Lewis Structure, Molecular Geometry, Hybridization, Polarity, and the MO Diagram are all examples of C2H2 Lewis Structure.

Acetylene, or C2H2, is the simplest alkyne and a colourless hydrocarbon with a garlic-like odour. Due to the existence of two carbon atoms connected with a triple bond, it is extremely reactive to ambient temperature and lacks oxygen, making it an unsaturated molecule. Acetylene is very flammable and explosive because it is reactive, unstable, and lighter than air.

Despite its toxicity, acetylene is utilised in welding because it is combustible. To humans, this molecule poses a threat since its presence in the atmosphere limits the amount of oxygen available. Not only does it have an impact on humans, but it also has an impact on other living species by disrupting numerous natural atmospheric cycles in which oxygen plays an important role.

In light of this, the recommended airborne exposure limit (REL) for acetylene has been set at 2500 ppm (Ceiling), over which a quantity of acetylene that becomes an asphyxiant gas can kill humans.

As a result, it’s critical to comprehend acetylene’s behavioural chemical features in order to comprehend why it behaves the way it does. Let’s start with an examination of acetylene’s Lewis structure.

Acetylene’s Lewis Structure (C2H2)

The Lewis Structure is a diagram that depicts how the valence electrons participate in bond formation.

To investigate this, you must first understand the electronic setup of the participating elements. The atomic number of carbon (C) is 6, and its electronic configuration is 1s2 2s2 2p2.

Hydrogen (H), on the other hand, has an atomic number of 1 and an electronic configuration of 1s1.

Each atom seeks to establish a stable situation by stabilising the number of valence electrons, which is 8 for Carbon and 2 for Hydrogen, according to the octet rule. As a result, Carbon has four valence electrons, while Hydrogen has one.

What are valence electrons and what do they do?

The valence electrons are the electrons in an atom that are the furthest away from the nucleus. The valence electrons in an atom’s outermost two shells are involved in bond formation either by sharing or by entirely donating themselves.

Furthermore, their number within the shell is determined by the octet rule, which states that an atom is best stable when it has a maximum of 8 valence electrons.

Let’s take a look at the Lewis structure of acetylene one step at a time:

Step 1: Determine the total number of valence electrons that an acetylene molecule already has: A single acetylene (C2H2) molecule has a value of ten.

Step 2: Determine how many extra valence electrons an acetylene molecule requires: A single acetylene (C2H2) molecule has a value of ten.

Step 3: To begin drawing the structure, locate the central atom: There will be no centre atom because both elements (carbon and hydrogen) are present in equal proportions. This explains why the structure’s shape will be linear.

Step 4: Look for the following types of bonds forming between the atoms: The Carbon (C) atoms make a triple bond, while the Hydrogen (H) and Carbon (C) atoms form a single bond.

Assemble all of the previously given points and draw the following structure:

According to the octet rule, why is the number of valence electrons for Hydrogen and Carbon different?

The maximal capacity of a shell and its flexibility in exceptional situations are used to fill the number of valence electrons in the outermost shells.

This is because hydrogen can only hold two electrons while carbon can hold eight. It may surprise you to learn that certain elements, such as sulphur, do not follow the octet rule and can contain 10 to twelve valence electrons.

Acetylene’s Molecular Geometry (C2H2)

The study of a molecule’s molecular geometry is a crucial step in chemistry for analysing a molecule’s behavioural features.

It aids in the determination of a molecule’s polarity, phase of matter, magnetism, reactivity, colour, and biological activity; in short, molecular geometry may be used to study anything and everything about a molecule.

Acetylene (C2H2) is a tetraatomic molecule made up of two distinct atoms that link in equal amounts. Furthermore, carbon bonds to carbon, giving acetylene (C2H2) a linear structure and a 180° bond angle.

The Valence Shell Electron Pair Repulsion (VSEPR) theory states that the valence electrons surrounding an atom in a pair reject each other until they achieve an arrangement where this repulsion is minimised the greatest, which can be used to study the molecular geometry of acetylene (C2H2).

We’ll go deeper into this when we talk about the polarity of acetylene (C2H2) in the next section.

Polarity of C2H2 (Acetylene)

Polarity is a chemical property that allows materials to form poles that separate negative and positive charges.

With the help of a hydrogen bond, the molecule has a strong attraction and repulsion behaviour as a result of the charge separation. Non-polar molecules, unlike unstable polar molecules, are rather stable due to the absence of charge separation.

As a result, the molecule has a difficult time bonding with neighbouring atoms to form a new molecule. Non-polar molecules are made up of weak Van der Waal forces that aren’t as strong as hydrogen bonds, therefore new atoms don’t form strong bonds.

Polarity behaviour is exclusively determined by the electronegativity values of the involved atoms. Electronegativity is a property of atoms that governs how strongly they will attract electrons to complete their octet. The stronger the attraction, the higher the electronegativity value, and vice versa. Carbon (C) and Hydrogen (H) have electronegativity values of 2.55 and 2.20, respectively.

As a result, the electronegativity difference between Carbon-Carbon (C-C) and Carbon-Hydrogen (C-H) bonds is 0 and 0.35, respectively.

It can be seen that both values are less than 0.4, indicating that acetylene (C2H2) is non-polar, but the Carbon-Hydrogen (C-H) bond is somewhat more polar than the Carbon-Carbon (C-C) bond due to its larger value.

However, acetylene (C2H2) is non-polar in general because the electronegativity values are less than 0.4, cancelling out the net dipole moment totally.

You should also read the article about the polarity of C2H2 that was written particularly for you.

Hybridization of C2H2 (Acetylene)

The idea that the atomic orbitals merge and overlap to fuse and form hybrid orbitals, which directly impacts the molecular geometry and bonding behaviour of the newly produced molecule, is influenced by the concept of hybridization waves path for the molecular orbital diagram.

This can be investigated using the Valence Bond Theory (VBT), which states that bonding occurs in such a way that each molecule achieves a stable energy level with no severe repulsion.

In the acetylene (C2H2) molecule, the carbon atoms are sp hybridised, but the hydrogen atoms have unhybridized 1s atomic orbitals. In sp hybridization, the central atom’s s orbital only interacts with one of its p orbitals.

Atoms that display sp hybridization always have a linear molecular geometry with two sp orbitals at 180 degrees apart. As a result, linear molecular geometry is similar to sp hybridization in that figuring out one will aid with figuring out the other.

Both carbon atoms are linked by a triple connection composed by one sigma () bond and two pi () bonds, as seen in the Lewis structure.

This indicates that both carbon atoms have two sets of unhybridized p atomic orbitals that overlap to form two pi bonds between the sigma () bound sp-hybridized carbon atoms.

When the carbon achieves the excited state, one electron from the 2s orbital transfers to the 2pz orbital, resulting in one electron in each of the 2px 2py 2pz orbitals. Only the 2s1 and 2pz1 orbitals are hybridised in the case of a carbon-hydrogen bond.

This hybridization results in the development of four new sp hybridised orbitals, with two new sp hybridised orbitals produced by the carbon-hydrogen bonding.

Diagram of the Molecular Orbital (MO) of C2H2.

The Molecular Orbital (MO) Diagram depicts the bonding that occurs between the electrons of the participating atoms to create new molecules.

The atomic orbitals mix and overlap in a specific way to form a corresponding number of molecular orbitals, which is the primary idea this diagram follows.

This happens when electrons travel to different orbitals depending on their excitation level, causing them to be disseminated and redistributed within the orbitals that are involved.

Bonding and antibonding orbitals are formed when these electrons migrate from their original positions, giving rise to the molecular orbital diagram unique to each molecule.

The Carbon-Carbon bond is depicted in the above-mentioned molecular orbital diagram of acetylene (C2H2).

The sp hybridised orbitals join and overlap to generate a bonding sigma () orbital and an antibonding sigma (*) orbital, which can be seen.

Furthermore, the four p orbitals merge and overlap to generate two and two* orbitals, respectively. If we draw the energy sequence from the lowest to the highest molecular orbital, we get: σ < π(y) = π(z) < π(y)* = π(z)* < σ*.

The first and only bonding orbitals are filled since there are only 6 electrons available to fill the orbitals.


Acetylene (C2H2) is a dangerous chemical for humans because it depletes oxygen levels in the air.

The Lewis structure, which states that acetylene (C2H2) is an unsaturated chemical that is compatible and reactive enough to connect with air molecules and become harmful to human health, can be used to learn a lot about the molecule.

The molecular geometry, hybridization, polarity, and molecular orbital (MO) diagram can all help to explain this phenomenon.

Misha Khatri
Misha Khatri is an emeritus professor in the University of Notre Dame's Department of Chemistry and Biochemistry. He graduated from Northern Illinois University with a BSc in Chemistry and Mathematics and a PhD in Physical Analytical Chemistry from the University of Utah.


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