Multiple Bonds — Double & Triple Bonds

Multiple Bonds!

When it comes to covalent bonds, there are three types of bonds that can form: Single bonds, double bonds, and triple bonds

$\color{magenta}{Remember:}$ Single bonds are a result of two atoms sharing a pair of electrons

How are multiple bonds formed?

• Double bonds are a result of two atoms sharing two pairs of electrons
• Triple bonds are a result of two atoms sharing three pairs of electrons

Double Bonds

• Double bonds are stronger and shorter than single bonds. But, they are weaker and longer than triple bonds

Triple Bonds

• Triple bonds are the strongest and shortest covalent bonds!

Reminder:

The main group elements generally obey the octet rule. So, when forming bonds, atoms are limited by the valence number of they can share.

• Example: Oxygen is missing 2 electrons, thus it can only form two single bonds OR one double bond in order to complete its octet.

• Remember, there are exceptions to the octet rule.

(Video) Double and Triple Covalent Bonds

by The Science Classroom

This quick video will teach you about two special types of covalent bonds. These bonds form because almost all atoms want to follow the octet rule and have eight valence electrons. He shows you two examples, Oxygen and Nitrogen. Both are diatomic elements that bond to themselves. He also demonstrates how to find the valence electrons by using the groups of the periodic table. Finally, he also illustrates the structural formulas to demonstrate the lewis structures of multiple bonds.

Something important to keep in mind with these bonds is that as bond strength increases, bond length decreases.

Double and Triple Bonds in Covalent Compounds

Multiple bonds are made by the sharing of more than one electron pair. Double bonds are made of two electron pairs. Triple bonds are made of three electron pairs.

Multiple bonds are made of different types of bonds and orbitals. Sigma bonds are made from the combination of s-orbitals. Pi bonds are made from the combination of p-orbitals.

Quantum Mechanics and Mulitple Bonds

When we studied quantum mechanics, we learned that electrons reside in atomic orbitals. But, scientists had a problem. They had only described the orbitals of atoms. How could the quantum mechanical model explain bonding?

The first theory was valence bond theory. It was proposed in the 1930s by Linus Pauling. In valence bond theory, bonds are created when atomic orbitals overlap. A direct overlap between s-orbitals results in a sigma ($\sigma$) bond. The overlap of p-orbitals results in a pi ($\pi$) bond. So, we could explain molecules with multiple bonds. It also explains some of their chemical properties. For example, single bonds can rotate, but double and triple bonds are rigid. That's because the p-orbital overlap is lateral. So, the atoms are locked in place.

As a side note, valence bond theory also explains molecules like methane. In methane, carbon is bonded to four hydrogens. Carbon's electron configuration is $1s^{2}2s^{2}2p^{2}$. So, it has two valence electrons in the 2s and 2p shells. But, all the bonds in methane are equivalent. We know that the p-shell has more energy than the s-shell. So, the bonds shouldn't be equivalent. Thankfully, within valence bond theory, physical chemists developed hybrid orbitals. Hybrid orbitals occur in molecules where s and p orbitals are overlapping. The merging of s and p orbitals explains the equivalent bond energies. It also helps us to explain molecular geometry theories, like VSEPR. Molecules with double and triple bonds are locked into trigonal planar and linear geometries.

The second theory was molecular orbital theory. It describes a molecule's orbitals as the sum of the atomic orbitals. Its benefit is electrons can be delocalized throughout a molecule. For example, let's say we had a molecule with resonance. We don't have to consider multiple structures. Instead, we have a more intuitive understanding. The delocalized electrons are spread evenly throughout the whole molecule. This better explains a molecule's ionization energies and absorption and emission spectrum.