Grade 11 → Chemical Bonding and Molecular Structure ↓
Hybridization and its types
The fascinating concept of hybridisation plays a vital role in understanding chemical bonding and molecular structure. Hybridisation is a theoretical model that describes the rearrangement of atomic orbitals to form new hybrid orbitals that are suitable for the pairing of electrons to form chemical bonds, usually covalent bonds. Let's delve deeper into the concept of hybridisation, explore its types and understand how it helps in predicting the geometry of molecules.
What is hybridization?
Hybridization is the process of combining atomic orbitals into new orbitals called hybrid orbitals. This concept was introduced to explain the structure of molecules such as methane (CH4
), which could not be explained satisfactorily using existing notions of atomic structure. Hybrid orbitals are degenerate, which means they have the same energy level, and they share characteristics of the constituent atomic orbitals.
Why is hybridization important?
Hybridisation helps to explain the following:
- Molecular geometry of compounds.
- Bond angles in molecules.
- Stability and strength of chemical bonds.
Basic principles of hybridization
There are several key principles about hybridisation that need to be understood:
- Hybrid orbitals are formed by the mixing of atomic orbitals located on the same atom.
- The number of hybrid orbitals formed is equal to the number of mixed atomic orbitals.
- The hybrid orbitals are oriented in such a way that they maintain maximum distance from each other, and minimize electron pair repulsion.
Types of hybridization
Hybridization can result in different geometries and types of bonds. The type of hybridization depends on the number and type of orbitals involved in the mixing. Here are the main types:
1. sp
hybridisation
In sp
hybridization, an s
orbital and a p
orbital combine to form two equivalent sp
hybrid orbitals. Each orbital is of 50% s
and 50% p
character. In this type the orbitals are oriented at 180 degrees relative to each other, forming a linear geometry.
Example: Acetylene (C2H2)
In acetylene, the carbon atoms are sp
hybridised. Each carbon atom forms two sp
hybridised orbitals. One of these forms a sigma bond with hydrogen and the other forms a sigma bond with the other carbon atom. The unhybridised p
orbitals overlap each other to form pi bonds, resulting in a triple bond between the carbon atoms.
C - H : sp ≡ H - C : sp
2. sp2
hybridisation
In sp2
hybridisation, one s
orbital mixes with two p
orbitals, resulting in the formation of three equivalent sp2
hybrid orbitals. These orbitals are arranged in a planar triangular configuration with an angle of 120° between them.
Example: Ethylene (C2H4)
In ethylene, each carbon atom is sp2
hybridized, and sp2
hybridized orbitals form sigma bonds with hydrogen and the other carbon. The remaining unhybridized p
orbitals overlap to form a pi bond, forming a double bond between the carbon atoms.
H - C = C - H sp2
3. sp3
hybridisation
In sp3
hybridization, one s
orbital mixes with three p
orbitals, forming four equivalent sp3
hybrid orbitals. These orbitals adopt a tetrahedral shape with a bond angle of 109.5°.
Example: methane (CH4)
In methane, carbon is sp3
hybridised, with four sp3
hybrid orbitals forming sigma bonds with hydrogen.
H - C - H sp3
4. sp3d
hybridisation
In sp3d
hybridization, one s
, three p
, and one d
orbital combine to form five sp3d
hybrid orbitals. These form a trigonal bipyramidal arrangement.
Example: Phosphorus pentachloride (PCl5)
In phosphorus pentachloride, phosphorus is sp3d
hybridized, and it forms five equivalent bonds with chlorine, forming a trigonal bipyramidal shape.
P / | Cl Cl Cl | | Cl Cl
5. sp3d2
hybridisation
In sp3d2
hybridization, one s
, three p
, and two d
orbitals combine to form six equivalent hybrid orbitals. This results in an octahedral arrangement.
Example: Sulfur hexafluoride (SF6)
In sulfur hexafluoride, sulfur is sp3d2
hybridized. It forms six bonds with fluorine in an octahedral shape, with 90° bond angles.
F | FSF | F | / F
How to determine hybridization?
Step 1: Count the valence electrons
Start with the number of valence electrons on the central atom. Consider the effect of surrounding atoms and any charge on the molecule or ion.
Step 2: Determine the molecular geometry
Use VSEPR (valence shell electron pair repulsion) theory to predict the geometry of the molecule. This helps to predict how many hybrid orbitals are needed.
Step 3: Use the hybridization shortcut
The regions of electron density around the central atom can help predict hybridization. Here are some guidelines:
- 2 area:
sp
- 3 region:
sp2
- 4 region:
sp3
- 5 region:
sp3d
- 6 region:
sp3d2
Hybridisation in different compounds
Water (H2O)
The structure of water is bent and its bond angle is 104.5°. Oxygen in water is sp3
hybridized, with two lone pairs and two single bonds with hydrogen.
H - O - H sp3
Ammonia (NH3)
The nitrogen atom in ammonia is sp3
hybridized. It forms three sigma bonds with hydrogen and retains a lone pair, resulting in a trigonal pyramidal shape.
H - N - H | H
Boron trifluoride (BF3)
Boron trifluoride shows sp2
hybridization, in which boron forms three sigma bonds with fluorine. This results in a planar triangular shape with a 120° bond angle.
F | BF | F
Importance of hybridization in chemistry
Hybridization is an important concept in understanding how molecules adopt certain shapes and bond types, which affect the physical and chemical properties of chemical reactions. It offered a coherent model that correlated with experimental data, helping to predict molecular structures and reactivity.
While hybridization is a valuable theoretical tool, it is important to understand that it oversimplifies the true quantum mechanical nature of electrons in atoms. Nevertheless, it remains a cornerstone in chemical education for explaining molecular geometry and bonding characteristics.
Conclusion
Hybridization laid the foundation for understanding how atoms combine to form molecules with specific geometries. By mixing atomic orbitals, hybridization explains the shapes, bond angles, and bond patterns of molecules, ultimately shaping the world of chemistry.