NAVODAYA VIDYALAYA SAMITI NOIDA

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E-CONTENT PREPARATION

FOR

CLASS 11^TH CHEMISTRY

CHAPTER-4: CHEMICAL BONDING AND MOLECULAR STRUCTURE

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PREPARED BY: C. RAJASEKHAR, PGT CHEMISTRY,

JNV CHIKKABALLAPURA, KARNATAKA

CHEMICAL BONDING AND MOLECULAR STRUCTURE-VSEPR THEORY

WE KNOW THAT LEWIS CONCEPT IS UNABLE TO EXPLAIN THE SHAPES OF MOLECULES. THE VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY

VSEPR THEORY PROVIDES A SIMPLE PROCEDURE TO PREDICT THE SHAPES OF COVALENT MOLECULES.THIS THEORY IS BASED ON THE REPULSIVE INTERACTIONS OF THE ELECTRON PAIRS IN THE VALENCE SHELL OF THE ATOMS.

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THE MAIN POSTULATES OF VSEPR THEORY

1. THE SHAPE OF A MOLECULE DEPENDS UPON THE NUMBER OF VALENCE SHELL ELECTRON PAIRS (BONDED OR NON –BONDED) AROUND THE CENTRAL ATOM
2. PAIRS OF ELECTRONS IN THE VALENCE SHELL REPEL ONE ANOTHER SINCE THEIR ELECTRON CLOUDS ARE NEGATIVELY CHARGED

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3. THESE PAIRS OF ELECTRONS OCCUPY SUCH POSITIONS IN SPACE THAT MINIMISE REPULSION AND THUS MAXIMISE DISTANCE BETWEEN THEM.

4. THE VALENCE SHELL IS TAKEN AS A SPHERE WITH THE ELECTRON PAIRS LOCALISING ON THE SPHERICAL SURFACE AT MAXIMUM DISTANCE BETWEEN THEM.

5. A MULTIPLE BOND IS TREATED AS IF IT IS A SINGLE ELECTRON PAIR AND THE TWO OR THREE ELECTRON PAIRS OF A MULTIPLE BOND ARE TREATED AS A SINGLE SUPER PAIR.

VSEPR MODEL IS APPLICABLE TO ANY ONE STRUCTURE OF RESONANCE STRUCTURES

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THE REPULSIVE INTERACTION OF ELECTRON PAIRS IN THE ORDER�

LONE PAIR(LP)- LONE PAIR(LP) > LONE PAIR(LP) –BOND PAIR(BP) >BOND PAIR(BP)-BOND PAIR(BP)

The lone pairs localized on the central atom ,each bonded pair is shared between two atoms. As a result , the lp electrons in a molecule occupy more space compared to bp of electrons This results in greater repulsion between lone pairs of electrons as compared to the lone pair –bond pair repulsions. These repulsion effects result in deviations from idealized shapes and alterations in bond angles in molecules.

� As �(i)molecules in which the central atom has no lone pair and � (ii) molecules in which central atom has one or more lone pairs

VSEPR THEORY DIVIDES MOLECULES IN TO TWO CATEGORIES�

Below are the examples for linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral

SUCH ARRANGEMENT CAN BE SEEN IN THE MOLECULES LIKE BF₃ ,CH₄ , AND PCL₅ AS DEPICTED BY THEIR BALL AND STICK MODELS

VALENCE BOND THEORY

• It is based on the knowledge of atomic orbitals, E.C of elements, the overlap criteria of AOs and the principle of variation and superposition . Let us consider the formation of H₂ molecule which is the simplest of all molecules.
• Consider two H atoms A and B approaching each other having nuclei N_A and N_B and electrons e_A and e_B .When the two atoms are at large distance from each other, there is no interaction between them.As these two atoms approach each other ,new attractive and repulsive forces operate.

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Attractive forces arise between :

• (i) N_A –e_A and N_B –e_B
• (ii)N_A -e_B , N_B – e_A

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Repulsive forces arise between

• (i) electrons of two atoms e_A – e_B ,
• (ii) nuclei of two atoms N_A – N_B

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Attractive forces tend to bring the two atoms close to each other where as repulsive forces tend to push them apart

ORBITAL OVERLAP CONCEPT

• Partial merging of atomic orbitals is called overlapping of atomic orbitals which results in the pairing of electrons.
• The extent of overlap decides the strength of a covalent bond. Greater the overlap the stronger is the bond formed between two atoms. Therefore the formation of a covalent bond between two atoms results by pairing of electrons present in the valence shell having opposite spins

OVERLAPPING OF ATOMIC ORBITALS

When two atoms come close to each other ,there is overlapping of AOs. The various arrangements of s and p orbitals resulting in +ve, -ve, and zero overlap are depicted in fig

HYBRIDISATION

• Acc. to Pauling the AOs combine to form new set of equivalent orbitals known as hybrid orbitals.
• The hybridisation is defined as the process of intermixing of slightly different energies so as to redistribute their energies resulting in the formation of new set of orbitals of equivalent energies and shape.

SALIENT FEATURES OF HYBRIDISATION

1. The no. of hybrid orbitals is equal to the no. of AOs that are hybridised

2. The hybrid orbitals are always equivalent in energy and shape

3. The hybrid orbitals are more effective in forming stable bonds

IMPORTANT CONDITIONS FOR HYBRIDISATION

(i) The orbitals present in the valence shell of the atom are hybridised.

(ii) The orbitals undergoing hybridisation should have almost equal energy.

(iii) Promotion of electron is not essential condition prior to hybridisation.

(iv) Half filled or even filled orbitals of valence shell take part in hybridisation.

TYPES OF HYBRIDISATION

SP hybridisation:

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• It involves the mixing of one s and one p orbital resulting in the formation of two equivalent SP hybrid orbitals.
• Each SP hybrid orbital has 50% s- character and 50% p- character and such a molecule possesses linear geometry and is known as diagonal hybridisation

EXAMPLE OF MOLECULE HAVING sp HYBRIDISATION�

BeCl₂ : In Be one of the 2s electron is excited to vacant 2p orbital to account for divalency and thus one 2s and one 2p orbitals hybridise to form sp hybrid orbitals and are oriented in opposite direction forming an angle of 180⁰

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sp² Hybridisation

It involves the mixing of one s and two p orbitals resulting three sp² orbitals.

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Ex: BCl₃ in which one of the 2s electron is excited to vacant 2p orbital resulting in three unpaired electrons and these orient in a trigonal planar arrangement and overlap with 2p orbitals of Cl atoms and the bond angle between ClBCl is 120⁰

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• In case of H₂ O molecule one 2s and three 2p orbitals form 4 sp³ orbitals ,out of which two contain one electron each and the other two contain a pair of electrons.
• These four hybrid orbitals acquire a tetrahedral geometry ,with two corners occupied by lone pairs and the other two corners occupied by bond pairs and the bond angle is reduced to 104.5⁰ and the molecule thus gets a V- shape or angular geometry

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sp³ Hybridisation

• In CH₄ mixing of one s orbital and three p orbitals takes place resulting in four SP³ hybrid orbitals having 25% s character and 75% p character and these four hybrid orbitals are directed towards the corners of a tetrahedron with a bond angle of 109.5^0. Other examples are NH₃ and H₂ o and the bond angles are 107⁰ and 104.5⁰ respectively

STRUCTURE OF NH₃ AND H₂O

• In NH₃ the valence E.C of nitrogen in the ground state is 2s² 2p_x¹ 2p_y¹ 2p_z¹ having three unpaired es in the sp³ hybrid orbitals and a lone pair of electrons in the 4^th one. These three hybrid orbitals overlap with 1s orbitals of three H atoms to form 3 N-H sigma bonds.
• The molecule thus gets distorted due to more repulsion between lp-bp than between two bond pairs of electrons and the molecule gets distorted tetrahedron with a bond angle of 107⁰

OTHER EXAMPLES OF sp³ , sp² and sp HYBRIDISATION

• sp³ in C₂H₆ molecule: In ethane both carbons assume sp³ hybrid state.

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• One orbital overlaps axially with other carbon to form sp³ –sp³ sigma bond while the other 3 hybrid orbitals form sp³ –s sigma bonds with H atoms. In ethane C-C bond length is 154 pm and each CH bond length is 109 pm

sp² HYBRIDISATION IN C₂ H₄

• In ethene one of the sp² orbitals of each carbon overlaps axially with each other to form a C-C sigma bond while the other two sp² orbitals of each carbon forms sp² –s sigma bond with two H atoms. The unhybrid orbital of each carbon overlaps side wise to form a pie bond which has equal electron clouds above and below the plane .

sp HYBRIDISATION IN C₂ H₂ MOLECULE

• In this both carbons undergo SP hybridisation having two unhybridised orbitals i.e. 2p_x and 2p_y

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• One sp orbital of each carbon overlaps with each other to form c-c sigma bond and other sp hybrid orbital overlaps axially with H atoms forming sigma bonds .The two un hybrid p orbitals present on each carbon overlap side wise to form two pie bonds.

HYBRIDISATION OF ELEMENTS INVOLVING d ORBITALS

Formation of PCl₅ molecule(SP³ d hybridisation):

• The energy of 3d orbitals are comparable to those of 3s , 3p as well as 4s and 4p orbitals .therefore the hybridisation involving either 3s ,3p and 3d or 3d, 4s and 4p is possible. For ex: trigonal bipyramidal-sp³ d
• Square planar- dsp² ,square pyramidal-sp³ d² and octahedral - sp³ d² or d² sp³

FORMATION OF PCl₅

On excitation one of 3s electrons is promoted to vacant 3d orbital making five unpaired electrons for bonding with five Cl atoms forming a trigonal bi-pyramidal structure in which three bonds called equatorial bonds lie in a plane and the other two bonds called axial bonds are longer than equatorial bonds due to repulsive interaction from equatorial bond pairs and are weaker .

FORMATION OF SF₆

On excitation six orbitals with unpaired electrons are available for bonding with six f atoms forming SP³ d² hybridisation with an octahedral geometry

MOLECULAR ORBITAL THEORY

Salient features:

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(i) The electrons in a molecule are present in various MOs

(ii) The AO s of comparable energy and proper symmetry combine to form MOs

(iii) AO is monocentric where as MO is polycentric

(iv) The no. of MOs formed are equal to the no. of AOs combined

(v)BMO has lower energy and greater stability than AMO

(vi)The electron probability around a group of nuclei is given by MO

(vii)The MOs are filled in accordance with Aufbau principle, Hund’s rule and Pauli’s exclusion principle.

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FORMATION OF MOLECULAR ORBITALS BY LINEAR COMBINATION OF ATOMIC ORBITALs�

• The formation of molecular orbitals can be understood in terms of the constructive or destructive interference of the electron waves of the combining atoms.
• The sigma MO is formed by the addition of atomic orbitals called bonding moleclular orbital and pie MO is formed by the subtraction of atomic orbitals called antibonding molecular orbital

• Electron density in BMO is located between the nuclei of bonded atoms and the repulsion between the nuclei is very less and in case of ABMO electron density is located away from the space between the nuclei.
• There is a nodal plane between the nuclei and the repulsion between the nuclei is very high.
• Electrons placed in a BMO tend to hold the nuclei together and stabilise a molecule.Therefore BMO always possesses lower energy and ABMO higher energy

• The electrons placed in ABMO destabilise the molecule and the energy is raised above the energy of the parent AOs and the energy of BMO is lowered than the parent AOs.
• The total energy of two MOs remains the same as that of original atomic orbitals

CONDITIONS FOR THE COMBINATION OF ATOMIC ORBITALS

1. The combining AOs must have the same or nearly the same energy means 1s can combine with 1s only

2. The combining AOs must have same symmetry about the molecular axis. For ex 2p_z can combine with 2p_z only but not with 2p_x or 2p_y

3. The combining AOs must overlap to the maximum extent. Greater the extent of overlap, the greater will be the electron density between the nuclei of a molecular orbital

TYPES OF MOLECULAR ORBITALS AS σ,π,δ etc

• The sigma MOs are symmetrical around the bond axis while pie MOs are not symmetrical.
• The linear combination of 1s orbitals of two atoms produces two MOs namely σ1s and σ^*1s .Similarly a linear combination of 2p_z orbitals of two atoms produces σ2p_z and σ^*2p_z .
• Molecular orbitals formed from 2p_x and 2p_z are not symmetrical around the bond axis because of the presence of positive lobes above and negative lobes below the molecular plane Such MOs are π and π* . The π* AMO has a node between the nuclei

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ENERGY LEVEL DIAGRAMS FOR MOLECULAR ORBITALS

The increasing order of energies of various molecular orbitals for O₂ and F₂ is given below:

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For molecules such as B2 , C2 , N2 , etc. the increasing order of energies of various molecular orbitals is

ELECTRONIC CONFIGURATION AND MOLECULAR BEHAVIOUR

Stability of molecules: If N_b is the no. of es in BMO and N_a is the no. of es in ABMO then

(i) the molecule is stable if N_b > N_a

(ii) the molecule is unstable if N_b < N_a

Bond order: = ½( N_b –N_a ) . A positive bond order means a stable molecule while a negative or zero bond order means an unstable molecule.

Bond length : The bond length decreases as bond order increases

Magnetic character : If all the MOs in a molecule are paired, the molecule is diamagnetic or if one or more MOs are singly occupied it is paramagnetic ( attracted by magnetic field) Ex: O₂ molecule

HYDROGEN BONDING

• It is defined as the attractive force which binds H atom of one molecule with the electronegative atom(F,N or O) of another molecule.

Cause of formation of hydrogen bond:

Due to high EN of X attached to H the electrons are displaced towards X (acquires a partial –ve charge ) and H acquires a partial +ve charge resulting in the formation of a polar molecule having electrostatic force of attraction as H-X---H-X—H_X----.The magnitude of H bonding is maximum in solid state and minimum in gaseous state and it has strong influence on the structure and properties of the compounds

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TYPES OF HYDROGEN BONDS

1. Intermolecular hydrogen bond: It is formed between two different molecules of the same or different compounds. For ex H-bond in case of HF molecule ,alcohol or water molecules etc.

2. Intramolecular hydrogen bond: It is formed when H atom is in between the two highly electronegative (F,O,N)atoms present within the same molecule. For ex ,in ortho-nitrophenol the hydrogen is in between the two oxygen atoms.