Hybridization of Methane

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CH₄ Hybridization

In Chemistry, hybridization is the process of combining atomic orbitals into new hybrid orbitals (with shapes, and energy different than the original atomic orbitals) appropriate for the pairing of electrons to form chemical bonds in valence bond theory. To understand the hybridization of methane (CH₄), we have to examine the atomic orbitals of distinct shapes and energy that are included in the hybridization process. In this article, we will explain how the CH₄ Hybridization occurs in detail, CH₄ shapes, CH₄ bond angles, the formation of CH₄, etc.

 

Name of the Molecule

Methane

Molecular formula

(CH₄)

Type of Hybridization

sp³

Bond Angle

109.50

Geometry

Tetrahedral

 

What is Hybridization of Methane?

In general, CH₄ is a combination of 1 carbon and 4 hydrogen atoms. However, to form this bond the central atom which includes 4 valence electrons obtains more electrons from 4 hydrogen atoms to complete its octet. The formation of covalent bonds gets more precise when the electrons are shared between carbon and hydrogen.

 

Now, if we talk about the hybridization of methane, the central carbon is sp³ hybridized. It is because three 2p orbitals and one 2s orbital in the valence shell of carbon combine to form four sp³ hybrid orbitals of carbon to form C-H sigma bonds which eventually leads to the formation of methane molecules.

 

Formation of Methane (CH₄)

Methane is an organic compound and is the most important component of natural gas. The structure of methane includes a central carbon atom with four single bonds to form hydrogen atoms. To maximize the distance from each other, the four groups of bonding atoms do not fall on the same plane. Alternatively, each carbon atom lies at the corners of a geometrical shape known as tetrahedral. The carbon atom lies in the middle of the tetrahedron. Each face of the tetrahedron is an equilateral triangle in shape. 

 

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The molecular geometry of methane is tetrahedral. The H-C-H bond angle of methane is 109.5 degrees and is greater than 90 degrees. While drawing the structural formula of methane, it is beneficial to represent the three-dimensional character of its shape. The structural formula of methane given below is a frame of reference. The dotted line bond is to be examined as moving back into the page while the solid triangle bond is to be examined as emerging out of the page.

 

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(CH₄) Shapes 

During the formation of  sp³ orbitals, they arrange themselves in such a way that they are as far away as possible from each other. This is known as a tetrahedral arrangement with a bond angle of 109.50

 

No changes can be seen in terms of shape when the hydrogen atom combines with the carbon atom, and so the methane molecule takes the shape of a tetrahedral with a  bond angle of 109.50. Hence, the CH₄ structure is tetrahedral.

 

CH₄ Bond Angles 

There are 4 pairs of outer electrons around the central atom in methane. These pairs of electrons repel each other.

 

The H-C-H bond angle in methane is the tetrahedral angle, 109.50. The angle is formed when all the four pairs of outer electrons repel each other equally. The bond angles in ammonia and in water are less than 109.50, due to the stronger repulsion by the lone pairs of electrons. Hence, the CH₄ bond angle is 109.50.

 

Important Points to Note

  1. Each sp³ hybrid orbital of carbon crossway 1s-orbital of hydrogen to form C-H sigma bonds.

  2. The hybridization contains the combination of 1 s orbital and 3 p orbitals and there are no lone pairs.

  3. The energy and shape of the sp³ hybrid orbitals are equal. They contain one unpaired electron each.

 

CH₄ Molecular Geometry and Bond Angles

Determining the CH₄ molecular geometry will be easier now as we have already discussed the bond formation and the process of hybridization above. In methane, the four hybrid orbitals are placed in such a way to minimize the force of repulsion between them. However, the four orbitals do repel each other and get disposed of at the corners of a tetrahedron. The shape of the CH₄ is tetrahedral. The sp³ hybrid orbital retains a bond angle of 109.50.

 

Valence Bond Theory

According to valence bond theory (VB), all bonds are localised bonds established between two atoms by the donation of one electron from each atom. This is an incorrect assumption since many atoms connect with delocalized electrons. The VB theory of molecular oxygen predicts that there are no unpaired electrons. The VB theory is effective in qualitatively describing the geometries of covalent compounds. Valence Bond theory covers the creation of covalent bonds as well as the electronic structure of molecules. 

 

The theory holds that electrons occupy the atomic orbitals of individual atoms inside a molecule and that electrons from one atom are attracted to the nucleus of another. As the atoms go closer, the attraction grows stronger until they reach a point where the electron density causes repulsion between them. The lowest potential energy is gained at the minimal distance between the two atoms, and this electron density may be thought to be what keeps the two atoms together in a chemical bond. The use of overlapping atomic orbitals in valence bond theory to describe how chemical bonds form works effectively in basic diatomic compounds like H2. When molecules with more than two atoms create stable bonds, a more comprehensive model is required. Methane is a nice example (CH4). 

 

Properties of Methane

  • Because methane is the simplest alkane as well as the simplest type of saturated hydrocarbons, it is critical to understand its characteristics. The following are some of the most important features of methane:

  • CH4 is the chemical formula for methane.

  • Methane gas has a specific gravity of 0.554, making it lighter than air.

  • The Molar Mass of methane is 16.04 g/mol.

  • It can only be dissolved in water.

  • When methane is burned, it produces a pale, luminous, and extremely hot flame.

  • Methane is one of the most significant greenhouse gases.

  • Because of its high energy density of 55.7 MJ/kg, pure methane is also utilised for home heating and cooking.

  • Methane has a boiling point of 161.50 0C. Methane has a melting point of 182.5 0C.

  • Methane is the conjugate acid of a methanide, which is one of its chemical characteristics.

  • Methane is a sossil fuel as well as a bacterial metabolite.

  • Methane can explode the container storing it and the rocker if exposed to fire or heat for an extended period of time.


Health Hazard of Methane

When breathed in large quantities, methane can have a negative impact on the human body. A high quantity of methane in enclosed spaces reduces oxygen levels, resulting in asphyxia, dizziness, headache, vomiting, loss of coordination, nausea, and loss of consciousness. If the quantity of methane in the air rises by 5 to 14 percent by volume, it becomes explosive. Explosions of this nature are common in coal mines and collieries. As a result, before entering the mines, fresh air is sent through to reduce the concentration of methane. Methane risks can occur during the manufacturing, usage, and transportation of methane. Although we absorb methane when we breathe, prolonged exposure to high levels of methane is hazardous.


Important Points to Remember in Hybridization

  • In isolated atoms, hybrid orbitals do not exist. Only covalently bound atoms may generate them.

  • The shapes and orientations of hybrid orbitals differ significantly from those of atomic orbitals in isolated atoms.

  • Combining atomic orbitals yields a set of hybrid orbitals. The number of hybrid orbitals in a set is equal to the number of atomic orbitals used to create the set.

  • The molecular geometry predicted by the VSEPR theory is created by the kind of hybrid orbitals produced in a bound atom.

  • In terms of form and energy, all orbitals in a set of hybrid orbitals are equivalent.

  • Lone pair electrons are frequently found in hybrid orbitals.

  • Bonds are formed when hybrid orbitals intersect.

FAQs (Frequently Asked Questions)

1. Explain the Chemical Properties of Methane.

Methane is a colourless, odourless gas that occurs largely in nature and as a product of human activities. Methane is the part of the paraffin series of hydrocarbons and it is among the strongest of the greenhouse gases. The chemical formula of methane is CH₄.


Chemical Properties of Methane  


Methane is much lighter than air, and it retains a specific gravity of 0.554. Methane is moderately soluble in water. The density of methane is 0.656 kg/m³. It can be easily burned in air forming carbon dioxide, and water. The flame of methane after burning is paler, very hot, and moderately luminous. The boiling and melting point of methane are -1620 C, and -182.50 C. The molecular weight of methane is 16.04 g/mol.

2. What are the uses of Methane?

Some of the uses of methane are as follows.

  • Methane is used in the generation of electricity.

  • It is used to sanitize different products.

  • It is used in gas cookers and gas-fired power stations.

  • The testing of gas appliances can be done easily using methane.

  • Methane in its refined liquid form can be used as rocket fuel.

  • Methane as a fuel can be used in automobiles, microwaves, and water heaters.

3. What is the impact of sp3 hybridization in methane?

The sp3 hybridization in the methane molecule causes a tetrahedral arrangement with a bond angle of 109.50. This hybridisation gives a stable structure to the molecule.

4. What is meant by orbital hybridization?

Orbital hybridisation (or simply called hybridization) is the mixing of two or more atomic orbitals to form a new hybrid orbital. That is, in the case of an impaired electron, chemical bonds can be created for bond valency. In the case of sp3 configuration, orbitals combine with three p orbitals into a tetrahedron – this is seen in the case of methane, ethane. Other examples of hybridization are sp2 as found in boron trichloride (BCl3), ethylene (C2H4). 

5. What is the configuration of valence electrons of carbon? 

2s22p2 is the orbital configuration of the valence electrons of carbon. 2s orbital, being lower in energy than 2p, will be filled first. That means that in the 2s orbital there are two electrons, and a single electron in two of the three 2p orbitals. There’s also an empty 2p orbital. This constitutes the configuration of the valence electrons of carbon.

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