Chemistry often seems like a realm of mystery, filled with curious interactions and elusive connections. Covalent bonding is one such mystery in the world of elements, which will hopefully be demystified by the end of this blog!
The Heart of Covalent Bonding
At a basic level, covalent bonding can be defined as the strong attraction between two atomic nuclei and their shared pair of electrons. For example, in HCl, Hydrogen and Chlorine both share one electron with each other, and this pair of electrons attract the two nuclei, binding them together.
An interesting question to ask now is, how are the two electrons shared between Hydrogen and Chlorine? Do they just sit in the middle?
If only it was that simple!
The way in which electron pairs are shared between atoms, comes down to a key concept in Chemistry called Electronegativity.
Simply put, electronegativity describes the ability of an atom/nucleus attract an electron. As you can imagine, if the nucleus has a higher positive charge, it will have a higher electronegativity, or if it has a lot of electron shells, it will have a lower electronegativity (since it would be more difficult to attract the outer electrons).
The way electrons are shared between the atoms in a covalent bond depends on the electronegativity difference between the atoms involved.
- If the difference is 0 (meaning both atoms attract the electrons equally), then the electrons will mostly be in the middle.
- If the electronegativity is non-zero but still small (typically less than 1.7), one atom will hold the shared pair of electrons closer to itself. This will create a partial charge as the atom which attracts the electrons more strongly will be slightly more negative than the other atom. This type of bonding is referred to as “polar covalent.”
- But what happens when the electronegativity difference becomes significant, typically greater than 1.7? The bond starts to become ionic.
This is a crucial point to grasp — covalent and ionic bonding aren’t distinct categories! They’re more like points on a continuum.
As the electronegativity difference increases, the bond becomes more ionic in character, with electrons transferring almost completely from one atom to another at large electronegativity differences
Strength of Covalent Bonding
The next question to think about is the bond strength of covalent bonding, and what affects them. The below is a list of factors you should be aware of:
- Number of Shared Electrons: The more electrons shared between atoms, the stronger the covalent bond. Double or triple bonds involve more shared electrons and are stronger than single bonds.
- Bond Length: A shorter bond length results in a stronger covalent bond because the atoms are closer together, so the shared electrons experience stronger attraction to both nuclei.
- Electronegativity Difference: Polar covalent bonds tend to be stronger than non-polar bonds for a few reasons. This include their partial ionic character which strengthens their attraction to each other, the shared pair of electrons being closer to a nucleus making their attraction stronger, and they typically have shorter bond lengths, which increases the bond strength, as mentioned before.
- Atomic Size: Smaller atoms tend to form stronger covalent bond due to the closer proximity of electrons to the nucleus.
- Bond Angle: The angle between atoms in a molecule can affect bond strength. Ideal bond angles lead to efficient overlap of orbitals, strengthening the bond.
- Molecular Shape: The shapes of molecules can impact bond strength. Bent or distorted shapes may weaken covalent bonds due to reduced orbital overlap.
- Hybridization: Hybridization of atomic orbitals can lead to the formation of stronger covalent bonds. Hybrid orbitals create more effective overlap, enhancing bond strength.
- Orbital Overlap: Effective overlap of atomic orbitals creates stronger covalent bonds. Sigma (σ) bonds involve direct head-on overlap, while pi (π) bonds involve sideways overlap (more on this later).
You should be aware of the factors that affect the strength of covalent bonding in order to identify trends in the periodic table, to have an idea of the strength of a given bond, etc.
Bond type: Sigma and Pi bonds
We mentioned earlier that covalent bonding arises as a result of sharing of an electron pair. But we should remember that electrons are not fixed particles orbiting an atom, as described in the Bohr model. Quantum mechanically, electrons have both particle and wave-like properties. So, it is more appropriate to visualise electrons as ‘clouds’ surrounding the nucleus (the electron is not within the cloud, but rather is a cloud – a cloud representing the probability distribution of where an electron can be at a point in time).
So, covalent bonding (sharing of electrons), is essentially an overlap of these electron “clouds”. Hence, we can imagine that the way in which two electron clouds overlap, will affect the strength of the resulting bond. This gives rise to two main types of overlap and bonding:
- Sigma (σ) bonds: The direct, head-on overlap of atomic orbitals
- Pi (π) bonds: Sideways overlap of atomic orbitals. Usually, the ‘second’ bond in a double bond is a pi-bond (and the ‘third’ bond).
Conclusion
So, there you have it!
The mesmerising story of covalent bonding, where electrons are shared, electronegativity dictates the nature of the bond, and bond length is a measure of strength.
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