As with any scientific theory, hybridization must be able to explain the experimental evidence.

Hybridization can explain double bonds too, such as in ethene (C 2H 4).

These hybrids can be drawn as a mix of s and p orbitals, where they do still look similar to the p orbitals they have 75% of the character of, with the s character shown as one lobe of the p orbitals enlarged, the other shrunk (due to destructive/constructive overlap).

This explains the observation of equal C-H bond angles and bond strength in methane. Like in VSEPR, the electrons in different orbitals repel one another so they will find maximum spacing around the atom, which for tetrahedral carbon is 109.5°. Sigma bonds are formed by head-on orbital overlap.

These sp 3 hybrid orbitals can combine by overlapping with orbitals in other atoms (like the 1s AO of hydrogen) and form a sigma bond.

The character of the hybrids is related to the AOs that made them, so they are 25% ‘s character’ and 75% ‘p character’. In methane the one 2s orbital and three 2p orbitals create four sp 3 hybrid orbitals.

These four orbitals mix together to create four hybrid orbitals of equal energy.

This results in four singly occupied orbitals three 2p orbitals equal in energy and one 2s orbital slightly lower in energy.

This is done by a 2s electron being excited to a slightly higher 2p orbital, the one that is currently empty.

The four n=2 atomic orbitals (AOs) mix together to create four equivalent hybrid orbitals that are equal in energy. If they did, there would be marked differences in bond energy.

Atomic orbital hybridization explains why in methane (CH 4), carbon’s four valence electrons do not bond in this 2s 2 2p 2 ground state.

Like a hybrid animal or car, hybrid atomic orbitals are mixes of different atomic orbitals that show the characteristics of both. We call this idea atomic orbital hybridization. So how does carbon in CH 4 make four bonds? How are the C-H bonds all equal energy and length? One current theory of bonding is that atomic orbitals can mix together or hybridize to become hybrid orbitals that are equal in energy. Methane is a stable compound that is easily stored at standard conditions.Ī simple comparison shows that compounds where carbon makes four bonds instead of two are generally far more common and stable.

In methane (CH 4),carbon is bonded to four hydrogens and the bond angles and strengths are all equal.

However, methylene is very reactive and unstable – most simple carbenes are. Carbenes have two electrons on carbon that are not shared the other two are making covalent bonds.

Methylene (:CH 2) is a type of compound called a carbene.

So it has four valence electrons and two are in a higher energy state than the other two.ĭoes carbon actually make two bonds with a lone pair, like an analogue of NH 3? Let’s look at two carbon compounds: The two 2s electrons are paired up and two p electrons are unpaired in the higher energy 2p subshell. For example, carbon’s electron configuration is 2s 2 2p 2. We write electron configurations, like 1s 2 2s 2 2p 2, to show electrons in different atomic orbitals that create molecules out of separate atoms by taking part in bonding.īut atomic orbitals on their own don’t accurately reflect the bond energies and bond lengths in compounds.

In the lesson Atomic orbitals and energy levels, we learned that we use atomic orbitals (AOs) to explain bonding.To explain how atomic orbital hybridization is consistent with the observed properties of carbon-carbon bonds.To correctly draw sp 3, sp 2 and sp hybridized atomic orbitals.To explain the bonding in simple molecules using atomic orbital hybridization.

Atomic orbitals explained