The mixing of one s – orbital and one p – orbital of the same atom of nearly same energy to form a set of diagonally arranged two identical hybrid orbitals of equivalent energy is called sp hybridization. These hybrid orbitals are arranged in a linear manner around central atom and are at an angle of 180^o to one another.

Formation of BeF₂ Molecule:

Ground State of Beryllium Atom:

Atomic number of beryllium is 4. Its configuration in ground state is 1s², 2s², 2p⁰.

Beryllium atom in ground state:

Excited state of Beryllium Atom:

During combination with fluorine, the 2s electron pair is split up and one electron is promoted to empty 2p_x orbital. This condition is called excited state of beryllium. In excited state one electron of 2s migrates to 2p orbital forming 2 – half filled orbitals.
Beryllium atom in excited state:

Hybridization of Beryllium Atom:

One 2s orbital and one 2p orbitals of beryllium mix up forming two hybrid orbitals of equivalent energy. These two new equivalent orbitals are called sp hybrid orbitals. They are identical in all respect.

Beryllium atom in excited state:

Angle and Geometry :

Two sp hybridized orbitals formed, repel each other and These hybrid orbitals are arranged in a linear manner around central beryllium atom and are at an angle of 180^o to one another. Each sp hybrid orbital contain unpaired electron. In each sp hybrid orbital, one of the lobes is bigger because of more concentration of electron density. Only bigger lobe is involved in bond formation.

Thus in BeF₂ molecule has linear or diagonal structure with boron atom at the centre and two fluorine atoms at the either sides of beryllium. F-Be-F bond angle is 180^o.

Bond:

Two sp hybrid orbitals of boron atom having one unpaired electron each overlap separately with 1p orbitals of two fluorine atoms along the axis forming two covalent bonds (sigma bonds). The bonds between beryllium and fluorine are sp- p overlap. Thus F – Be — F bond angles are 180^o. Molecule is diagonal or linear. Both Be-F bonds in beryllium difluoride are of equal strength.

Diagram:

Type and Geometry of Beryllium difluoride Molecule:

BeF₂ Is linear molecule, while H₂O is angular molecule.

In BeF₂ molecule, beryllium undergoes sp hybridization and achieve electronic configuration 1s², 2s¹2p_x¹ in excited state. During formation of beryllium difluoride boron undergoes sp hybridization One 2s orbital and one 2p orbital of beryllium mix up forming two hybrid orbitals of equivalent energy. These two new equivalent orbitals are called sp hybrid orbitals. They are identical in all respect. Two sp hybridized orbitals formed, repel each other and they are directed diagonally opposite in space.  Angle between them is 180⁰. Two sp hybrid orbitals of boron atom having one unpaired electron each overlap separately with 1p orbitals of two fluorine atoms along the axis forming two covalent bonds (sigma bonds). Thus BeF₂ Is linear molecule.

In water, the oxygen atom is sp³ hybridized. Two hybrid orbitals have paired electrons (lone pair) and they are non – bonding orbitals.  Other two orbitals are half – filled (singly occupied) and they are bonding orbitals. These hybridized orbitals are in four directions of four corners of tetrahedron. Four sp³ hybridized orbitals formed, repel each other and they should be directed towards the four corners of a regular  tetrahedron and  Angle between them should be 109.5⁰.  The non bonding electron repel each other strongly and occupy more space than the electron pairs involved in bonding.  The force of repulsion between electron pair decreases in following order. lone pair – lone pair > lone pair- bond pair >  bond pair – bond pair. Hence the bond angle between two lone pairs i.e. H-O-H angle decreases from109.5⁰ to 104.5⁰. Thus water molecule is angular.

Formation of acetylene (C₂H ₂) Molecule:

Ground State of Carbon Atom:
Atomic number of carbon is 6. Its configuration in ground state is 1s², 2s², 2p² i.e. 1s² 2s², 2p_x¹ 2p_y¹ 2p_z⁰

Carbon atom in ground state:

Excited state of Carbon Atom:

During combination with hydrogen, the 2s electron pair is split up and one electron is promoted to empty 2p_z orbital. This condition is called excited state of carbon. In excited state one electron of 2s migrates to 2p orbital forming 4 – half filled orbitals.
Carbon atom in excited state:

Hybridization:

In acetylene there is sp hybridization. One 2s orbital and one 2p orbitals of carbon mix up forming two hybrid orbitals of equivalent energy. These two new equivalent orbitals are called sp hybrid orbitals. They are identical in all respect. Two ‘p’ orbitals of each carbon atom remains unhybridized.

Angle and Geometry:

Two sp hybridized orbitals formed, repel each other and These hybrid orbitals are arranged along x- axis in a linear manner around central carbon atom and are at an angle of 180⁰ to one another. The unhybridized p_y and p_z orbital remain perpendicular to hybrid orbitals along y-axis and z- axis ,
mutually perpendicular. Each sp hybrid orbital and unhybrid orbitals contain unpaired electron. In each sp hybrid orbital, one of the lobes is bigger because of more concentration of electron density. Only bigger lobe is involved in bond formation.

As all the four atoms in C₂H₂ molecule being in the same line, the molecule is diagonal or linear. H-C-C bond angle is 180^o.

Formation of Bonds:

1. Sigma Bond Formation :

A covalent bond formed by collinear or coaxial or in the line of internuclear axis. Overlapping of orbitals is known as sigma bond. One Sp hybrid orbital of one carbon atom overlaps with One hybrid orbital of other carbon atom by head on collision forming sigma bond. One (Sp- Sp ) overlap. Remaining one hybrid orbitals of each carbon atom overlap with ‘s’ orbital of two hydrogen atoms separately forming two sigma bonds. (2 C – H). Two (Sp– s)overlaps. Both C-H bond in Acetylene are of equal strength. Thus there are Three sigma bonds. Sigma bonds are stronger.

2) Formation of pi Bond :
The covalent bond formed by collateral or sidewise overlapping is called pi bond. The unhybridized 2 p_x and 2 p_z orbitals of each carbon atom being perpendicular to each other and to the plane of H-C-C-H axis overlap laterally with one another to form two week pi bond between two
carbon atoms by two p – p overlap. (one 2 p_y -2 p_y )and (one 2p_z-2 p_z) . Thus two (p-p) -overlaps.

In ethylene molecule there are 3 sigma bonds and 2 pi bonds. There is a triple bond between carbon and carbon consisting one sigma and two pi bonds.

Diagram:

Type and Geometry of Acetylene Molecule:

There is only one pi bond in ethylene molecule but there are two pi bonds in the acetylene molecule:

In ethylene molecule carbon atom shows sp² hybridization in excited state. The resulting three hybrid orbitals form two C-H bonds and One pi bond of sigma type. Thus each carbon atom is left with unhybridized p_z orbitals with lobes above and below the plane of hybridized orbitals. These two unhybridized orbitals overlap each other laterally and form a single pi bond between two carbon atoms.

In acetylene molecule carbon atom shows sp hybridization in excited state. The resulting two orbitals are linear and form one C-H bond and one C-C bond of sigma type. Thus each carbon is left with two unhybridized p_y and p_z orbitals, which are mutually perpendicular to H-C-C-H axis. These unhybridized orbitals overlap each other laterally and form two pi bonds between two carbon atoms.

The post SP Hybridization appeared first on The Fact Factor.

Mixing of one ‘s’ orbital and two ‘p’ – orbitals of nearly same energy forming set of trigonally arranged three identical orbitals of equal energy is known as sp² hybridization.

Geometry sp² Hybridization:

The hybrid orbitals are arranged around the central atom in a plane at an angle of 120° to one another. When these three orbitals overlap with appropriate orbitals of three other atoms, three bonds are formed and the resulting molecule has a trigonal planar structure. Each bond angle is 120°.

Diagram:

Formation of Boron Trifluoride Molecules:

Ground State of Boron Atom:

Atomic number of boron is 5. Its configuration in ground state is 1s², 2s², 2p¹

Boron atom in ground state:

Excited State of Boron Atom:

During combination with fluorine, the 2s electron pair is split up and one electron is promoted to empty 2p_y orbital. This condition is called excited state of boron. In excited state one electron of 2s migrates to 2p orbital forming 3 – half filled orbitals.

Boron atom in excited state:

Hybridization of Boron Atom:
One 2s orbital and two 2p orbitals of boron mix up forming three hybrid orbitals of equivalent energy. These three new equivalent orbitals are called sp² hybrid orbitals. They are identical in all respect

Boron atom in hybridized state:

Angle and Geometry:

• Three sp² hybridized orbitals formed, repel each other and they are directed towards the three corners of an equilateral triangle. The angle between them is 120°.
• Each sp² hybrid orbital contains an unpaired electron.
• In each sp² hybrid orbital, one of the lobes is bigger because of more concentration of electron density. Only bigger lobe is involved in bond formation.
• Thus BF₃ molecule has a trigonal planar structure with boron atom at the centre and three fluorine atoms at the four corners of equilateral triangle. F-B-F bond angle is 120°.

Bond Formation:

Three sp² hybrid orbitals of boron atom having one unpaired electron each overlap separately with 2p orbitals of three fluorine atoms along the axis forming three covalent bonds (sigma). Thus in BF₃ molecule has a planar structure with a boron atom at the centre and three fluorine atoms at the three corners of an equilateral triangle. F-B-F bond angle is 120°.

Bond:

• There are three sigma bonds..
• The bonds between boron and fluorine are sp²– p.
• Thus F – B — F bond angles are 120°. The molecule is trigonal planar.
• All B-F bonds in boron trifluoride are of equal strength.

Diagram:

Type and Geometry of Boron trifluoride Molecule:


The BF₃ molecule has a planar structure while NH₃ molecule has a pyramidal structure.
During formation of boron trifluoride, boron undergoes sp² hybridization and achieve the electronic configuration 1s², 2s¹2p_x¹2p_y¹. In the excited state, one 2s orbital and two 2p orbitals of boron mix up forming three hybrid orbitals of equivalent energy. These three new equivalent orbitals are called sp² hybrid orbitals. They are identical in all respect. Three sp² hybridized orbitals formed, repel each other and they are directed towards the three corners of an equilateral triangle (in one plane). Angle between them is 120^o. Three sp² hybrid orbitals of boron atom having one unpaired electron each overlap separately with 1p orbitals of three fluorine atoms along the axis forming three covalent bonds (sigma bonds). Thus boron trifluoride has triangular planar structure.

During formation of ammonia nitrogen undergoes sp³ hybridization. One 2s orbital and three 2p orbitals of mix up forming four hybrid orbitals of equivalent energy. These four new equivalent orbitals are called sp³ hybrid orbitals. They are identical in all respect and they should be directed towards the four corners of a regular tetrahedron and Angle between them should be 109.5^o. Three hybridized orbitals contain unpaired electron. The fourth hybridized orbital has lone pair of electron. The three half filled (containing unpaired electron) sp³ hybrid orbitals of nitrogen overlap axially with three half filled 1s orbitals of three hydrogen atoms separately to form three covalent N-H bonds (sigma bonds). The fourth hybrid orbital containing lone pair of electron remains non bonded. The non bonding electron repel each other strongly and occupy more space than the electron pairs involved in bonding. The force of repulsion between lone pair- bond pair is greater than bond pair – bond pair. Hence the
bond angle i.e. H-N-H angle decreases from109.5^o to 107^o.

Formation of Ethylene Molecule:

Ground State of Carbon Atom:
Atomic number of carbon is 6. Its configuration in ground state is 1s², 2s², 2p² i.e. 1s² 2s², 2p_x¹ 2p_y¹ 2p_z⁰

Carbon atom in ground state:

Excited state of Carbon Atom:

During combination with hydrogen, the 2s electron pair is split up and one electron is promoted to empty 2p_z orbital. This condition is called excited state of carbon. In excited state one electron of 2s migrates to 2p orbital forming 4 – half filled orbitals.
Carbon atom in excited state:

Hybridization of Carbon Atom:

In ethylene there is sp² hybridization. One 2s orbital and two 2p orbitals of carbon mix up forming three hybrid orbitals of equivalent energy. These three new equivalent orbitals are called sp² hybrid orbitals. They are identical in all respect. One ‘p’ orbital of each carbon atom remains unhybridized

Carbon atom in hybridized state:

Angle and Geometry:

Three sp² hybridized orbitals formed, repel each other and they are directed in a plane towards the three corners of an equilateral triangle. Angle between them is 120^o. The unhybridized p_z orbital remain perpendicular to this plane. Each sp² hybrid orbital contain unpaired electron. In each sp² hybrid orbital, one of the lobes is bigger because of more concentration of electron density. Only bigger lobe is involved in bond formation.

As all the six atoms in C₂H₆ molecule being in the same plane, the molecule is planar. H-C-H bond angle is 120^o. H-C-C bond angle is 120^o.

Bonds Formation:
1. Sigma Bond Formation :

A covalent bond formed by collinear or coaxial or in the line of internuclear axis. Overlapping of orbitals is known as sigma bond. One Sp² hybrid orbital of one carbon atom overlaps with One hybrid orbital of other carbon atom by head on collision forming sigma bond. One (Sp²– Sp² ) overlap.

Remaining two hybrid orbitals of each carbon atom overlap with ‘s’ orbital of four hydrogen atoms separately forming four sigma bonds. (C – H). Four (Sp²– s) overlaps. All C-H bond in ethylene are of equal strength.

Thus there are five sigma bonds. Sigma bonds are stronger.

2. Formation of pi Bond:

The covalent bond formed by collateral or sidewise overlapping is called pi bond. The unhybridized 2 p_z orbitals of each carbon atom being perpendicular to the plane of four hydrogen atoms and carbon atoms overlap laterally with one another to form a week pi bond between two carbon atoms by p – p overlap. One (p-p) -pi bond. This bond consists of two equal electron cloud one lying above the plane of the atom and other lying below this plane.

Hence, in ethylene molecule there are 5 sigma bonds and 1 pi bond.

Diagram :

Type and Geometry of Ethylene Molecule :

The post SP2 Hybridization appeared first on The Fact Factor.

The mixing of one s-orbital and three p- orbitals of the same atom having nearly the same energy to form four orbitals of equal in all respects and tetrahedrally arranged is known as sp3 hybridization.

Geometry of sp3 Hybridization:

sp³ hybridized orbitals repel each other and they are directed to four corners of a regular tetrahedron. The angle between them is 109.5° and the geometry of the molecule is tetrahedral (non-planar). This type of hybridization is also known as tetrahedral hybridization. The sp³ hybridization is shown pictorially in the figure.

Formation of Methane Molecule ( CH₄ ):

Step -1: Formation of the excited state of a Carbon atom:

The carbon atom in the ground state takes up some energy and goes to the excited state.  In this process, a pair of electrons in 2s orbital splits up and one of the electron from this pair is transferred to empty 2pz orbital.  Thus the excited state has four half-filled orbitals.

Step – 2: Hybridization of Orbitals:

One orbital of 2s and three orbitals of 2p mix up forming four hybrid orbitals of equivalent energy. These four new equivalent orbitals are called sp³ hybrid orbitals. They are identical in all respect.

Angle and Geometry:

Four sp³ hybridized orbitals formed, repel each other and they are directed towards the four corners of a regular tetrahedron.    The Angle between them is 109.5°. Each sp³ hybrid orbital contains one unpaired electron.

In each sp³ hybrid orbital, one of the lobes is bigger because of more concentration of electron density. Only a bigger lobe is involved in bond formation. Due to this maximum overlapping is achieved.

Four sp³ hybrid orbitals of carbon atom having one unpaired electron each overlap separately with 1s orbitals of four hydrogen atom along the axis forming four covalent bonds. Thus in CH₄ molecule has a tetrahedral structure with a carbon atom at the centre and four hydrogens at the four corners of a regular tetrahedron. H-C-H bond angle is 109.5°.

Bonds:

Four sp³ hybrid orbitals of carbon atom having one unpaired electron each overlap separately with 1s orbitals of four hydrogen atom along the axis forming four covalent bonds (sigma bonds). The bonds between carbon and hydrogen are sp³– s. Thus H – C – H bond angles are 109.5°. The molecule is tetrahedral. All C-H bonds in methane are of equal strength.

Type and Geometry of Methane Molecule:

Formation of Ammonia Molecule ( NH₃):

Need of Hybridization in Ammonia Molecule:

In nitrogen atom, there are three half-filled 2p orbitals and the valency should be 3 and it is three. These three ‘p’ orbitals are perpendicular to each other and if they form bonds, the angle between them should be 90°. But the actual angle between bonding orbitals in ammonia is 107°. To explain this difference in the angle the concept of hybridization is required.

Electron Configuration:

The Atomic number of Nitrogen is 7.  Its configuration in ground state is 1s², 2s², 2p³

Hybridization:

In the formation of ammonia one 2s orbital and three 2p orbitals of nitrogen mix up forming four hybrid orbitals of equivalent energy. These four new equivalent orbitals are called sp³ hybrid orbitals. They are identical in all respect. One hybrid orbital has paired electrons (lone pair) and it is nonbonding orbital. The other three orbitals are half-filled and they are bonding orbitals. The nonbonding pair of hybridized orbitals is called as a lone pair. These hybridized orbitals are in the directions of four corners of a regular tetrahedron.

Angle and Geometry:

Four sp³ hybridized orbitals formed, repel each other and they should be directed towards the four corners of a regular tetrahedron and the angle between them should be 109.5°. One hybrid orbital has paired electrons and it is nonbonding orbital. The other three orbitals are half-filled and they are bonding orbitals. But due to greater repulsion exerted by lone pair on bonding orbitals, the angle is reduced to 107°. Thus ammonia has distorted tetrahedral shape.

The three half-filled (containing unpaired electron) sp3 hybrid orbitals of nitrogen overlap axially with three half-filled 1s orbitals of three hydrogen atoms separately to form three covalent N-H bonds Hence geometry is tetrahedral pyramidal in which the nitrogen lies at the centre, three hydrogen atoms form the base and one pair of electrons forms the apex of the pyramid. H-N-H angle is 107°.

Bonds:

The three half-filled (containing unpaired electron) sp3 hybrid orbitals of nitrogen overlap axially with three half-filled 1s orbitals of three hydrogen atoms separately to form three covalent N-H bonds (sigma bonds). The fourth hybrid orbital containing lone pair of the electron remains non bonded. The bonds between Nitrogen and hydrogen are sp3- s. H – N – H bond angles are 107°. All N-H bonds in ammonia are of equal strength.

Type and Geometry:

Formation of Water Molecule:

Need of Hybridization:

In oxygen atom there are two half-filled 2p orbitals and valency should be 2 and it is two.  These three ‘p’ orbitals are perpendicular to each other and if they form bonds, the angle between them should be 90°.  But the actual angle between bonding orbitals in water is 104.5°.  To explain this variation the concept of hybridization is required.

Electron Configuration:

The atomic number of Oxygen is 8.  Its configuration in ground state is 1s², 2s², 2p⁴

Hybridization :

In the formation of water molecule one 2s orbital and three 2p orbitals of Oxygen mix up forming four hybrid orbitals of equivalent energy. These four new equivalent orbitals are called sp³ hybrid orbitals. They are identical in all respect. The two hybrid orbitals have paired electrons and they are non – bonding orbitals.  Other two orbitals are half-filled and they are bonding orbitals. The nonbonding pairs of hybridized orbitals are called lone pairs. These hybridized orbitals are in the directions of four corners of a regular tetrahedron.

Angle and Geometry:

Four sp3hybridised orbitals formed, repel each other and they should be directed towards the four corners of a regular tetrahedron and the angle between them should be 109.5°. But due to greater repulsion exerted by lone pair on bonding orbitals, the angle is reduced to 104.5⁰. Thus water has distorted the tetrahedral shape.

The two half-filled (containing unpaired electron) sp³ hybrid orbitals of oxygen overlap axially with two half-filled 1s orbitals of two hydrogen atoms separately to form three O-H bonds. Thus geometry is angular or V-shaped, in which the oxygen lies at the centre, two hydrogen atoms occupy two corners of the tetrahedron and lone pair of electrons occupy remaining two corners of the tetrahedron. H-O-H angle is 104.5⁰.

Bonds:

The two half-filled (containing unpaired electron) sp3 hybrid orbitals of oxygen overlap axially with two half-filled 1s orbitals of two hydrogen atoms separately to form two O-H bonds (sigma bond).  The remaining two hybrid orbitals containing a lone pair of the electron remains nonbonded.

Type and Geometry:

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > sp3 Hybridization

The post sp3 Hybridization appeared first on The Fact Factor.

In this article, we shall study the concept of hybridization of orbitals. This concept overcomes the limitations of valence bond theory.

Need for Hybridization Concept:

A simple approach based on the overlap of s and p orbitals can be applied to many molecules, but it fails to explain the formation of compounds of Beryllium (Be), Boron (B), and carbon (C). The electronic configurations of Be, B, and C in the ground state are as follows.

The atoms of Beryllium Be (Z = 4)  Electronic configuration is 1s² 2s² (With no unpaired electrons).

The atom of Boron B (Z = 5)   Electronic configuration is 1s² 2s² 2p¹ (With one unpaired electron).

The atom of Carbon C (Z = 6)   Electronic configuration is 1s² 2s² 2p² (With two unpaired electrons).

According to valence bond theory valency of an element depends on a number of unpaired electrons in the orbitals. Thus Beryllium, Boron, and Carbon should be zero-valent, monovalent and divalent respectively.  However, these elements form the compounds having valency 2, 3, and 4 respectively. Valence bond theory fails to explain this phenomenon. This is explained by hybridization.

Valence Bond Theory fails to explain the observed geometry of the molecules of water and ammonia e.g. in the formation of H₂O molecule, the H – O – H bond angle should be 90°. But the measured bond angle is 104.3° and the molecule is V-Shaped. Valence bond theory failed to explain this change. To explain The equivalence of bonds we have to use the concept of a process of mixing and recasting of atomic orbitals. For e.g. all the four C-H bonds in methane molecule are equivalent in terms of strength, energy, etc. It can be explained on the basis of hybridization.

Note: The above paragraphs give limitations of the valence bond theory.

Hybridization:

Mixing and recasting or orbitals of an atom (same atom) with nearly equal energy to form new equivalent orbitals with maximum symmetry and definite orientation in space is called hybridization.

Steps Involved in Hybridization:

Step -1: Formation of excited state:

The atom in the ground state takes up some energy and goes to the excited state.  In this process, usually, a pair of electrons in lower energy orbital is split up and one of the electron from this pair is transferred to some empty slightly higher but almost equal energy orbital.  Thus the excited state has a larger number of half-filled orbitals. This new number of half-filled orbitals decides the number of covalent bonds an atom can form.

Step = 2: Mixing and recasting of atomic orbitals:

The orbitals of nearly the same energy in an excited state now hybridize i.e. unite and redistribute themselves giving hybrid orbitals of the same energy and definite orientation in space. Mixing and recasting or orbitals of an atom(same atom) with nearly equal energy to form new equivalent orbitals with maximum symmetry and definite orientation in space is called hybridization.

One 2-s orbital and three p- orbitals mix together and recast themselves to form new four sp³ hybridized orbitals.

Step – 3: Proper Orientation of the Hybrid Orbitals in Space:

The hybrid orbitals then get arranged in space in such a way to minimize mutual repulsion. Each hybrid orbital is more concentrated on one side of the nucleus. Due to this greater overlap is achieved and a stronger bond is formed.

Thus in carbon, the four hybrid sp3 orbitals arrange themselves at four corners of a tetrahedron to minimize mutual repulsion.

Essential Condition for Hybridization:

The orbitals participating in hybridization should have nearly the same energy. Thus in the formation of methane, the 2s and 2p orbitals of carbon have nearly the same energies, so that the recasting of orbitals is possible.  But hybridization of 2s and 3p is not possible because there is much difference between their energies.

Characteristics or Rules of hybridization:

• Atomic orbitals undergoing hybridization should belong to the same atom or ion.
• Atomic orbitals participating in hybridization should have nearly the same energy. Thus 2s and 2p can hybridize, 3s and 3p can also hybridize, but 2s and 3p cannot.
• The total number of hybrid orbitals formed is equal to the number of atomic orbitals involved in the hybridization process.
• All hybrid orbitals are identical with respect to energy and directional character.
• The hybrid orbitals may differ from one other in their orientations.
• The shape of the hybrid orbitals is different from that of the original atomic orbital.
• Similar to atomic orbitals, each hybrid orbital can have a maximum of two electrons.
• The hybrid orbitals are concentrated in one particular direction to achieve greater overlapping.
• The hybrid orbitals have maximum symmetry and definite orientation in space so that the mutual force of repulsion of electrons is avoided.

Types of Hybridization and Geometry of Molecules:

The hybridization involving s and p orbitals are of the following three types:

• Tetrahedral or sp³ hybridization e.g. CH₄, NH₃, H₂O
• Trigonal or sp² e.g. BF₃, C₂H₄.
• Diagonal or sp hybridization e.g. BeF₂, C₂H₂

Their names indicate the orientation of the orbitals in space and the designation (sp², sp³, etc) indicates the number and types of atomic orbitals involved in hybridization.

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Hybridization of Orbitals

The post Hybridization of Orbitals appeared first on The Fact Factor.

According to Valence Bond Theory, covalent bonds are of two types a) Sigma bond (σ) and b) Pi bond (π)

Sigma bond (σ):

The covalent bond formed as a result of an end to end overlap i.e. head-on collision of atomic orbitals Or A covalent bond formed by collinear or coaxial i.e. in a line of internuclear axis overlapping of an atomic orbital is known as a sigma bond. Example: In hydrogen molecule, there is a sigma bond between two overlapping ‘s’ orbitals.

Formation of Sigma Bond:

A covalent bond formed by collinear or coaxial i.e. in a line of internuclear axis overlapping of an atomic orbital is known as a sigma bond. S orbitals are non-directional hence they can overlap in any side. Thus s-s overlap always forms a sigma bond. In order to form sigma bond p orbitals must lie along the internuclear axis. They are formed by s-s overlap e.g. H-H, s-p overlap e.g. H-F, p-p overlap e.g. F-F

Characteristics of Sigma Bond:

• It is a covalent bond formed by a coaxial overlap of bonding orbitals.
• It is a very strong bond, due to a greater extent of overlapping.
• It is possible between any two orbitals s-s, p-p or s-p and also hybrid orbitals.
• Bond energy is more.

Pi Bond (π):

A covalent bond is formed by lateral or sideway or collateral overlapping of pure orbitals is known as a pi bond. Example: Ethylene has one pz – pz  π bond. Acetylene has two π bonds.

Formation of pi Bond:

A covalent bond is formed by lateral or sideway or collateral overlapping of pure orbitals is known as a pi bond. A pi bond is formed by a lateral overlap of two p orbitals oriented mutually parallel but perpendicular to the internuclear axis.

Characteristics of Pi Bond:

• It is a covalent bond formed by a lateral or sideways overlap of atomic orbitals.
• It is a weak bond due to the lower extent of overlapping.
• It is possible between two ‘p’ orbitals only.
• Bond energy is less.
• It is to be noted that the pi bond is formed when the sigma bond already exists. The formation of only a pi bond is not possible.

Sigma Bond is Stronger Than a pi Bond:

The extent of overlapping of orbitals along the same axes is always greater than the extent of overlapping at an angle. Also, the larger the overlapping of orbitals, the stronger is the covalent bond.

A sigma bond is formed by co-axial overlapping of atomic orbitals. Due to the large overlap of atomic orbitals, the bond involves more evolution of energy than Pi bond. Due to the large overlap of atomic orbitals, the bond involves more evolution of energy than Pi bond. In this bond, the electron density between two nuclei on the internuclear axis is very high.

A pi bond is formed by collateral or sidewise overlapping of atomic orbitals. Due to the less overlap of atomic orbitals, the bond involves less evolution of energy than the sigma bond. In the case of a pi bond, the electron density is higher above and below the internuclear axis and not on the nuclear axis.

Hence the sigma bond in which overlapping of orbitals take place along the same axis is stronger than the pi bond in which overlapping of orbitals takes place at an angle (laterally).

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Sigma bonds and Pi Bonds

The post Sigma bonds and Pi Bonds appeared first on The Fact Factor.

In this article, we shall study the formation of bonds by overlapping of orbitals and different types of overlaps.

Formation of Hydrogen Molecule on The Basis Of Orbital Overlap:

The electronic configuration of a hydrogen atom is 1s¹. It contains one unpaired electron in its valence shell. Therefore 1s orbital in the hydrogen atom is bonding orbital. When the two hydrogen atoms having valence electrons with opposite spins approach each other, the attractive force dominates the repulsive force between the electrons in the initial stage. Thus, as the distance between two hydrogen atoms decreases the potential energy of the system gradually decreases.

A state of minimum potential energy is reached when the forces of attraction are balanced by the forces of repulsion between two hydrogen atoms. At this point the energy of the system is minimum and the stability is maximum. At this stage orbital of two hydrogen atoms, overlap and spins of electrons are neutralized and a stable covalent bond is formed between hydrogen atoms, forming H₂, molecule. In this case, the overlap of orbitals is maximum. This overlap is called s-s overlap and bond formed are coaxial overlapping of two s orbital is called s-s sigma (σ) bond.

In the formation of a molecule, the release of energy comes from two sources. The neutralization of spin magnetic moments of the two electrons and the accumulation of electron density between two nuclei.

Diagram :

1 S Orbital          1 S Orbital         S-S overlap

Reason: Helium does not form a diatomic molecule:

Helium with atomic number 2 has electron configuration 1s2. Thus helium contains two paired electrons in its 1s orbital. The pairing of electrons in 1s orbital neutralizes the spin magnetic moments of each other. This gives stability to the helium atom. There are no empty orbitals in the first shell for unpairing and promoting these paired electrons to higher energy orbital. Thus helium has no unpaired electron i.e. no bonding orbital. Therefore according to valence bond theory, it cannot form a covalent bond with another helium atom.

Each helium atom has paired electrons in 1s is orbital.  If two such orbitals come close to each other, there is net repulsion between them because repulsive forces are stronger than the attractive forces.  If they do so it increases the potential energy of the system and it becomes unstable.  Thus overlapping of orbitals cannot take place. Therefore helium cannot form a diatomic molecule. It exists as a monoatomic gas.

Types of Overlap of Atomic Orbitals:

The term overlap refers to the overlap of the atomic orbitals of the two approaching atoms as they enter into the bond formation stage. In the case of s and p orbitals, there can be three types of overlap.

s – s orbital overlap ( formation of H₂ molecule):

The mutual overlap between the half-filled s orbitals of two atoms is called s – s overlap and the covalent bond formed is known as sigma (s) bond. e.g. formation of a hydrogen molecule from two hydrogen atoms.

s – orbital is spherical in shape and overlapping takes place to some extent in all directions.  Hence s -s bond is non – directional.

Hydrogen (1s¹) atom has 1s orbital containing a single electron i.e. it is half-filled.  Two such 1s orbitals from the two hydrogen atoms having electrons with opposite spins approach each other, then the potential energy of the system decreases. The two ‘s’ orbitals overlap each other when they acquire minimum potential energy, forming H-H sigma bond or s-s overlap. H-H bond is a nonpolar covalent bond.

As the two orbitals are overlapping such that the overlapped region lies on the line joining the two nuclei of the overlapping orbitals (axial overlapping) bond formed is sigma bond.

Diagram :

1 S Orbital          1 S Orbital         S-S overlap

p – p orbital overlap (Formation of Fluorine F₂ molecule):

The mutual overlap between two half-filled p – orbitals of two atoms is called p – p overlap and the covalent bond formed is known as p – p bond.  If the overlapping takes place along the internuclear axis the bond is called sigma bond and if the overlapping takes place literally the bond is known as pi bond. e.g. formation of fluorine molecule from two fluorine atoms. CI₂, Br₂and I₂ are also formed by p – p overlap.

The atomic number of fluorine is 9. The electronic configuration of the fluorine atom is 1s². 2s². 2p_x², 2p_y², 2p_z¹. Each F atom has one unpaired electron in p – orbital (p_z orbital). When two fluorine atoms each containing unpaired electron with opposite spins approach each other, then the potential energy of the system decreases. The two ‘p’ orbitals overlap each other when they acquire minimum potential energy.

As the two orbitals are overlapping such that the overlapped region lies on the line joining the two nuclei of the overlapping orbitals (axial overlapping) bond formed is sigma bond. As p orbitals are dumbbell-shaped they overlap in a particular direction. Therefore the p-p bond is directional. The back lobe of the overlapping p orbital is distorted and its size is reduced. This is due to the tendency of the electron to remain in the overlapping region.

F-F bond is a non-polar, as shared pair of electrons is attracted by both the atoms.

Diagram :

1 P Orbital                    1 P Orbital                          P-P overlap

s – p orbital overlap (Formation of Hydrogen Fluoride Molecule):

The overlap between the half-filled s – orbital of one atom and the half-filled p – orbital of another atom is called s – p overlap and the covalent bond formed is known as s – p sigma bond. E.g.: Formation of HF molecule, H – X bond in HCI, HBr, and HI are also formed by s-p overlap.

The electronic configuration of a hydrogen atom is 1s1, while that of fluorine atom is 1s2. 2s2. 2px2, 2py2, 2pz1. Thus 1s orbital of a hydrogen atom and the 2p orbital of fluorine atom containing unpaired electron can overlap and the bond formation is possible provided the spin of electrons in overlapping orbitals is opposite. The bond in hydrogen fluoride is a s – p bond. As the two orbitals are overlapping such that the overlapped region lies on the line joining the two nuclei of the overlapping orbitals (axial overlapping) bond formed is sigma bond. The back lobe of the overlapping p orbital is distorted and its size is reduced. This is due to the tendency of the electron to remain in the overlapping region.

Fluorine is more electronegative than hydrogen, hence H-F bond is polar.

Diagram :                                

Reason: H – F bond is polar.

HF molecule is formed by the s-p overlap of orbitals. On Pauling scale, the electronegativity of H is 2.1 and that of F is 4. Thus fluorine is more electronegative than H, the electronic cloud in this bond is displaced more towards fluorine.  In other words, the electron pair shared between the two atoms is attracted more towards fluorine. This gives fluorine atom a partial negative charge (δ -) and hydrogen atom a partial charge   (δ +).

Thus the H – F bond is not purely covalent.  It possesses a partial ionic character and is said to be a polar bond.  The bond has 43% ionic character. H – X bond in HCI, HBr, and HI are also polar since CI, Br and I are more electronegative than H.

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Overlapping of Orbitals

The post Overlapping of Orbitals appeared first on The Fact Factor.

Valence bond theory was proposed by Heitler and London in (1927) and it was extended by Pauling and Slater (1931). It is based on the following concepts.

• The pairing of electrons.
• Neutralization of opposite electron spins.
• Overlapping of orbitals containing unpaired electrons to give a region of common electron density to the combining atoms.

Postulates of Valence Bond Theory:

 Condition of bond formation: 

A covalent bond is formed when the atomic orbital of one atom overlaps the atomic orbital of the other atom. The sharing of electrons takes place due to the overlapping of half-filled orbitals of the outermost shell of the two atoms. Only atomic orbital with unpaired electrons with opposite spin can participate in overlapping or bonding.  They are known as Bonding Orbitals.

Binding force of covalent bond: 

During the overlapping of half-filled orbitals, spins get neutralized and electron density between the two nuclei increases. As a result of this, the force of repulsion between the two nuclei decreases, while the force of attraction between the nucleus of one and the electron of other increases.

Strength of a bond:

Bond is formed when the system attains a stable lowest energy level.  This occurs at a certain minimum equilibrium distance between two nuclei known as bond length. The strength of the covalent bond depends upon the extent of overlapping between two atoms.  Greater the overlap, stronger is the bond. Complete overlapping is not possible as there are repulsive forces between two nuclei.

Number of covalent bonds:

One bonding orbital can form only one covalent bond.  Hence the number of bonding orbitals restricts the number of covalent bonds that can be formed by an atom. The number of unpaired electrons possessed by an atom determines its valency.

Directional nature of covalent bond:

S- orbitals are non-directional but p, d, and f  – orbitals have a particular direction. Due to their directional nature, the geometry of the molecule depends upon the orientation of the overlapping of the orbitals.

Geometry of a Molecule on the Basis of Valence Bond Theory:

A covalent bond is formed when the atomic orbital of one atom overlaps the atomic orbital of the other atom. The bond has maximum electron density in the region of overlap. The line joining the nuclei of two atoms determines the direction of the bond.

‘S’ orbitals are spherical and non – directional, but p, d, f orbitals are directionally oriented and produce overlap in the internuclear axis resulting in a bond formation having s a specific direction. The geometry of molecules and bond angles therein are determined by the directional orientation of orbitals involved in the overlap. The electron pairs in the valence shell of the central atom arrange themselves symmetrically so as to get as far apart as possible due to repulsion between them.

e.g. Methane has tetrahedral geometry, Boron trifluoride has planar geometry while Beryllium difluoride has linear geometry, etc.

Energy Changes in Covalent Bond Formation on the Basis of Valence Bond Theory:

Interactive Force Between Approaching Molecules:

It must be realized that when any two atoms approach close to each other new forces of attraction and repulsion set in.

The forces of attraction are between the nucleus of one atom and the electron of the other and vice – versa. The forces of repulsion are between two nuclei amongst themselves as well as between the electrons of the two atoms amongst themselves.  The force of attraction and repulsion are represented in the figure.

If the net result is attraction i.e. attractive forces are stronger than the repulsive forces the total energy of the system decreases and a chemical bond results.   If the net result is repulsion i.e. the repulsive forces are stronger than attractive forces, the total energy of the system increases and no chemical bonding is not possible.

Variation in Energy of System Due to Approach of Orbitals for Overlapping:

When two Hydrogen atoms with unpaired electrons with parallel spin approach each other, repulsion is greater than attraction and energy of system increases and bonding does not take place. In this case, the energy goes on increasing as inter-nuclear distance starts decreasing.

When two Hydrogen atoms with unpaired electrons with opposite spin approach each other, the attraction is greater than repulsion and energy of system continuously decreases till it becomes minimum at equilibrium distance between the two nuclei. There is overlap and leads to covalent bond formation between two hydrogens.

Formation of Hydrogen Molecule in terms of Decrease of Potential Energy:

The way in which the potential energy of the system changes as two hydrogen atoms having electrons with opposite spins approaches each other to form a covalent bond is represented in the potential energy curve. Graphical representation of the change in potential energy as a function of internuclear distance is known as the Potential Energy Curve. The electronic configuration o the hydrogen atom is 1s¹.  It contains one unpaired electron in its valence shell.

Explanation of Curve – 1

• State A: When two hydrogen atoms are far apart from each other, then the potential energy of one atom is independent of the other.  This potential energy is taken arbitrarily as zero
• State B: As the two hydrogen atoms having electrons with opposite spins approach each other, the electrons of one atom begin to feel the attractive forces of the nucleus of the other atom.  The attractive force dominates the repulsive force between the electrons.  Thus, as the distance between two hydrogen atoms decreases the potential energy of the system gradually decreases.
• State C: A state of minimum potential energy is reached when the forces of attraction are balanced by the forces of repulsion between two hydrogen atoms.  At this stage orbital of two hydrogen atoms, overlap and spins of electrons are neutralized and a stable covalent bond is formed between hydrogen atoms, forming a hydrogen molecule. In this case, the overlap of orbitals is maximum. The minimum equilibrium distance between the two hydrogen nuclei at this stage is called bond length which is 74 picometer. The amount of energy released at this stage is called bond energy and it is equal to 431.8 kJ/mole.8.
• State D:  When two hydrogen atoms are forced to come still closer, then the potential energy of the system shows a sharp rise.  This is due to the increases of repulsive forces between the two nuclei at such small inter-nuclear distance and hence bond formed becomes unstable.

Explanation of Curve – 2

• State E: It two hydrogen atoms having electrons with parallel spins approach each other, potential energy of system increases and no bond is formed between the two hydrogen atoms.

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Valence Bond Theory

The post Valence Bond Theory appeared first on The Fact Factor.