Alkenes - Introduction (Edexcel International A Level Chemistry): Revision Note

Alkenes - Introduction

• All alkenes contain a double carbon bond, which is shown as two lines between two of the carbon atoms i.e. C=C

• All alkenes contain a double carbon bond, which is the functional group and is what allows alkenes to react in ways that alkanes cannot

• Alkenes have the general molecular formula C_nH_2n

• They are said to be unsaturated hydrocarbons

  • They contain carbon-carbon double bonds

  • They are made up of hydrogen and carbon atoms only

• Alkenes are named using the nomenclature rule alk + ene

• In molecules with a straight chain of 4 or more carbon atoms, the position of the C=C double bond must be specified

• The carbon atoms on the straight chain must be numbered, starting with the end closest to the double bond

• The lowest-numbered carbon atom participating in the double bond is indicated just before the -ene:

The First Five Members of the Alkene Family

Bonding in Alkenes

• Each carbon atom has four electrons in its outer shell (electronic configuration: 1s²2s²2p²)

• Carbon atoms share these four electrons in four covalent bonds with other atoms to achieve a full outer shell configuration

• These electrons are found in orbitals within the respective atoms

• When forming a covalent bond, the orbitals overlap in such a way to form two types of bonds

  • Sigma bonds (σ)

  • Pi bonds (π)

• When carbon atoms use only three of their electron pairs to form a σ bond, each carbon atom will have a p orbital which contains one spare electron

• When the p orbitals of two carbon atoms overlap with each other, a π bond is formed (the π bond contains two electrons)

• The two orbitals that form the π bond lie above and below the plane of the two carbon atoms to maximise bond overlap

σ bonds

• Sigma (σ) bonds are formed from the end to end overlap of atomic orbitals

• s orbitals overlap this way as well as p orbitals

Sigma orbitals can be formed from the end to end overlap of s orbitals 

• The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond

• The pair of electrons is found between the nuclei of the two atoms

• The electrostatic attraction between the electrons and nuclei bonds the atoms to each other

Hydrogen

• The hydrogen atom has only one s orbital

• The s orbitals of the two hydrogen atoms will overlap to form a σ bond

π bonds

• Pi (π) bonds are formed from the sideways overlap of adjacent p orbitals

• The two lobes that make up the π bond lie above and below the plane of the σ bond

• This maximises overlap of the p orbitals

• A single π bond is drawn as two-electron clouds, one arising from each lobe of the p orbitals

• The two clouds of electrons in a π bond represent one bond containing two electrons

π orbitals can be formed from the sideways overlap of p orbitals

Ethene

• Each carbon atom uses three of its four electrons to form σ bonds

  • Two σ bonds are formed with the hydrogen atoms

  • One σ bond is formed with the other carbon atom

  • The fourth electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a π bond

  • This means that the C-C is a double bond: one σ and one π bond

Each carbon atom in ethene forms two sigma bonds with hydrogen atoms and one σ bond with another carbon atom. The fourth electron is used to form a π bond between the two carbon atoms