BONDING IN METHANE AND ETHANE
Methane, CH4
The simple view of the bonding in methane
You will be familiar with drawing methane using dots and crosses
diagrams, but it is worth looking at its structure a bit more closely.
There is a serious mismatch between this structure and the modern
electronic structure of carbon, 1s2 2s2 2px1 2py1. The modern structure
shows that there are only 2 unpaired electrons for hydrogens to share
with, instead of the 4 which the simple view requires.
You can see this more readily using the electrons-in-boxes notation.
Only the 2-level electrons are shown. The 1s2 electrons are too deep
inside the atom to be involved in bonding. The only electrons directly
available for sharing are the 2p electrons. Why then isn't methane
CH2?
Promotion of an electron
When bonds are formed, energy is released and the system becomes
more stable. If carbon forms 4 bonds rather than 2, twice as much
energy is released and so the resulting molecule becomes even more
stable.
There is only a small energy gap between the 2s and 2p orbitals, and
so it pays the carbon to provide a small amount of energy to promote
an electron from the 2s to the empty 2p to give 4 unpaired electrons.
The extra energy released when the bonds form more than
compensates for the initial input.
Note: People sometimes worry that the promoted electron is drawn
as an up-arrow, whereas it started as a down-arrow. It simply makes
the diagram look tidier – nothing very sophisticated is going on!
Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane
all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of
orbitals. You aren't going to get four identical bonds unless you start from four identical orbitals.
Hybridization
The electrons rearrange themselves again in a process called
hybridization. This reorganizes the electrons into four identical hybrid
orbitals called sp3 hybrids (because they are made from one s orbital
and three p orbitals). You should read "sp3" as "s p three" - not as "s p
cubed".
Based on:
Clark, Jim. “chemguide. 2000. 6 Dec 2001 <http://www.chemguide.co.uk/>
Page 1 of 5
The sp3 hybrid orbitals look a bit like half a p orbital, and they arrange
themselves in space so that they are as far apart as possible. You can
picture the nucleus as being at the centre of a tetrahedron (a triangularly
based pyramid) with the orbitals pointing to the corners. For clarity, the
nucleus is drawn far larger than it really is.
What happens when the bonds are formed?
Remember that hydrogen's electron is in a 1s orbital – a spherically symmetric region of space
surrounding the nucleus where there is some fixed chance (say 95%) of finding the electron.
When a covalent bond is formed, the atomic orbitals (the orbitals in the individual atoms) merge
to produce a new molecular orbital which contains the electron pair which creates the bond.
Four molecular orbitals are formed, looking rather like the original sp3 hybrids, but with a
hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we've
previously drawn as a dot and a cross.
The principles involved - promotion of electrons if necessary, then hybridization, followed by the
formation of molecular orbitals – can be applied to any covalently-bounded molecule.
Ethane, C2H6
The formation of molecular orbitals in ethane
Ethane isn't particularly important in its own right, but is included because it is a simple example
of how a carbon-carbon single bond is formed.
Each carbon atom in the ethane promotes an electron and then forms sp3 hybrids exactly as
we've described in methane. So just before bonding, the atoms look like this:
Based on:
Clark, Jim. “chemguide. 2000. 6 Dec 2001 <http://www.chemguide.co.uk/>
Page 2 of 5
The hydrogens bond with the two carbons to produce molecular orbitals just as they did with
methane. The two carbon atoms bond by merging their remaining sp3 hybrid orbitals end-toend to make a new molecular orbital. The bond formed by this end-to-end overlap is called a
sigma bond. The bonds between the carbons and hydrogens are also sigma bonds.
In any sigma bond, the most likely place to find the pair of electrons is on a line between the two
nuclei.
Free rotation about the carbon-carbon single bond
The two ends of this molecule can spin quite freely about the sigma bond so that there are, in a
sense, an infinite number of possibilities for the shape of an ethane molecule. Some possible
shapes are:
In each case, the left hand CH3 group has been kept in a constant position so that you can see
the effect of spinning the right hand one.
Other alkanes
All other alkanes will be bonded in the same way:
• The carbon atoms will each promote an electron and then hybridise to give sp3 hybrid
orbitals.
• The carbon atoms will join to each other by forming sigma bonds by the end-to-end
overlap of their sp3 hybrid orbitals.
• Hydrogen atoms will join on wherever they are needed by overlapping their 1s1 orbitals
with sp3 hybrid orbitals on the carbon atoms.
Based on:
Clark, Jim. “chemguide. 2000. 6 Dec 2001 <http://www.chemguide.co.uk/>
Page 3 of 5
BONDING IN ETHENE
Ethene, C2H4
The simple view of the bonding in ethene
At a simple level, you will have drawn ethene showing two bonds between
the carbon atoms. Each line in this diagram represents one pair of shared
electrons.
Ethene is actually much more interesting than this.
An orbital view of the bonding in ethene
Ethene is built from hydrogen atoms (1s1) and carbon atoms (1s 2 2s2 2px1 2py1). The carbon
atom doesn't have enough unpaired electrons to form the required number of bonds, so it needs
to promote one of the 2s2 pair into the empty 2pz orbital. This is exactly the same as happens
whenever carbon forms bonds - whatever else it ends up joined to.
Now there's a difference, because each carbon is only joining to three other atoms rather than
four – as in methane or ethane. When the carbon atoms hybridize their outer orbitals before
forming bonds, this time they only hybridize three of the orbitals rather than all four. They use
the 2s electron and two of the 2p electrons, but leave the other 2p electron unchanged.
The new orbitals formed are called sp2 hybrids, because they are
made by an s orbital and two p orbitals reorganizing themselves.
sp2 orbitals look rather like sp3 orbitals that you have already come
across in the bonding in methane, except that they are shorter and
fatter. The three sp2 hybrid orbitals arrange themselves as far apart
as possible – which is at 120° to each other in a plane. The
remaining p orbital is at right angles to them.
The two carbon atoms and four hydrogen atoms would look like
this before they joined together:
The various atomic orbitals which are pointing towards each other now
merge to give molecular orbitals, each containing a bonding pair of
electrons. These are sigma bonds – just like those formed by end-to-end
overlap of atomic orbitals in, say, ethane.
Based on:
Clark, Jim. “chemguide. 2000. 6 Dec 2001 <http://www.chemguide.co.uk/>
Page 4 of 5
The p orbitals on each carbon aren't pointing towards each other, and so we'll leave those for a
moment. In the diagram, the black dots represent the nuclei of the atoms.
Notice that the p orbitals are so close that they are overlapping sideways.
This sideways overlap also creates a molecular orbital, but of a different
kind. In this one the electrons aren't held on the line between the two
nuclei, but above and below the plane of the molecule. A bond formed in
this way is called a pi bond.
For clarity, the sigma bonds are shown using lines – each line representing one pair of shared
electrons. The various sorts of line show the directions the bonds point in. An ordinary line
represents a bond in the plane of the screen (or the paper if you've printed it), a broken line is a
bond going back away from you, and a wedge shows a bond coming out towards you.
Note: The really interesting bond in ethene is the pi bond. In almost all cases where you will
draw the structure of ethene, the sigma bonds will be shown as lines.
Be clear about what a pi bond is. It is a region of space in which you can find the two electrons
which make up the bond. Those two electrons can live anywhere within that space. It would be
quite misleading to think of one living in the top and the other in the bottom.
All double bonds (whatever atoms they might be joining) will consist of a sigma bond and a pi
bond.
Based on:
Clark, Jim. “chemguide. 2000. 6 Dec 2001 <http://www.chemguide.co.uk/>
Page 5 of 5