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Advanced Higher Chemistry

Advanced Higher Chemistry. Unit 3 Bonding and reactions of alkanes. Alkanes. General Formula C n H 2n+2 Saturated hydrocarbons Each carbon atom forms 4 single bonds. Each carbon atom has 4 bonding pairs of electrons.

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Advanced Higher Chemistry

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  1. Advanced Higher Chemistry Unit 3 Bonding and reactions of alkanes

  2. Alkanes • General Formula CnH2n+2 • Saturated hydrocarbons • Each carbon atom forms 4 single bonds. • Each carbon atom has 4 bonding pairs of electrons. • The 4 bonding pairs of electrons repel one another resulting in a tetrahedral arrangement.

  3. Bonding in Alkanes Definitions :- Covalent Bond Two half filled atomic orbitals from two different atoms overlap and attract the nuclei of each atom. This gives rise to a MOLECULAR ORBITAL. Molecular Orbital Orbital of 2 bonding electrons moving under the influence of 2 nuclei.

  4. Electron arrangement of a ground state carbon atom 1s2 2s2 2p2 • How many unpaired electrons are there in a carbon atom (Remember Hund’s Rule)? Answer - There are 2 unpaired electrons in a carbon atom (the two electrons in the p orbital). Carbon should therefore only form 2 covalent bonds. • How many covalent bonds does carbon usually form? Answer - carbon usually forms 4 covalent bonds. WHY?

  5. Hund’s Rule of Maximum Multiplicity “When electrons occupy degenerate orbitals, the electrons fill each orbital singly, keeping their spins parallel before spin pairing occurs”

  6. Hybrid Orbitals A 2s electron is promoted to the third 2p orbital Carbon atom now has 4 singly occupied orbitals that ‘mix’ to form 4 hybrid orbitals of equal energy.

  7. 2s Orbital Three 2p orbitals (px, py and pz) Hybridisation Four sp3 hybrid orbitals

  8. Hybridisation is possible as the 2s and 2p sublevels are close in energy. • Hybridised orbitals are more directional in shape than unhybridised orbitals, this provides better overlap when bonds form. • In diagrams the small lobe of the hybridized orbital is usually omitted. • The sp3 hybridised carbon atom is often referred to as a tetrahedral carbon atom.

  9. Bonding in Alkanes • sp3 hybridised carbons form bonds when each of the four sp3 orbitals (each with one electron) overlap with an orbital of another atom that contains only one electron. e.g. methane

  10. Ethane • When orbitals from one atom overlap with the orbital from another atom a MOLECULAR ORBITAL is formed. • The four molecular orbitals in methane (the four C-H bonds) can be described as • sp3-s molecular orbitals

  11. A molecular orbital between two carbon atoms (a C-C bond) could be described as a sp3-sp3 molecular orbital • All bonds in alkanes are formed by this end-on overlap of orbitals. • All the molecular orbitals formed lie along the axis between the atoms (i.e. if you drew a line between the two atoms the molecular orbital would lie on that line). • When covalent bonds form along the axis between the atoms they are called sigma bonds (). • All covalent bonds in alkanes are sigma bonds.

  12. Reactions of Alkanes • Alkanes are relatively unreactive. e.g. paraffin (from Latin - unreactive) is a mixture of alkanes used to store alkali metals. • Only two types of reaction to consider 1) Combustion (alkanes are important fuels) e.g. natural gas petrol kerosene diesel 2) Reaction with chlorine and bromine.

  13. Halogenation • Reaction of bromine with alkanes is relatively slow, but it does happen. • Reaction will only start in the presence of sunlight. • Reaction is described as a substitution reaction or a free radical chain reaction. • A chain reaction consists of 3 steps. • Initiation • Propagation • Termination • Free Radical - An atom or molecule with an unpaired electron (e.g. Br•, CH3•, Cl•), always very reactive.

  14. Br—Br  Br• + •Br R—H •Br R• + H-Br Br—Br •R Br• + R-Br Step 2 - Propagation 1 : Reaction of Br radical with alkane Step 1- Initiation : Homolytic fission of bromine molecule.(energy supplied by UV light from sunlight) Propagation 2 : Reaction of alkane radical with Br2 • The new Br radical can now react with another alkane molecule, and so on and so on and so on. • Propagation will continue until one of the reactants is used up.

  15. Br• •Br  Br—Br R• •R  R—R OR R• •Br  R—Br OR Step 3 - Termination : When radicals meet.

  16. Notes on Chain Reactions • Chain reactions are very difficult to control and will produce a mixture of products. e.g. CH4 + Br2 CH3 Br + CH2 Br2 + CHBr3 + CBr4 • The more complicated the alkane the more complicated the mixture. • Controlling the proportions of reactants can give some control e.g. CH4 + excess Br2 Mainly CBr4 excess CH4 + Br2 Mainly CH3Br

  17. The ease of formation for halogen free radicals is related to bond strength. I-I Br-Br Cl-Cl  Stronger bond  Harder to form radicals  More reactive radicals • Bromination is slow, but chlorination can be very fast e.g. chlorination of methane when initiated using a flash gun .

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