In chemistry, a triple bond is defined as a covalent linkage, where two atoms share three pairs of electrons, as in the nitrogen molecule, acetylene, C2H2, or N2. One of the main electron pairs exists in a sigma bond, which is concentrated in the region along the line joining the two nuclei; the remaining two pairs are available in pi bonds, each of which occupies the two parallel regions of space on the opposite sides of the line that is determined by the two atoms.
The triple bond symbol is ☰.
Bonding
The bonding types may be explained in terms of orbital hybridization. In the acetylene case, every carbon atom contains two p-orbitals and two sp-orbitals. The two p-orbitals are perpendicular on both the y-axis and z-axis. At the same time, the two sp-orbitals lie linear with 180° angles and occupy the x-axis (cartesian coordinate system). When the carbon atoms approach each other, the sp orbitals overlap to make an sp-sp sigma bond. The Pz-orbitals, at the same time, approach, and together, they produce a Pz-Pz pi-bond. Similarly, the other pair of Py-orbitals produce a Py-Py pi-bond. The result is the formation of one sigma-bond and two pi-bonds.
Example of a Triple Bond Formation
Formation of Triple bond in Carbon
One of the perfect triple bond examples representing the value of hybrid orbitals is Carbon. The ground state configuration of Carbon is given as follows:
As per the Valence Bond Theory, carbon should produce two covalent bonds by resulting in a CH2 due to the reason it has two unpaired electrons in its electronic configuration. However, some experiments have already shown that CH2 is highly reactive and it cannot exist outside of a reaction. Thus, this does not explain how the CH4 can exist. To produce four bonds, the carbon’s configuration must contain four unpaired electrons.
CH4 can be explained in one way, one 2s and three 2p orbitals combine to form four equal energy sp3 hybrid orbitals. That would give the configuration as follows:
Now that the carbon compound has four unpaired electrons so that it can have four equal energy bonds. The orbital hybridization is favored due to the hybridized orbitals being more directional that leads to greater overlap when forming the bonds. Thus the bonds formed are stronger. This results in very stable compounds when hybridization takes place.
Let us see the explanation of various types of hybridization further and how every type helps explain the structure of certain molecules.
sp3 Hybridization
sp3 hybridization may explain the tetrahedral structure of the molecules. In it, both the 2s orbitals and all the three of 2p orbitals hybridize to make four sp3 orbitals, each consisting of 75% of p character and 25% of s character. And, the frontal lobes align themselves in a manner as represented below. In this particular structure, electron repulsion is minimized.
Energy changes occurring in the hybridization is given below:
The s-orbital’s hybridization with all the three p orbitals (Px, Py, and Pz) results in four sp3 hybrid orbitals. And, the sp3 hybrid orbitals are oriented at a bond angle of 109.5° from each other. This particular 109.5° arrangement gives the tetrahedral geometry as figured above.
Stability of Triple Bond Compared to a Single Bond
In the case of covalent molecules, more is the sharing of electrons between the atoms; stronger is: a single bond 2 electrons are shared, in a manner, 4 in double bond and 6 in a triple bond. Therefore, a triple bond is the strongest and most difficult to break. Now, the stronger the bond between the two atoms, the stabler (or more stable) the molecule. Thus, a triple bond is said to be more stable.
In terms of strength, the order is Triple >> Double >> Single,
In terms of bond length, the order is Single >> Double >> Triple,
In terms of stability, However, it is Single >> Double >> Triple.
Why? Because mostly, it’s much harder to break the sigma bonds compared to break the pi bonds, and single bonds contain zero pi bonds, whereas the triple bonds have two of them, which makes them easier to break.
Rotation of a Triple Bond
Yes, the rotation of a triple bond takes place. Only a single p-orbital is involved in the sp hybridization in acetylene. The other two mutually perpendicular P-orbitals, say Pz, Py can rotate, replacing each other’s axis (on rotation, Pz becomes Py and Py becomes Pz). However, practically it is meaningless and merely detected because there is no formation of a new compound and the groups attached to do carbon just to rotate about their axis, bringing no change in the overall symmetry and geometry. (The angle of rotation should be 90°** or its multiples)
Example
Let us take an example of 2-butyne. Rotating the triple bond raises many questions. First, it is not said to be meaningful. Raman or Infrared spectroscoPy are the two methods to notice the rotations. The end methyl groups’ rotations are visible. So, What visible is the stretching of a triple bond. If you grab the two methyl groups on a speculative paper experiment, the triple bond could rotate, but the main thing is how you would see it.