane and all other alkanes these problems are solved by orbital hybridisation. This theory assumes that the 2s orbital and the three p orbitals mix or hybridise to form four new hybrid orbitals known as sp 3 orbitals. This can also be shown as an energy diagram. The hybrid orbitals are degenerate and have identical shapes. four sp 3 hybridised orbitals

7 As there is only a small energy gap between the 2s and 2p orbitals it is energetically favourable for carbon to promote a 2s electron to the empty 2p orbital. This Carbon will now have four unpaired electrons and so have a normal valency of four. It is energetically favourable because the energy required to promote one electron is more than compensated for when bonds are formed. If carbon forms four bonds instead of two then twice as much energy will be released. When orbitals overlap bonds are formed. In a similar way naming molecular orbitals, if the atomic orbitals overlap end on the bond is called a SIGMA(σ)BOND. This diagram shows an sp 3 hybrid orbital forming a sigma bond with the 1s orbital of a hydrogen atom In methane, all four hybrid orbitals are used to make 4 sigma bonds with hydrogen atoms. This gives the familiar tetrahedral arrangement for CH 4 molecules. All other alkanes bond in a similar way using sp 3 hydridisation. Sigma bonds form between carbon atoms as well as forming between hydrogen atoms and carbon atoms. The bonding in ethane is shown below. It is useful to remember that any carbon atom bonded to FOUR other groups is an sp 3 hybridised carbon atom.

8 sp3 hybridisation does not explain the bonding and structure of ethene. Ethene has bond angles of 120 degrees and all the atoms lie in the same plane it is definitely not tetrahedral. Furthermore it has a carbon to carbon double bond which, unlike single bonds, does not allow free rotation along the internuclear axis In alkenes the bonding observed is also due to hybridisation. As with alkanes, an electron from the 2s shell is promoted to the empty 2p orbital. This time the 2s orbital mixes with only TWO of the p orbitals forming THREE hybrid sp 2 orbitals. One of the p orbitals remains unhybridised. one unhybridised p orbital three sp 2 hybrid orbitals The carbon atom now has four electrons to use in bonding 3 at 120 degrees apart (on the same plane) in the sp 2 hybrid orbitals and one at 90 degrees to these in the unhybridised p orbital.

9 Have you ever thought it strange that when ethene reacts with bromine one of the carbon to carbon bonds breaks while the other does not. The reason for this is that the two bonds are NOT the same type of bond one of them is much stronger than the other. The diagram shows that when two sp 2 hybridised carbon atoms form a bond with each other the do so by end-on overlap a sigma bond the four hydrogen atoms also form sigma bonds with carbon the same way they did in ethane. Now consider the remaining bonding electron in the unhybridised p orbital. It is clear from the position they approach each other that they will overlap in a side-on fashion. When this happens a pi(π) bond will form pi bonds are weaker than sigma bonds Why? In simple terms, after forming a sigma-bond (a pre-requisite for pibonds), the two atoms get locked along the internuclear axis. As a result, the orbitals available for pi-bonding can only partially overlap, thus forming a weaker bond. The diagram below shows how sigma bonds form along the internuclear axis while pi bonds form above and below this axis. It is useful to remember that any double bond consists of one sigma and one pi bond.

10 Benzene, C 6 H 6, is an aromatic compound a compound which contains a ring of delocalised electrons. We now have to know exactly what this really means the answer lies in hybridisation. The benzene molecule is planar with carbon to carbon bond angles of 120 degrees. The carbon atoms are sp 2 hybridised and joined by sigma bonds. The hydrogen atoms are as usual joined to the carbon atoms by sigma bonds There are six unhybridised p orbitals (one per carbon atom) and these orbitals meet side on forming a doughnut shaped pi bonding system above and below the plane of carbon atoms. The six electrons will be somewhere in this pi system they are delocalised. This idea is shown in many different diagrams.

11 While many chemical compounds are coloured because they absorb visible light, most organic molecules appear colourless. Remember that colour in compounds generally arises due to the fact that electrons in the substance absorb certain wavelengths of light and move to higher energy levels. The observed colour of the compound is the complementary colour to the colour that is absorbed. Some organic molecules are coloured. Indeed, these molecules are often present in many of the coloured substances in nature. This molecule is responsible for the orange/ yellow colour of some fruits and vegetables Carrots, melons and peppers all contain carotene. H 3C CH3 H 3C CH CH 3 3 Vitamin A is a yellow coloured compound. This vitamin is required for good eyesight. Vitamin A deficiency can lead to an inability to see in the dark. Carotene is converted to vitamin A by the body and this is why there is some truth to the fact that carrots help you see in the dark. H 3C H 3C CH3 OH

12 Lycopene provides the red colour in fruit and vegetables. There is some evidence that lycopene can help prevent some forms of cancer. Absorption of visible light by organic molecules Why are these molecules coloured? Energy from photons (light) is used to promote electrons from bonding or non-bonding orbitals into higher energy anti-bonding orbitals. Several transitions are possible. The σ* and π* anti-bonding orbitals are normally empty. When absorptions occur, electrons are excited and promoted from a filled orbital (an electron in a σ or π bonding orbital or from a lone pair in a non-bonding orbital) into a higher energy anti-bonding orbital. Consider the transitions shown in the diagram (The diagram is not to scale). σ* π * Organic compounds that contain only σ bonds are colourless. The σ bonding orbital is the highest occupied molecular orbital (HOMO), and the lowest unoccupied molecular orbital (LUMO) is the σ* anti-bonding orbital. The transition between these orbitals (as shown above) is quite large (high energy) and corresponds to the UV part of the spectrum. Therefore no visible light is absorbed and the compound is colourless.

13 Excitations of electrons in compounds containing simple π bonds, like those shown below, still involve a large transition to promote an electron from HOMO (π bonding orbital) to LUMO (σ* anti-bonding orbital), and thus these compounds also absorb in the UV region of the spectrum. All these molecules are colourless. Consider the coloured molecules we looked at previously. H 3C OH H 3C CH3 H 3C All these molecules have a chain of alternating double(π) and single (σ) bonds. This is called CONJUGATION. The molecules shown above are conjugated. When this happens the electrons in the conjugated part of the molecules are DELOCALISED along the length of the conjugated chain. Delocalised electrons LOWER the energy gap between the HOMO and the LUMO to such an extent that compounds with enough conjugation will absorb light in the visible region of the spectrum as electrons are promoted from HOMO to LUMO. H 3C

14 The greater the number of atoms spanned by the delocalised electrons, the smaller the energy gap will be and the compound will absorb light of lower energy (towards the red end of the visible part of the spectrum) Compound Number of C=C in conjugated system Main colour absorbed Colour compound appears Vitamin A 5 Violet Yellow β-carotene 11 Blue Orange Lycopene 11 Green Red If a compound absorbs any portion of the spectrum in the visible light region, it will exhibit an observable colour. Since violet light has higher energy than blue or green, when it is absorbed we observe the yellow light that is transmitted. As molecules with greater conjugation absorb lower energy light, the greater the degree of conjugation, the more likely the compound is to have a red colour. Similarly, less conjugation would result in compounds appearing yellow. The part of the molecule responsible for causing the colour is termed the chromophore. Coloured compounds arise because visible light is absorbed by electrons in the chromophore, which are then promoted to a higher energy molecular orbital. The chromophore of vitamin A is highlighted in red.

15 Many ph indicators are coloured due to conjugated electron systems. Phenolphthalein is colourless in acid but pinky/purple in alkali. Why? + 2H + No conjugated system Big energy gap between HOMO LUMO Absorbs light in UV spectrum Colourless Conjugated system Lower energy gap between HOMO LUMO Absorbs light in Visible spectrum See complementary colour

16 1. A student was asked to draw a diagram to illustrate the bonding in ethene, C2H4. The diagram drawn by the student is shown below. Use your knowledge of chemistry to comment on this diagram. 2. How many sigma and how many pi bonds do the following molecules have? a. b. c. d. 3. The electronic spectra of molecules can be described in terms of the wavelength of maximum absorbance, max. The table below shows a number of compounds with their corresponding max values. a. Compound 1 is buta-1,3-diene. Name compound 2. b. Draw the most likely structure for the compound with max = 291nm. c. The compounds shown have a system of alternation double and single bonds. What word is used to describe this type of system? d. Explain why compound 4 has the highest max value.

17 e. - carotene, max = 452 nm gives the orange colour to carrots and has the structure whereas - carotene max = 434 nm is found in oranges and has the structure: Explain why there is a difference in the max values for these two structures. f. The pink colour of cooked salmon and lobster is due to astaxanthin which has the structure (i) (ii) Circle the chromophore in astaxanthin. The molecule is optically active. Circle any asymmetric carbon atom responsible for this optical activity. 4. The carbon atoms in benzene are sp 2 hybridised. a. Describe what is meant by sp 2 hybridised. b. Describe how a pi bond is formed. c. How many electrons are found in the delocalised pi system in benzene? 5. Why is there a conflict between the electronic configuration of carbon and the formula of carbon tetrachloride? 6. Ethanal and ethene both have sp 2 hybrid orbitals. a. Draw both molecules extended structural formulae and circle the sp 2 carbon atom in ethanal.

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