3l. Explaining the covalent bonding in the ethene molecule C₂H₄

Covalent bonding diagram for ETHENE covalent molecule, molecular formula C₂H₄

 Two atoms of carbon (electronic structure 2.4) combine with four atoms of hydrogen (electronic structure 1) to form ethene C₂H₄ (only the outer shell of carbon's electrons are shown).

On the right, the Lewis diagram of ethene - a simplified 'dot and cross' electronic diagram for the covalently bonded ethene molecule.

Electronically, hydrogen (1) becomes like helium (2) and carbon (2.4) becomes like neon (2.8), so ALL the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

Carbon is four electrons short of a full stable outer shell and hydrogen one.

The molecule can be shown as  (displayed formula) with one carbon = carbon double bond and four carbon – hydrogen single covalent bonds (it has a planar shape, its completely flat!, the H–C=C and H–C–H bond angles are 120^o). The valency of carbon is still 4.

The ethane molecule is held together by the four strong C–H carbon–hydrogen single covalent bonds and one C=C carbon–carbon covalent double bond.

 

The above diagrams are other styles of 'dot and x' electronic diagrams for ethene.

The Venn style diagram above on the right is the best style for the ethene molecule, clearly showing the sharing of the pairs of electrons for the C=C double covalent bond, and the pairs of electrons for the C-H single bond in the ethene molecule (in a sort of Venn diagram style).

  AND 

ball and stick model of ethene  AND  space filling model of ethene

See also advanced A-level chemistry notes on the shapes of molecules

Comments on ethene, a member of the alkene series of hydrocarbons

Melting point of ethene -169 ^oC

Boiling point of ethene -104 ^oC

You would expect low values because of the weak intermolecular forces between small covalent molecules like ethene.

A colourless gas at room temperature, with a strong hydrocarbon odour.

Use to manufacture many chemicals including the plastic polymer poly(ethene).

GCSE level links for alkenes like ethene

ALKENES - unsaturated hydrocarbons - reaction with bromine & hydrogen

CRACKING - a problem of supply and demand, other products

Addition polymers: poly(ethene), poly(propene), polystyrene, PVC, PTFE, structure & uses

More on the uses of plastics, issues with using plastics, solutions & recycling methods

Advanced level link

The chemistry of alkenes

See also advanced A-level chemistry notes on the shapes of molecules

Extra advanced level notes on the structure and covalent bonding of ethene

This gives four unpaired electrons, all of which can pair up with an electron from another atom to form four covalent bonds, but they may be of two varieties if a double or triple bond is involved in the carbon based molecule - read on ...

Reminders (if needed): The diagrams below show the bonding situation in alkenes, that conveniently involve two of the most important types of covalent bond between atoms, including carbon.

In carbon based compounds a single bond (sigma bond, σ bond) is formed by the overlap of two orbitals which can be either an s orbital and p orbital, (illustrated above).

Two electrons (an electron pair) are mutually attracted to the positive nuclei on either side.

A molecular orbital is formed and the axis of the bonding orbital lies on a central line between the two nuclei.

So we have a 'central' C-C sigma bond shown by the black line and the other atoms C, H or Cl etc. are also bonded to the C=C carbon atoms with sigma bonds too i.e. a planar >C-C< sigma bond system of single covalent bonds.

So, we now have to account for the 4th valency electron of the C=C carbon atoms.

In the case of alkenes, the pi bond (π bond) is formed by the overlap of two 2p orbitals of the carbon atoms, but, due to repulsion with the bonded pairs of the sigma bond electrons, they cannot form another sigma bond molecular orbital along the same central axis.

Instead, two pi orbital (in yellow, containing one electron each), are formed above and below the planar arrangement of the sigma bond linking the two carbon atoms (and other C or H atoms), giving the planar >C=C< bond arrangement.

Incidentally, the presence of the pi orbitals inhibits rotation around the double bond in unsaturated alkenes, because it requires a lot of energy to twist the orbitals around and break the pi bond and this accounts for the possibility of E/Z 'geometrical' isomerism in alkenes)

However, in alkanes everything is free to rotate around the sigma bonds of the C-C bonds in saturated alkane molecules.