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3l. The covalent bonding in the ethene molecule

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Part 3 Covalent Bonding: small molecules & properties

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Covalent bonding diagram for ETHENE covalent molecule, molecular formula C2H4

* metals \ non-metals (zig-zag line)

Pd metals Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0
1 1H  Note that hydrogen does not readily fit into any group but is a non-metal 2He
2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
The covalent molecule ethene from carbon combining with hydrogen


(c) doc b Two atoms of carbon (2.4) combine with four atoms of hydrogen (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.

The molecule can be shown as (c) doc b (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 120o). 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.

dot and cross diagram of the ethene molecule

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

  ball and stick model of ethene

space filling model of ethene



Melting point of ethene -169 oC

Boiling point of ethene -104 oC

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

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

From the quantum level rules, carbon's electron configuration is 1s22s22p2

In terms of the separate orbitals you can express it as 1s2, 2s2, 2px1, 2py1, 2pz0

This can be further expressed as an electron box diagram: 1s2s2p with two unpaired electrons.

However, as you should know by now, carbon usually forms four bonds (valency of 4) rather than two.

This is because it is energetically favourable to promote one electron from the 2s orbital into the third empty 2pz orbital, the energy required for this 'promotion' is far less than that released when the carbon atom forms four bonds rather than two..

This gives a theoretical electron configuration of 1s2, 2s1, 2px1, 2py1, 2pz1  or  1s2s2p

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.

alkenes structure and naming (c) doc b      diagram of sigma and pi bonds in alkenes covalent bonding in alkenes explained advanced level chemistry  Sigma and pi covalent bonds in alkenes

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 isomerism in alkenes)

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

What next?

Recommend next: The covalent bonding in the nitrogen molecule

Explaining the properties of small covalently bonded molecules

Sub-index for Part 3. Covalent Bonding: small molecules & properties

Index for ALL chemical bonding and structure notes

Perhaps of interest?

The basic chemistry of alkenes like ethene

Advanced A level notes on alkenes

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