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This document covers key concepts in chemical bonding and molecular structure, including the octet rule, types of bonds (covalent, ionic), limitations of the octet rule, bond parameters, resonance structures, and factors affecting bond polarity. It explains the formation and characteristics of covalent and ionic bonds, the concept of formal charge, dipole moments, and the application of Fajans' rules to determine the covalent character of ionic bonds. The document is structured to provide a deep understanding of the theories and principles governing chemical bonding.

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Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Octet Rule • Covalent Bond • Limitations of Octet Rule By-Vijay Kumar Sethi
Chemical bond • The attractive force which holds various constituents (atoms, ions, etc.) together in different chemical species is called a chemical bond. Valence electrons:-Electrons present in the outermost shell are known as valence electrons Lewis symbols:-Valence electrons of an atom are represented by dots around the symbol of the element Significance of Lewis Symbols • Valence electrons helps to calculate the common or group valence of the element. • The group valence of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.
Octet Rule • Theory of chemical combination between atoms is known as electronic theory of chemical bonding. • It was developed by Kössel and Lewis in 1916. According to this, • atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. • This is known as octet rule.
Covalent Bond • Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. • Langmuir introduced the term covalent bond. • According to Lewis-Langmuir theory a covalent bond is formed by sharing a pair of electron between two bonded atoms
Single covalent bond • If covalent bond is formed by sharing a pair of electron it is called a single covalent bond and it is represented by a line between the two atoms Double Bond:- If two atoms share two pairs of electrons Triple Bond:- If two atoms share three pairs of electrons • Multiple bonds- double and triple bonds collectively known as multiple bonds
Limitations of the Octet Rule 1) The incomplete octet of the central atom:- • This is especially the case with elements having less than four valence electrons. • Examples are LiCl, BeH2 and BCl3. • Li, Be and B have 1,2 and 3 valence electrons only. • Some other such compounds are AlCl3 and BF3.
Continue… Limitations of the Octet Rule 2) Odd-electron molecules • In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO2, the octet rule is not satisfied for the N atoms
Continue… Limitations of the Octet Rule 3) The expanded octet:- • Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. • In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet. sulphur also forms many compounds in which the octet rule is obeyed. For example sulphur dichloride SCl2
Continue… Limitations of the Octet Rule 4) octet rule is based upon the chemical inertness of noble gases. However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like XeF2, KrF2, XeOF2 etc., 5) This theory does not account for the shape of molecules. 6) It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Lewis dot structures • Formal Charge By-Vijay Kumar Sethi
Lewis dot structures • The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule
Writing of Lewis dot structures The Lewis dot structures can be written by adopting the following steps: Step 1:- Count the total number of valence electrons. For anions, add one electron for each negative charge and for cation subtract one electron for each positive charge. Step 2:- Write skeletal structure (In general the least electronegative atom occupies the central position) Step 3:-Draw a single bond by one shared electron pair between bonded atoms. Remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs. If necessary rearrange the electrons as multiple bonds or lone pairs to complete the octet.
Problem 4.1 Write the Lewis dot structure of CO molecule. Solution • Step 1. Count the total number of valence electrons of carbon and oxygen atoms. The valence electrons available are 4 + 6 =10. • Step 2. The skeletal structure of CO is written as: C O • Step 3. Draw a single bond (one shared electron pair) between C and O and complete the octet on O, the remaining two electrons are the lone pair on C. • This does not complete the octet on carbon and hence we have to resort to multiple bonding (triple bond) between C and O atoms.
Formal Charge • In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom. • It is, however, feasible to assign a formal charge on each atom. • The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure.
Continue… Formal Charge:- Consider the ozone molecule (O3). The Lewis structure of O3 may be drawn as • Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species. • Generally the lowest energy structure is the one with the smallest formal charges on the atoms.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Ionic Or Electrovalent Bond • Lattice Enthalpy By-Vijay Kumar Sethi
Ionic Or Electrovalent Bond Kössel, drew attention to the following facts: ❖ In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases; ❖ The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of an electron by the respective atoms; ❖ The negative and positive ions thus formed attain stable noble gas electronic configurations. The noble gases (with the exception of helium which has a duplet of electrons) have a particularly stable outer shell configuration of eight (octet) electrons, ns2 np6 . ❖ The negative and positive ions are stabilized by electrostatic attraction.
Continue… Ionic Or Electrovalent Bond • For example, the formation of NaCl • Formation of CaF2 The bond formed, as a result of the electrostatic attraction between the positive and negative ions is termed as the electrovalent bond. The electrovalence is equal to the number of unit charge(s) on the ion. Thus, calcium is assigned a positive electrovalence of two, while chlorine a negative electrovalence of one.
Ionic Bonds: One Big Greedy Thief Dog!
Continue… Ionic Or Electrovalent Bond ❑ The formation of ionic compounds would depend upon: ❑ The ease of formation of the positive and negative ions from the respective neutral atoms; ❑ The arrangement of the positive and negative ions in the solid, that is, the lattice of the crystalline compound. • Ionic bonds will be formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy. • Lattice enthalpy plays a key role in the formation of ionic compounds
Lattice Enthalpy • The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. • For example, the lattice enthalpy of NaCl is 788 kJ mol–1 . This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+ (g) and one mole of Cl– (g) to an infinite distance. NaCl(s) → Na+ (g) + Cl– (g) • The solid crystal being three dimensional; it is not possible to calculate lattice enthalpy directly from the interaction of forces of attraction and repulsion only. It is determined indirectly using Born-Haber Cycle .
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Bond Parameters-Bond Length, Bond Angle, Bond Enthalpy, Bond Order By-Vijay Kumar Sethi
Bond Parameters Bond Length • Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. • Bond lengths are measured by spectroscopic, X-ray diffraction and electron-diffraction techniques • Each atom of the bonded pair contributes to the bond length . • In the case of a covalent bond, the contribution from each atom is called the covalent radius of that atom.
• The covalent radius is half of the distance between two similar atoms joined by a covalent bond in the same molecule. • The van der Waals radius is half of the distance between two similar atoms which are not bonded Covalent radius and van der Waals radius
Continue… Bond Parameters Bond Angle • It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. • Bond angle is expressed in degree which can be experimentally determined by spectroscopic methods. • It helps in determining the shape of molecule/complex ion. • For example H–O–H bond angle in water is 104.5o
Continue… Bond Parameters Bond Enthalpy • It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. • The unit of bond enthalpy is kJ mol–1 . Larger the bond dissociation enthalpy, stronger will be the bond in the molecule. Enthalpy of atomization
Continue… Bond Parameters Bond Enthalpy • In case of polyatomic molecules like water H2O • Two O-H bond have different values of bond enthalpy because of changed chemical environment. • Therefore in polyatomic molecules the term mean or average bond enthalpy is used. • It is obtained by dividing total bond dissociation enthalpy by the number of bonds broken
Molecule No. of Shared electron pairs Number of covalent bonds Bond Order Hydrogen (H2) One 1 1 Oxygen (O2) Two 2 2 Nitrogen (N2) Three 3 3 Carbon monoxide (CO) Three 3 3
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Concept of Resonance By-Vijay Kumar Sethi
Resonance Structures • A single Lewis structure of a molecule is not sufficient to describe its experimentally determined parameters. For example, Ozone O3. • The normal O–O and O=O bond lengths are 148 pm and 121 pm respectively. • Experimentally determined oxygen-oxygen bond lengths in the O3 molecule are same (128 pm).
Continue… Resonance Structures • According to the concept of resonance , whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately. • Resonance is represented by a double headed arrow.
Problem 4.4:- Explain the structure of CO2 molecule. Solution:- • The experimentally determined carbon to oxygen bond length in CO2 is 115 pm. • The lengths of a normal (C=O) and (C≡O) are 121 pm and 110 pm respectively. • So a single Lewis structure cannot depict this position and it becomes necessary to write more than one Lewis structures and to consider that the structure of CO2 is best described as a hybrid of the canonical or resonance forms I, II and III.
Continue… Resonance Structures • In general, it may be stated that ❖ Resonance stabilizes the molecule as the energy of the resonance hybrid is less than the energy of any single canonical structure; and, ❖ Resonance averages the bond characteristics as a whole. Resonance Energy:- • Energy difference between energy of the resonance hybrid and energy of the most stable canonical structure is known as resonance energy
Continue… Resonance Structures Remember that : ❖ The canonical forms have no real existence. ❖ The molecule does not exist for a certain fraction of time in one canonical form and for other fractions of time in other canonical forms. ❖ There is no equilibrium between the canonical . ❖ The molecule as such has a single structure which is the resonance hybrid of the canonical forms and which cannot as such be depicted by a single Lewis structure.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Polar covalent bond • Dipole moment • Covalent Character in Ionic Bond-Fajans Rule By-Vijay Kumar Sethi
Polarity of Bonds- Polar covalent bond • When covalent bond is formed between two similar atoms, for example in H2, O2, Cl2, N2 or F2, the shared pair of electrons is equally attracted by the two atoms. As a result electron pair is situated exactly between the two identical nuclei. The bond so formed is called nonpolar covalent bond. • When covalent bond is formed between two atoms of different elements (heteronuclear molecule) like HF, the shared electron pair between the two atoms gets displaced more towards more electronegative atom (fluorine).The resultant covalent bond is a polar covalent bond. • As a result of polarisation, the molecule possesses the dipole moment
Covalent Bond, Polar covalent bond and Ionic Bond
Dipole moment • The product of the magnitude of the charge and the distance between the centres of positive and negative charge is known as dipole moment. • It is usually designated by a Greek letter ‘μ’ (Mu). • Mathematically, • Dipole moment (μ) = charge (Q) × distance of separation (r) • Dipole moment is usually expressed in Debye units (D). • 1 D = 3.33564 × 10–30 C m where C is coulomb and m is meter.
Continue… Dipole moment • Dipole moment is a vector quantity and by convention it is depicted by a small arrow with tail on the negative centre and head pointing towards the positive centre. • But in chemistry presence of dipole moment is represented by the crossed arrow put on Lewis structure of the molecule. • The cross is on positive end and arrow head is on negative end. • For example This arrow symbolises the direction of the shift of electron density in the molecule
Dipole moment and geometry • In case of polyatomic molecules the dipole moment depends not only on bond dipoles but also on the spatial arrangement of various bonds in the molecule. • In such case, the dipole moment of a molecule is the vector sum of the dipole moments of various bonds. Bond dipole:- • Dipole moment of individual bond • For example H2O molecule, bent , bond angle of 104.5o . • Net dipole moment =1.85 D • It is the resultant of the dipole moments of two O–H bonds.
Continue… Dipole moment and geometry • The dipole moment in case of BeF2 is zero. • This is because the two equal bond dipoles point in opposite directions and cancel the effect of each other. • In BF3, the dipole moment is zero although the B – F bonds are oriented at an angle of 120o to one another, the three bond moments give a net sum of zero as the resultant of any two is equal and opposite to the third.
Dipole Moment in NH3 and NF3 molecule. • Both have pyramidal shape with a lone pair of electrons on nitrogen atom. • Dipole moment of NH3 ( 4.90 × 10–30 C m) is greater than that of NF3 (0.8 × 10–30 C m). In case of NH3 the orbital dipole due to lone pair is in the same direction as the resultant dipole moment of the N – H bonds, whereas in NF3 the orbital dipole is in the direction opposite to the resultant dipole moment of the three N–F bonds. Fluorine is more electronegative than nitrogen
Covalent Character in Ionic Bond • The cation polarises the anion, pulling the electronic charge toward itself and thereby increasing the electronic charge between the two. • In a covalent bond, electron charge density is in between the nuclei. • Factors affecting the percent covalent character of the ionic bond:- • The polarising power of the cation, • the polarisability of the anion and • the extent of distortion (polarisation) of anion
Continue… Covalent Character in Ionic Bond • Polarising power • It is the extent to which a cation can polarise an anion. • It is proportional to charge density. • Charge density is the ratio of charge to volume. • Polarising power α Charge density. • More the charge density, greater is the polarising power for that cation. • Polarisability • It is the extent to which an anion can be polarised. • It can also be called as the ease with which an ion can be polarised. • Polarisation is the distortion of a spherically symmetric electron cloud to an unsymmetric cloud.
Fajans Rule • The partial covalent character of ionic bonds was discussed by Fajans in terms of the following rules: • Size of the ion: The smaller the size of the cation and the larger the size of the anion, the greater the covalent character of an ionic bond. • The charge of Cation: The greater the charge on the cation, the greater the covalent character of the ionic bond. • Electronic Configuration: For cations of the same size and charge, the one, with electronic configuration (n-1)dn nso , typical of transition metals, is more polarising than the one with a noble gas configuration, ns2 np6 , typical of alkali and alkaline earth metal cations.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • The Valence Shell Electron Pair Repulsion (VSEPR) Theory • Shapes of Molecules in which the central atom has no lone pair By-Vijay Kumar Sethi
• This theory provides a simple procedure to predict the shapes of covalent molecules. • Proposed by:-Sidgwick and Powell in 1940, • Based on the repulsive interactions of the electron pairs in the valence shell of the atoms. • Further developed and redefined by Nyholm and Gillespie (1957). The main postulates of VSEPR theory:- 1) The shape of a molecule depends upon the number of valence shell electron pairs (bonded or nonbonded) around the central atom. 2) Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged. The Valence Shell Electron Pair Repulsion (VSEPR) Theory
Continue… The main postulates of VSEPR theory:- 3) These pairs of electrons tend to occupy such positions in space that minimize 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 from one another. 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. 6) Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure
• The repulsive interaction of electron pairs decrease in the order: • lp – lp > lp –bp > bp – bp • lone pairs are localised on the central atom and each bonded pair is shared between two atoms. • As a result, the lone pair electrons in a molecule occupy more space as compared to the bonding pairs of electrons. • This results in greater repulsion between lone pairs of electrons as compared to the lone pair - bond pair and bond pair - bond pair repulsions. lp= Lone pair bp=Bond pair
Geometry is regular- • linear, • trigonal planar, • tetrahedral, • Trigonal bipyramidal and • octahedral Molecules in which the central atom has no lone pair
Continue… Molecules in which the central atom has no lone pair
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Shapes of Molecules in which the central atom has one or more lone pairs- According to VSEPR Theory By-Vijay Kumar Sethi
Molecules in which the central atom has one or more lone pairs Regular Geometry - • linear, • trigonal planar, • tetrahedral, • Trigonal bipyramidal • octahedral
Continue… Molecules in which the central atom has one or more lone pairs
Continue… Molecules in which the central atom has one or more lone pairs
The Effect of Nonbonding Electrons • Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. • Therefore, the bond angle decreases as the number of lone pairs increases
Some Specific Examples • electron pairs of a multiple bond are treated as a single super pair. • lone pair- bond pair repulsion is much more as compared to the bond pair-bond pair repulsion. • So the angle is reduced to 119.5° from 120°.
Continue… Some Specific Examples • In (a) the lp is present at axial position so there are three lp—bp repulsions at 90°. • In(b) the lp is in an equatorial position, and there are two lp—bp repulsions. • Hence, arrangement (b) is more stable. • The shape shown in (b) is described as a distorted tetrahedron, a folded square or a see-saw. • If bond angle is more than 100o , repulsive interactions are not significant
Continue… Some Specific Examples • In (a) the lp are at equatorial position so there are less lp-bp repulsions as compared to others in which the lp are at axial positions. • So structure (a) is most stable. (T-shaped).
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Valence bond theory(VBT) By-Vijay Kumar Sethi
Valence bond theory • Introduced by Heitler and London (1927) and developed further by Pauling and others. • based on the • knowledge of atomic orbitals, • Electronic configurations of elements , • The overlap criteria of atomic orbitals, • The hybridization of atomic orbitals and the • principles of variation and superposition • A covalent bond is formed by the overlapping of atomic orbitals. • Before overlapping atomic orbitals undergo hybridization. • A bond is formed when attractive interactions balance the repulsive interactions and net energy is lowered
Continue… Valence bond theory • consider the formation of hydrogen molecule • Consider two hydrogen atoms A and B approaching each other having nuclei NA and NB and electrons present in them are represented by eA and eB. • Attractive forces arise between: • nucleus of one atom and its own electron that is NA – eA and NB– eB. • nucleus of one atom and electron of other atom i.e., NA– eB, NB– eA. • Similarly repulsive forces arise between • electrons of two atoms like eA – eB, • nuclei of two atoms NA – NB.
• Attractive forces tend to bring the two atoms close to each other whereas repulsive forces tend to push them apart Continue… consider the formation of hydrogen molecule
• Magnitude of new attractive force is more than the new repulsive forces. • As a result, two atoms approach each other and potential energy decreases. • Ultimately a stage is reached where the net force of attraction balances the force of repulsion and system acquires minimum energy. • At this stage two hydrogen atoms are said to be bonded together to form a stable molecule having the bond length of 74 pm. • Since the energy gets released when the bond is formed between two hydrogen atoms, the hydrogen molecule is more stable than that of isolated hydrogen atoms. • The energy so released is called as bond enthalpy Continue… consider the formation of hydrogen molecule
When H – H distance = 74 pm, Repulsion = Attraction → strongest bond → optimal overlap → lowest energy → bond enthalpy -435.8 kJ 74 pm
At H – H distance > 74 pm, Repulsion < Attraction → weaker bond → too little overlap → atoms come closer 74 pm > 74 pm
At H – H distance < 74 pm, Repulsion > Attraction → weaker bond → too much overlap → atoms get further apart 74 pm < 74 pm
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Overlapping of Atomic Orbitals • Sigma and pi Bonds By-Vijay Kumar Sethi
Orbital Overlap Concept • In the formation of hydrogen molecule, there is a minimum energy state when two hydrogen atoms are so near that their atomic orbitals undergo partial interpenetration. • The 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. • In general, greater the overlap the stronger is the bond formed between two atoms. • Therefore, according to orbital overlap concept, 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 orbitals of two atoms come close to form bond, their overlap may be positive, negative or zero depending upon the sign (phase) and direction of orientation of amplitude of orbital wave function in space • Positive and negative sign on boundary surface diagrams show the sign (phase) of orbital wave function and are not related to charge. • Orbitals forming bond should have same sign (phase) and orientation in space. • This is called positive overlap.
Overlapping of Atomic Orbitals • A negative overlap results when two orbitals with different phase overlap with each other. • A zero overlap will results when orbitals do not overlap at all or do not overlap effectively. • A negative and zero overlap of orbitals does not result in bond formation
Nature of Covalent Bonds • The covalent bond may be classified into two types depending upon the types of overlapping: (i) Sigma(σ) bond, and (ii) pi(π) bond (i) Sigma(σ) bond : • This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. This is called as head on overlap or axial overlap. • In this the overlapping of orbitals takes place to a larger extent. Hence, it is stronger • It may be formed by s-s overlapping, s-p overlapping, p–p overlapping • In this electron density is concentrated in between the nuclei of bonded atoms.
Continue… (i) Sigma(σ) bond : s-s overlapping : In this case, there is overlap of two half filled s-orbitals along the internuclear axis
Continue… (i) Sigma(σ) bond : s-p overlapping: This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom.
Continue… (i) Sigma(σ) bond : p–p overlapping : This type of overlap takes place between half filled p- orbitals of the two approaching atoms.
(ii) pi(π ) bond : • In the formation of π bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. • This bond is formed by sidewise or lateral overlapping of atomic orbitals. • In this bond electron charged clouds is above and below the plane of the participating atoms. • Extent of overlapping is smaller so it is weaker.
Differences between Sigma and pi Bonds
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • HYBRIDISATION By-Vijay Kumar Sethi
Why?
• The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised. • The hybridised orbitals are always equivalent in energy and shape. • The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. • These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement. the type of hybridisation indicates the geometry of the molecules Salient features of hybridisation
• The orbitals present in the valence shell and penultimate shell of the atom are hybridised. • The orbitals undergoing hybridisation should have almost equal energy. • Promotion of electron is not essential condition prior to hybridisation. • It is not necessary that only half filled orbitals participate in hybridisation. In some cases, even filled orbitals and empty orbitals of valence shell and penultimate shell take part in hybridisation. Important conditions for hybridisation
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • HYBRIDISATION-Involving s and p orbitals By-Vijay Kumar Sethi
• The molecule thus gets distorted and the bond angle is reduced to 107° from 109.5°. • The geometry is pyramidal
• The bond angle in this case is reduced to 104.5° from 109.5° Repulsion lp-lp > lp-bp > bp-bp • Water molecule has V-shape or angular geometry or bent shape.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Hybridisation of Carbon in Ethane, Ethene, Ethyne By-Vijay Kumar Sethi
Hybridisation in Ethane C2H6 molecule • Each C is sp3 hybrid • C-C bond , sp3 -sp3 sigma bond . • Each C-H sp3 –s sigma bonds . • In ethane C–C bond length is 154 pm and each C–H bond length is 109 pm.
Hybridisation in Ethene C2H4 molecule • Each C is sp2 hybrid • C-C bond , sp2 -sp2 sigma bond . • Each C-H sp2 –s sigma bonds .
Continue… Hybridisation in Ethene C2H4 molecule • The unhybridised orbital (2px or 2py) of one carbon atom overlaps sidewise with the similar orbital of the other carbon atom to form weak π bond, which consists of two equal electron clouds distributed above and below the plane of carbon and hydrogen atoms. C=C bond length 134 pm. C–H bond length 108 pm. H–C–H bond angle is 117.6° H–C–C angle is 121°.
Hybridisation in Ethyne C2H2 molecule • Each C is sp hybrid • C-C bond , sp-sp sigma bond . • Each C-H sp –s sigma bonds . • Each of the two unhybridised p orbitals of both the carbon atoms overlaps sidewise to form two π bonds between the carbon atoms.
Continue… Hybridisation in Ethyne C2H2 molecule • The triple bond between the two carbon atoms is made up of one sigma and two pi bonds C≡C bond length 120 pm. C–H bond length 106 pm.
Summarization
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Hybridisation of Elements involving d Orbitals • Formation of PCl5 and SF6 By-Vijay Kumar Sethi
Formation of PCl5 (sp3 d hybridisation): Five sp3 d hybrid orbitals are directed towards the five corners of a trigonal bipyramidal
Continue… Formation of PCl5 (sp3 d hybridisation): • All the bond angles in trigonal bipyramidal geometry are not equivalent. • Three P–Cl bond lie in one plane and make an angle of 120° with each other; these bonds are termed as equatorial bonds. • The remaining two P–Cl bonds–one lying above and the other lying below the equatorial plane, make an angle of 90° with the plane. These bonds are called axial bonds. As the axial bond pairs suffer more repulsive interaction from the equatorial bond pairs, therefore axial bonds are slightly longer and hence slightly weaker than the equatorial bonds; which makes PCl5 molecule more reactive.
Formation of SF6 (sp3 d2 hybridisation) sp3 d2 hybrid orbitals are projected towards the six corners of a regular octahedron in SF6.
• These six sp3 d2 hybrid orbitals overlap with singly occupied orbitals of fluorine atoms to form six S–F sigma bonds. • Thus SF6 molecule has a regular octahedral geometry Continue… Formation of SF6 (sp3 d2 hybridisation)
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Molecular Orbital Theory(MOT) • Linear Combination of Atomic Orbitals (LCAO) By-Vijay Kumar Sethi
MOLECULAR ORBITAL THEORY • Developed by F. Hund and R.S. Mulliken in 1932. • The salient features of this theory are : 1) The electrons in a molecule are present in the various molecular orbitals. 2) The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals. 3) Molecular orbital is polycentric. The electron in a molecular orbital is influenced by two or more nuclei depending upon the number of atoms in the molecule
Continue… MOLECULAR ORBITAL THEORY 4) The number of molecular orbital formed is equal to the number of combining atomic orbitals. When two atomic orbitals combine, two molecular orbitals are formed. One is known as bonding molecular orbital while the other is called antibonding molecular orbital. 5) The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital. 6) The electron probability distribution around a group of nuclei in a molecule is given by a molecular orbital. 7) The molecular orbitals like atomic orbitals are filled in accordance with the aufbau principle obeying the Pauli’s exclusion principle and the Hund’s rule
Formation of Molecular Orbitals Linear Combination of Atomic Orbitals (LCAO) • This is an approximate method to obtain wave functions for molecular orbitals. • Consider hydrogen molecule consisting of two atoms of hydrogen ,A and B. • Each hydrogen atom in the ground state has one electron in 1s orbital. • The atomic orbitals of these atoms may be represented by the wave functions ψA and ψB. • Mathematically, the formation of molecular orbitals may be described by the linear combination of atomic orbitals that can take place by addition and by subtraction of wave functions of individual atomic orbitals as shown :
Continue… Linear Combination of Atomic Orbitals (LCAO) • The two molecular orbitals σ and σ* are formed as : • The molecular orbital σ formed by the addition of atomic orbitals is called the bonding molecular orbital while • the molecular orbital σ* formed by the subtraction of atomic orbital is called antibonding molecular orbital
Bonding Molecular Orbital Anti-bonding Molecular Orbital Formed by constructive interference of the electron waves of the combining atoms. Formed by destructive interference of the electron waves of the combining atoms. The electron density is located between the nuclei of the bonded atoms because of which the repulsion between the nuclei is very less Most of the electron density is located away from the space between the nuclei hence the repulsion between the nuclei is high. Electrons placed in a bonding molecular orbital tend to hold the nuclei together and stabilise a molecule Electrons placed in the antibonding molecular orbital destabilise the molecule. Possesses lower energy than either of the atomic orbitals that have combined to form it. Possesses higher energy than either of the atomic orbitals that have combined to form it. It does not has nodal plane It has is a nodal plane
Conditions for the Combination of Atomic Orbitals 1. The combining atomic orbitals must have the same or nearly the same energy. 2. The combining atomic orbitals must have the same symmetry about the molecular axis. By convention z-axis is taken as the molecular axis. • For example, 2pz orbital of one atom can combine with 2pz orbital of the other atom but not with the 2px or 2py orbitals because of their different symmetries. 3. The combining atomic orbitals 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.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Types of Molecular Orbitals By-Vijay Kumar Sethi
Types of Molecular Orbitals • Molecular orbitals of diatomic molecules are designated as σ (sigma), π (pi), δ (delta), etc. • The sigma (σ) molecular orbitals are symmetrical around the bond-axis while pi (π) molecular orbitals are not symmetrical.
• For example, the linear combination of 1s orbitals centered on two nuclei produces two molecular orbitals which are symmetrical around the bond- axis. • Such molecular orbitals are of the σ type and are designated as σ1s and σ*1s Continue… Types of Molecular Orbitals
Continue… Types of Molecular Orbitals • If internuclear axis is taken to be in the z-direction, a linear combination of 2pz- orbitals of two atoms also produces two sigma molecular orbitals designated as σ2pz and σ*2pz.
Continue… Types of Molecular Orbitals • Molecular orbitals obtained from 2px and 2py orbitals are not symmetrical around the bond axis because of the presence of positive lobes above and negative lobes below the molecular plane. • Such molecular orbitals, are labelled as π and π* . (π 2px , π* 2px and π 2py , π* 2py) • A π bonding MO has larger electron density above and below the inter -nuclear axis. The π* antibonding MO has a node between the nuclei.
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Energy Level Diagram for Molecular Orbitals By-Vijay Kumar Sethi
Energy Level Diagram for Molecular Orbitals • 1s atomic orbitals on two atoms form two molecular orbitals designated as σ1s and σ*1s. • 2s and 2p atomic orbitals (eight atomic orbitals on two atoms) give rise to the following eight molecular orbitals: • The increasing order of energies of various molecular orbitals for O2 and F2 is : • For molecules such as B2, C2, N2, etc. the increasing order of energies is
for O2 and F2 For B2, C2, N2
Electronic configuration of the molecule • The distribution of electrons among various molecular orbitals is called the electronic configuration of the molecule Stability of Molecules: • If Nb is the number of electrons occupying bonding orbitals and Na the number occupying the antibonding orbitals, then • (i) the molecule is stable if Nb is greater than Na, and • (ii) the molecule is unstable if Nb is less than Na.
Bond order • Bond order (b.o.) is defined as one half the difference between the number of electrons present in the bonding and the antibonding orbitals i.e., Bond order (b.o.) = ½ (Nb–Na) • Nb > Na Positive bond order means a stable molecule • Nb< Na Negative bond order , unsatble molecule • Nb = Na Zero bond order , unstable molecule. • Integral bond order values of 1, 2 or 3 correspond to single, double or triple bonds respectively • The bond length decreases as bond order increases.
Magnetic nature • If all the molecular orbitals in a molecule are doubly occupied, the substance is diamagnetic (repelled by magnetic field). No unpaired electron • However if one or more molecular orbitals are singly occupied it is paramagnetic (attracted by magnetic field), e.g., O2molecule. Presence of unpaired electron
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Bonding In Some Homonuclear Diatomic Molecules as per MOT By-Vijay Kumar Sethi
Hydrogen molecule (H2 ) electronic configuration of H2:- (σ1s)2 Magnetic property:- Diamagnetic ( no unpaired electron)
Helium molecule (He2 ): • electronic configuration of He2:- (σ1s)2 (σ*1s)2 • He2 molecule does not exist. • Similarly, it can be shown that Be2 molecule also does not exist. (σ1s)2 (σ*1s)2 (σ2s)2 (σ*2s)2 Bond order = 0
Lithium molecule (Li2 ): • The electronic configuration of lithium is 1s2 2s1 . • There are six electrons in Li2. • The electronic configuration of Li2 molecule is • Li2 : (σ1s)2 (σ*1s)2 (σ2s)2 • The above configuration is also written as KK(σ2s)2 where KK represents the closed K shell structure (σ1s)2 (σ*1s)2 . • Li2 Diamagnetic and stable. • Li2 molecules are known to exist in the vapour phase.
Carbon molecule (C2 ): • The electronic configuration of carbon is 1s2 2s2 2p2 . • There are twelve electrons in C2. • The electronic configuration of C2 molecule is C2 diamagnetic, detected in vapour phase. Double bond in C2 consists of both pi bonds because of the presence of four electrons in two pi molecular orbitals.
Oxygen molecule (O2 ): • The electronic configuration of oxygen atom is 1s2 2s2 2p4 . each atom 8 electrons • There are 16 electrons in O2molecule.. • The electronic configuration of O2 molecule is
Oxygen molecule (O2 ): it contains two unpaired electrons in p * 2px and p * 2py molecular orbitals, therefore it is paramagnetic
Class XI Chemistry UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • HYDROGEN BONDING By-Vijay Kumar Sethi
HYDROGEN BONDING • Nitrogen, oxygen and fluorine are the highly electronegative elements. • When they are attached to a hydrogen atom to form covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. • This partially positively charged hydrogen atom of one molecule forms a bond with the more electronegative atom of other molecule. This bond is known as hydrogen bond and is weaker than the covalent bond. • For example HF • Hydrogen bond is represented by a dotted line (– – –) while a solid line represents the covalent bond. • Thus, hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule.
Cause of Formation of Hydrogen Bond • An electrostatic force of attraction between hydrogen atom of one molecule and electronegative atom of second molecule. Conditions for the Formation of Hydrogen Bond • (1) Hydrogen atom should be bonded to highly electronegative atom say F, O orN. (2) The electronegative atoms should be small in size. • For example in HF and HCl , H atom is bonded to more electronegative atom F and Cl respectively but H-bond exist in HF not in HCl because size of Cl is larger
Types of H-Bonds There are two types of H-bonds (i) Intermolecular hydrogen bond (ii) Intramolecular hydrogen bond (1) Intermolecular hydrogen bond : • It is formed between two different molecules of the same or different compounds. • For example, H-bond in case of HF molecule, alcohol or water molecules, etc.
(2) Intramolecular hydrogen bond : • It is formed when hydrogen atom is in between the two highly electronegative (F, O, N) atoms present within the same molecule. • For example, in o-nitrophenol the hydrogen is in between the two oxygen atoms.

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class11chapter4chemicalbondingandmolecularstructure-231220155442-2e63aebc.pptx

  • 1.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Octet Rule • Covalent Bond • Limitations of Octet Rule By-Vijay Kumar Sethi
  • 2.
    Chemical bond • Theattractive force which holds various constituents (atoms, ions, etc.) together in different chemical species is called a chemical bond. Valence electrons:-Electrons present in the outermost shell are known as valence electrons Lewis symbols:-Valence electrons of an atom are represented by dots around the symbol of the element Significance of Lewis Symbols • Valence electrons helps to calculate the common or group valence of the element. • The group valence of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.
  • 3.
    Octet Rule • Theoryof chemical combination between atoms is known as electronic theory of chemical bonding. • It was developed by Kössel and Lewis in 1916. According to this, • atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. • This is known as octet rule.
  • 4.
    Covalent Bond • Lewispostulated that atoms achieve the stable octet when they are linked by chemical bonds. • Langmuir introduced the term covalent bond. • According to Lewis-Langmuir theory a covalent bond is formed by sharing a pair of electron between two bonded atoms
  • 6.
    Single covalent bond •If covalent bond is formed by sharing a pair of electron it is called a single covalent bond and it is represented by a line between the two atoms Double Bond:- If two atoms share two pairs of electrons Triple Bond:- If two atoms share three pairs of electrons • Multiple bonds- double and triple bonds collectively known as multiple bonds
  • 7.
    Limitations of theOctet Rule 1) The incomplete octet of the central atom:- • This is especially the case with elements having less than four valence electrons. • Examples are LiCl, BeH2 and BCl3. • Li, Be and B have 1,2 and 3 valence electrons only. • Some other such compounds are AlCl3 and BF3.
  • 8.
    Continue… Limitations of theOctet Rule 2) Odd-electron molecules • In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO2, the octet rule is not satisfied for the N atoms
  • 9.
    Continue… Limitations of theOctet Rule 3) The expanded octet:- • Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. • In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet. sulphur also forms many compounds in which the octet rule is obeyed. For example sulphur dichloride SCl2
  • 10.
    Continue… Limitations of theOctet Rule 4) octet rule is based upon the chemical inertness of noble gases. However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like XeF2, KrF2, XeOF2 etc., 5) This theory does not account for the shape of molecules. 6) It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.
  • 11.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Lewis dot structures • Formal Charge By-Vijay Kumar Sethi
  • 12.
    Lewis dot structures •The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule
  • 13.
    Writing of Lewisdot structures The Lewis dot structures can be written by adopting the following steps: Step 1:- Count the total number of valence electrons. For anions, add one electron for each negative charge and for cation subtract one electron for each positive charge. Step 2:- Write skeletal structure (In general the least electronegative atom occupies the central position) Step 3:-Draw a single bond by one shared electron pair between bonded atoms. Remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs. If necessary rearrange the electrons as multiple bonds or lone pairs to complete the octet.
  • 14.
    Problem 4.1 Writethe Lewis dot structure of CO molecule. Solution • Step 1. Count the total number of valence electrons of carbon and oxygen atoms. The valence electrons available are 4 + 6 =10. • Step 2. The skeletal structure of CO is written as: C O • Step 3. Draw a single bond (one shared electron pair) between C and O and complete the octet on O, the remaining two electrons are the lone pair on C. • This does not complete the octet on carbon and hence we have to resort to multiple bonding (triple bond) between C and O atoms.
  • 16.
    Formal Charge • Incase of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom. • It is, however, feasible to assign a formal charge on each atom. • The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure.
  • 17.
    Continue… Formal Charge:- Considerthe ozone molecule (O3). The Lewis structure of O3 may be drawn as • Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species. • Generally the lowest energy structure is the one with the smallest formal charges on the atoms.
  • 18.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Ionic Or Electrovalent Bond • Lattice Enthalpy By-Vijay Kumar Sethi
  • 19.
    Ionic Or ElectrovalentBond Kössel, drew attention to the following facts: ❖ In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases; ❖ The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of an electron by the respective atoms; ❖ The negative and positive ions thus formed attain stable noble gas electronic configurations. The noble gases (with the exception of helium which has a duplet of electrons) have a particularly stable outer shell configuration of eight (octet) electrons, ns2 np6 . ❖ The negative and positive ions are stabilized by electrostatic attraction.
  • 20.
    Continue… Ionic Or ElectrovalentBond • For example, the formation of NaCl • Formation of CaF2 The bond formed, as a result of the electrostatic attraction between the positive and negative ions is termed as the electrovalent bond. The electrovalence is equal to the number of unit charge(s) on the ion. Thus, calcium is assigned a positive electrovalence of two, while chlorine a negative electrovalence of one.
  • 21.
    Ionic Bonds: OneBig Greedy Thief Dog!
  • 22.
    Continue… Ionic Or ElectrovalentBond ❑ The formation of ionic compounds would depend upon: ❑ The ease of formation of the positive and negative ions from the respective neutral atoms; ❑ The arrangement of the positive and negative ions in the solid, that is, the lattice of the crystalline compound. • Ionic bonds will be formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy. • Lattice enthalpy plays a key role in the formation of ionic compounds
  • 23.
    Lattice Enthalpy • TheLattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. • For example, the lattice enthalpy of NaCl is 788 kJ mol–1 . This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+ (g) and one mole of Cl– (g) to an infinite distance. NaCl(s) → Na+ (g) + Cl– (g) • The solid crystal being three dimensional; it is not possible to calculate lattice enthalpy directly from the interaction of forces of attraction and repulsion only. It is determined indirectly using Born-Haber Cycle .
  • 24.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Bond Parameters-Bond Length, Bond Angle, Bond Enthalpy, Bond Order By-Vijay Kumar Sethi
  • 25.
    Bond Parameters Bond Length •Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. • Bond lengths are measured by spectroscopic, X-ray diffraction and electron-diffraction techniques • Each atom of the bonded pair contributes to the bond length . • In the case of a covalent bond, the contribution from each atom is called the covalent radius of that atom.
  • 26.
    • The covalentradius is half of the distance between two similar atoms joined by a covalent bond in the same molecule. • The van der Waals radius is half of the distance between two similar atoms which are not bonded Covalent radius and van der Waals radius
  • 27.
    Continue… Bond Parameters Bond Angle •It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. • Bond angle is expressed in degree which can be experimentally determined by spectroscopic methods. • It helps in determining the shape of molecule/complex ion. • For example H–O–H bond angle in water is 104.5o
  • 28.
    Continue… Bond Parameters Bond Enthalpy •It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. • The unit of bond enthalpy is kJ mol–1 . Larger the bond dissociation enthalpy, stronger will be the bond in the molecule. Enthalpy of atomization
  • 29.
    Continue… Bond Parameters Bond Enthalpy •In case of polyatomic molecules like water H2O • Two O-H bond have different values of bond enthalpy because of changed chemical environment. • Therefore in polyatomic molecules the term mean or average bond enthalpy is used. • It is obtained by dividing total bond dissociation enthalpy by the number of bonds broken
  • 30.
    Molecule No. ofShared electron pairs Number of covalent bonds Bond Order Hydrogen (H2) One 1 1 Oxygen (O2) Two 2 2 Nitrogen (N2) Three 3 3 Carbon monoxide (CO) Three 3 3
  • 32.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Concept of Resonance By-Vijay Kumar Sethi
  • 33.
    Resonance Structures • Asingle Lewis structure of a molecule is not sufficient to describe its experimentally determined parameters. For example, Ozone O3. • The normal O–O and O=O bond lengths are 148 pm and 121 pm respectively. • Experimentally determined oxygen-oxygen bond lengths in the O3 molecule are same (128 pm).
  • 34.
    Continue… Resonance Structures • Accordingto the concept of resonance , whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately. • Resonance is represented by a double headed arrow.
  • 36.
    Problem 4.4:- Explainthe structure of CO2 molecule. Solution:- • The experimentally determined carbon to oxygen bond length in CO2 is 115 pm. • The lengths of a normal (C=O) and (C≡O) are 121 pm and 110 pm respectively. • So a single Lewis structure cannot depict this position and it becomes necessary to write more than one Lewis structures and to consider that the structure of CO2 is best described as a hybrid of the canonical or resonance forms I, II and III.
  • 37.
    Continue… Resonance Structures • Ingeneral, it may be stated that ❖ Resonance stabilizes the molecule as the energy of the resonance hybrid is less than the energy of any single canonical structure; and, ❖ Resonance averages the bond characteristics as a whole. Resonance Energy:- • Energy difference between energy of the resonance hybrid and energy of the most stable canonical structure is known as resonance energy
  • 38.
    Continue… Resonance Structures Remember that: ❖ The canonical forms have no real existence. ❖ The molecule does not exist for a certain fraction of time in one canonical form and for other fractions of time in other canonical forms. ❖ There is no equilibrium between the canonical . ❖ The molecule as such has a single structure which is the resonance hybrid of the canonical forms and which cannot as such be depicted by a single Lewis structure.
  • 39.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Polar covalent bond • Dipole moment • Covalent Character in Ionic Bond-Fajans Rule By-Vijay Kumar Sethi
  • 40.
    Polarity of Bonds-Polar covalent bond • When covalent bond is formed between two similar atoms, for example in H2, O2, Cl2, N2 or F2, the shared pair of electrons is equally attracted by the two atoms. As a result electron pair is situated exactly between the two identical nuclei. The bond so formed is called nonpolar covalent bond. • When covalent bond is formed between two atoms of different elements (heteronuclear molecule) like HF, the shared electron pair between the two atoms gets displaced more towards more electronegative atom (fluorine).The resultant covalent bond is a polar covalent bond. • As a result of polarisation, the molecule possesses the dipole moment
  • 41.
    Covalent Bond, Polarcovalent bond and Ionic Bond
  • 42.
    Dipole moment • Theproduct of the magnitude of the charge and the distance between the centres of positive and negative charge is known as dipole moment. • It is usually designated by a Greek letter ‘μ’ (Mu). • Mathematically, • Dipole moment (μ) = charge (Q) × distance of separation (r) • Dipole moment is usually expressed in Debye units (D). • 1 D = 3.33564 × 10–30 C m where C is coulomb and m is meter.
  • 43.
    Continue… Dipole moment • Dipolemoment is a vector quantity and by convention it is depicted by a small arrow with tail on the negative centre and head pointing towards the positive centre. • But in chemistry presence of dipole moment is represented by the crossed arrow put on Lewis structure of the molecule. • The cross is on positive end and arrow head is on negative end. • For example This arrow symbolises the direction of the shift of electron density in the molecule
  • 44.
    Dipole moment andgeometry • In case of polyatomic molecules the dipole moment depends not only on bond dipoles but also on the spatial arrangement of various bonds in the molecule. • In such case, the dipole moment of a molecule is the vector sum of the dipole moments of various bonds. Bond dipole:- • Dipole moment of individual bond • For example H2O molecule, bent , bond angle of 104.5o . • Net dipole moment =1.85 D • It is the resultant of the dipole moments of two O–H bonds.
  • 45.
    Continue… Dipole moment andgeometry • The dipole moment in case of BeF2 is zero. • This is because the two equal bond dipoles point in opposite directions and cancel the effect of each other. • In BF3, the dipole moment is zero although the B – F bonds are oriented at an angle of 120o to one another, the three bond moments give a net sum of zero as the resultant of any two is equal and opposite to the third.
  • 46.
    Dipole Moment inNH3 and NF3 molecule. • Both have pyramidal shape with a lone pair of electrons on nitrogen atom. • Dipole moment of NH3 ( 4.90 × 10–30 C m) is greater than that of NF3 (0.8 × 10–30 C m). In case of NH3 the orbital dipole due to lone pair is in the same direction as the resultant dipole moment of the N – H bonds, whereas in NF3 the orbital dipole is in the direction opposite to the resultant dipole moment of the three N–F bonds. Fluorine is more electronegative than nitrogen
  • 47.
    Covalent Character inIonic Bond • The cation polarises the anion, pulling the electronic charge toward itself and thereby increasing the electronic charge between the two. • In a covalent bond, electron charge density is in between the nuclei. • Factors affecting the percent covalent character of the ionic bond:- • The polarising power of the cation, • the polarisability of the anion and • the extent of distortion (polarisation) of anion
  • 48.
    Continue… Covalent Character inIonic Bond • Polarising power • It is the extent to which a cation can polarise an anion. • It is proportional to charge density. • Charge density is the ratio of charge to volume. • Polarising power α Charge density. • More the charge density, greater is the polarising power for that cation. • Polarisability • It is the extent to which an anion can be polarised. • It can also be called as the ease with which an ion can be polarised. • Polarisation is the distortion of a spherically symmetric electron cloud to an unsymmetric cloud.
  • 49.
    Fajans Rule • Thepartial covalent character of ionic bonds was discussed by Fajans in terms of the following rules: • Size of the ion: The smaller the size of the cation and the larger the size of the anion, the greater the covalent character of an ionic bond. • The charge of Cation: The greater the charge on the cation, the greater the covalent character of the ionic bond. • Electronic Configuration: For cations of the same size and charge, the one, with electronic configuration (n-1)dn nso , typical of transition metals, is more polarising than the one with a noble gas configuration, ns2 np6 , typical of alkali and alkaline earth metal cations.
  • 50.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • The Valence Shell Electron Pair Repulsion (VSEPR) Theory • Shapes of Molecules in which the central atom has no lone pair By-Vijay Kumar Sethi
  • 51.
    • This theoryprovides a simple procedure to predict the shapes of covalent molecules. • Proposed by:-Sidgwick and Powell in 1940, • Based on the repulsive interactions of the electron pairs in the valence shell of the atoms. • Further developed and redefined by Nyholm and Gillespie (1957). The main postulates of VSEPR theory:- 1) The shape of a molecule depends upon the number of valence shell electron pairs (bonded or nonbonded) around the central atom. 2) Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged. The Valence Shell Electron Pair Repulsion (VSEPR) Theory
  • 52.
    Continue… The main postulatesof VSEPR theory:- 3) These pairs of electrons tend to occupy such positions in space that minimize 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 from one another. 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. 6) Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure
  • 53.
    • The repulsiveinteraction of electron pairs decrease in the order: • lp – lp > lp –bp > bp – bp • lone pairs are localised on the central atom and each bonded pair is shared between two atoms. • As a result, the lone pair electrons in a molecule occupy more space as compared to the bonding pairs of electrons. • This results in greater repulsion between lone pairs of electrons as compared to the lone pair - bond pair and bond pair - bond pair repulsions. lp= Lone pair bp=Bond pair
  • 54.
    Geometry is regular- •linear, • trigonal planar, • tetrahedral, • Trigonal bipyramidal and • octahedral Molecules in which the central atom has no lone pair
  • 55.
    Continue… Molecules in whichthe central atom has no lone pair
  • 57.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Shapes of Molecules in which the central atom has one or more lone pairs- According to VSEPR Theory By-Vijay Kumar Sethi
  • 58.
    Molecules in whichthe central atom has one or more lone pairs Regular Geometry - • linear, • trigonal planar, • tetrahedral, • Trigonal bipyramidal • octahedral
  • 59.
    Continue… Molecules in whichthe central atom has one or more lone pairs
  • 60.
    Continue… Molecules in whichthe central atom has one or more lone pairs
  • 61.
    The Effect ofNonbonding Electrons • Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. • Therefore, the bond angle decreases as the number of lone pairs increases
  • 62.
    Some Specific Examples •electron pairs of a multiple bond are treated as a single super pair. • lone pair- bond pair repulsion is much more as compared to the bond pair-bond pair repulsion. • So the angle is reduced to 119.5° from 120°.
  • 63.
    Continue… Some Specific Examples •In (a) the lp is present at axial position so there are three lp—bp repulsions at 90°. • In(b) the lp is in an equatorial position, and there are two lp—bp repulsions. • Hence, arrangement (b) is more stable. • The shape shown in (b) is described as a distorted tetrahedron, a folded square or a see-saw. • If bond angle is more than 100o , repulsive interactions are not significant
  • 64.
    Continue… Some Specific Examples •In (a) the lp are at equatorial position so there are less lp-bp repulsions as compared to others in which the lp are at axial positions. • So structure (a) is most stable. (T-shaped).
  • 65.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Valence bond theory(VBT) By-Vijay Kumar Sethi
  • 66.
    Valence bond theory •Introduced by Heitler and London (1927) and developed further by Pauling and others. • based on the • knowledge of atomic orbitals, • Electronic configurations of elements , • The overlap criteria of atomic orbitals, • The hybridization of atomic orbitals and the • principles of variation and superposition • A covalent bond is formed by the overlapping of atomic orbitals. • Before overlapping atomic orbitals undergo hybridization. • A bond is formed when attractive interactions balance the repulsive interactions and net energy is lowered
  • 67.
    Continue… Valence bond theory •consider the formation of hydrogen molecule • Consider two hydrogen atoms A and B approaching each other having nuclei NA and NB and electrons present in them are represented by eA and eB. • Attractive forces arise between: • nucleus of one atom and its own electron that is NA – eA and NB– eB. • nucleus of one atom and electron of other atom i.e., NA– eB, NB– eA. • Similarly repulsive forces arise between • electrons of two atoms like eA – eB, • nuclei of two atoms NA – NB.
  • 68.
    • Attractive forcestend to bring the two atoms close to each other whereas repulsive forces tend to push them apart Continue… consider the formation of hydrogen molecule
  • 69.
    • Magnitude ofnew attractive force is more than the new repulsive forces. • As a result, two atoms approach each other and potential energy decreases. • Ultimately a stage is reached where the net force of attraction balances the force of repulsion and system acquires minimum energy. • At this stage two hydrogen atoms are said to be bonded together to form a stable molecule having the bond length of 74 pm. • Since the energy gets released when the bond is formed between two hydrogen atoms, the hydrogen molecule is more stable than that of isolated hydrogen atoms. • The energy so released is called as bond enthalpy Continue… consider the formation of hydrogen molecule
  • 71.
    When H –H distance = 74 pm, Repulsion = Attraction → strongest bond → optimal overlap → lowest energy → bond enthalpy -435.8 kJ 74 pm
  • 72.
    At H –H distance > 74 pm, Repulsion < Attraction → weaker bond → too little overlap → atoms come closer 74 pm > 74 pm
  • 73.
    At H –H distance < 74 pm, Repulsion > Attraction → weaker bond → too much overlap → atoms get further apart 74 pm < 74 pm
  • 74.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Overlapping of Atomic Orbitals • Sigma and pi Bonds By-Vijay Kumar Sethi
  • 75.
    Orbital Overlap Concept •In the formation of hydrogen molecule, there is a minimum energy state when two hydrogen atoms are so near that their atomic orbitals undergo partial interpenetration. • The 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. • In general, greater the overlap the stronger is the bond formed between two atoms. • Therefore, according to orbital overlap concept, the formation of a covalent bond between two atoms results by pairing of electrons present in the valence shell having opposite spins.
  • 76.
    Overlapping of AtomicOrbitals • When orbitals of two atoms come close to form bond, their overlap may be positive, negative or zero depending upon the sign (phase) and direction of orientation of amplitude of orbital wave function in space • Positive and negative sign on boundary surface diagrams show the sign (phase) of orbital wave function and are not related to charge. • Orbitals forming bond should have same sign (phase) and orientation in space. • This is called positive overlap.
  • 77.
    Overlapping of AtomicOrbitals • A negative overlap results when two orbitals with different phase overlap with each other. • A zero overlap will results when orbitals do not overlap at all or do not overlap effectively. • A negative and zero overlap of orbitals does not result in bond formation
  • 78.
    Nature of CovalentBonds • The covalent bond may be classified into two types depending upon the types of overlapping: (i) Sigma(σ) bond, and (ii) pi(π) bond (i) Sigma(σ) bond : • This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. This is called as head on overlap or axial overlap. • In this the overlapping of orbitals takes place to a larger extent. Hence, it is stronger • It may be formed by s-s overlapping, s-p overlapping, p–p overlapping • In this electron density is concentrated in between the nuclei of bonded atoms.
  • 79.
    Continue… (i) Sigma(σ) bond: s-s overlapping : In this case, there is overlap of two half filled s-orbitals along the internuclear axis
  • 80.
    Continue… (i) Sigma(σ) bond: s-p overlapping: This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom.
  • 81.
    Continue… (i) Sigma(σ) bond: p–p overlapping : This type of overlap takes place between half filled p- orbitals of the two approaching atoms.
  • 82.
    (ii) pi(π )bond : • In the formation of π bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. • This bond is formed by sidewise or lateral overlapping of atomic orbitals. • In this bond electron charged clouds is above and below the plane of the participating atoms. • Extent of overlapping is smaller so it is weaker.
  • 83.
  • 84.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • HYBRIDISATION By-Vijay Kumar Sethi
  • 85.
  • 87.
    • The numberof hybrid orbitals is equal to the number of the atomic orbitals that get hybridised. • The hybridised orbitals are always equivalent in energy and shape. • The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. • These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement. the type of hybridisation indicates the geometry of the molecules Salient features of hybridisation
  • 88.
    • The orbitalspresent in the valence shell and penultimate shell of the atom are hybridised. • The orbitals undergoing hybridisation should have almost equal energy. • Promotion of electron is not essential condition prior to hybridisation. • It is not necessary that only half filled orbitals participate in hybridisation. In some cases, even filled orbitals and empty orbitals of valence shell and penultimate shell take part in hybridisation. Important conditions for hybridisation
  • 90.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • HYBRIDISATION-Involving s and p orbitals By-Vijay Kumar Sethi
  • 94.
    • The moleculethus gets distorted and the bond angle is reduced to 107° from 109.5°. • The geometry is pyramidal
  • 95.
    • The bondangle in this case is reduced to 104.5° from 109.5° Repulsion lp-lp > lp-bp > bp-bp • Water molecule has V-shape or angular geometry or bent shape.
  • 96.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Hybridisation of Carbon in Ethane, Ethene, Ethyne By-Vijay Kumar Sethi
  • 97.
    Hybridisation in EthaneC2H6 molecule • Each C is sp3 hybrid • C-C bond , sp3 -sp3 sigma bond . • Each C-H sp3 –s sigma bonds . • In ethane C–C bond length is 154 pm and each C–H bond length is 109 pm.
  • 98.
    Hybridisation in EtheneC2H4 molecule • Each C is sp2 hybrid • C-C bond , sp2 -sp2 sigma bond . • Each C-H sp2 –s sigma bonds .
  • 99.
    Continue… Hybridisation in EtheneC2H4 molecule • The unhybridised orbital (2px or 2py) of one carbon atom overlaps sidewise with the similar orbital of the other carbon atom to form weak π bond, which consists of two equal electron clouds distributed above and below the plane of carbon and hydrogen atoms. C=C bond length 134 pm. C–H bond length 108 pm. H–C–H bond angle is 117.6° H–C–C angle is 121°.
  • 100.
    Hybridisation in EthyneC2H2 molecule • Each C is sp hybrid • C-C bond , sp-sp sigma bond . • Each C-H sp –s sigma bonds . • Each of the two unhybridised p orbitals of both the carbon atoms overlaps sidewise to form two π bonds between the carbon atoms.
  • 101.
    Continue… Hybridisation in EthyneC2H2 molecule • The triple bond between the two carbon atoms is made up of one sigma and two pi bonds C≡C bond length 120 pm. C–H bond length 106 pm.
  • 102.
  • 103.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Hybridisation of Elements involving d Orbitals • Formation of PCl5 and SF6 By-Vijay Kumar Sethi
  • 105.
    Formation of PCl5(sp3 d hybridisation): Five sp3 d hybrid orbitals are directed towards the five corners of a trigonal bipyramidal
  • 106.
    Continue… Formation of PCl5(sp3 d hybridisation): • All the bond angles in trigonal bipyramidal geometry are not equivalent. • Three P–Cl bond lie in one plane and make an angle of 120° with each other; these bonds are termed as equatorial bonds. • The remaining two P–Cl bonds–one lying above and the other lying below the equatorial plane, make an angle of 90° with the plane. These bonds are called axial bonds. As the axial bond pairs suffer more repulsive interaction from the equatorial bond pairs, therefore axial bonds are slightly longer and hence slightly weaker than the equatorial bonds; which makes PCl5 molecule more reactive.
  • 107.
    Formation of SF6(sp3 d2 hybridisation) sp3 d2 hybrid orbitals are projected towards the six corners of a regular octahedron in SF6.
  • 108.
    • These sixsp3 d2 hybrid orbitals overlap with singly occupied orbitals of fluorine atoms to form six S–F sigma bonds. • Thus SF6 molecule has a regular octahedral geometry Continue… Formation of SF6 (sp3 d2 hybridisation)
  • 109.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Molecular Orbital Theory(MOT) • Linear Combination of Atomic Orbitals (LCAO) By-Vijay Kumar Sethi
  • 110.
    MOLECULAR ORBITAL THEORY •Developed by F. Hund and R.S. Mulliken in 1932. • The salient features of this theory are : 1) The electrons in a molecule are present in the various molecular orbitals. 2) The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals. 3) Molecular orbital is polycentric. The electron in a molecular orbital is influenced by two or more nuclei depending upon the number of atoms in the molecule
  • 111.
    Continue… MOLECULAR ORBITAL THEORY 4)The number of molecular orbital formed is equal to the number of combining atomic orbitals. When two atomic orbitals combine, two molecular orbitals are formed. One is known as bonding molecular orbital while the other is called antibonding molecular orbital. 5) The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital. 6) The electron probability distribution around a group of nuclei in a molecule is given by a molecular orbital. 7) The molecular orbitals like atomic orbitals are filled in accordance with the aufbau principle obeying the Pauli’s exclusion principle and the Hund’s rule
  • 112.
    Formation of MolecularOrbitals Linear Combination of Atomic Orbitals (LCAO) • This is an approximate method to obtain wave functions for molecular orbitals. • Consider hydrogen molecule consisting of two atoms of hydrogen ,A and B. • Each hydrogen atom in the ground state has one electron in 1s orbital. • The atomic orbitals of these atoms may be represented by the wave functions ψA and ψB. • Mathematically, the formation of molecular orbitals may be described by the linear combination of atomic orbitals that can take place by addition and by subtraction of wave functions of individual atomic orbitals as shown :
  • 113.
    Continue… Linear Combination ofAtomic Orbitals (LCAO) • The two molecular orbitals σ and σ* are formed as : • The molecular orbital σ formed by the addition of atomic orbitals is called the bonding molecular orbital while • the molecular orbital σ* formed by the subtraction of atomic orbital is called antibonding molecular orbital
  • 114.
    Bonding Molecular OrbitalAnti-bonding Molecular Orbital Formed by constructive interference of the electron waves of the combining atoms. Formed by destructive interference of the electron waves of the combining atoms. The electron density is located between the nuclei of the bonded atoms because of which the repulsion between the nuclei is very less Most of the electron density is located away from the space between the nuclei hence the repulsion between the nuclei is high. Electrons placed in a bonding molecular orbital tend to hold the nuclei together and stabilise a molecule Electrons placed in the antibonding molecular orbital destabilise the molecule. Possesses lower energy than either of the atomic orbitals that have combined to form it. Possesses higher energy than either of the atomic orbitals that have combined to form it. It does not has nodal plane It has is a nodal plane
  • 115.
    Conditions for theCombination of Atomic Orbitals 1. The combining atomic orbitals must have the same or nearly the same energy. 2. The combining atomic orbitals must have the same symmetry about the molecular axis. By convention z-axis is taken as the molecular axis. • For example, 2pz orbital of one atom can combine with 2pz orbital of the other atom but not with the 2px or 2py orbitals because of their different symmetries. 3. The combining atomic orbitals 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.
  • 116.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Types of Molecular Orbitals By-Vijay Kumar Sethi
  • 117.
    Types of MolecularOrbitals • Molecular orbitals of diatomic molecules are designated as σ (sigma), π (pi), δ (delta), etc. • The sigma (σ) molecular orbitals are symmetrical around the bond-axis while pi (π) molecular orbitals are not symmetrical.
  • 118.
    • For example,the linear combination of 1s orbitals centered on two nuclei produces two molecular orbitals which are symmetrical around the bond- axis. • Such molecular orbitals are of the σ type and are designated as σ1s and σ*1s Continue… Types of Molecular Orbitals
  • 119.
    Continue… Types of MolecularOrbitals • If internuclear axis is taken to be in the z-direction, a linear combination of 2pz- orbitals of two atoms also produces two sigma molecular orbitals designated as σ2pz and σ*2pz.
  • 120.
    Continue… Types of MolecularOrbitals • Molecular orbitals obtained from 2px and 2py orbitals are not symmetrical around the bond axis because of the presence of positive lobes above and negative lobes below the molecular plane. • Such molecular orbitals, are labelled as π and π* . (π 2px , π* 2px and π 2py , π* 2py) • A π bonding MO has larger electron density above and below the inter -nuclear axis. The π* antibonding MO has a node between the nuclei.
  • 121.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Energy Level Diagram for Molecular Orbitals By-Vijay Kumar Sethi
  • 122.
    Energy Level Diagramfor Molecular Orbitals • 1s atomic orbitals on two atoms form two molecular orbitals designated as σ1s and σ*1s. • 2s and 2p atomic orbitals (eight atomic orbitals on two atoms) give rise to the following eight molecular orbitals: • The increasing order of energies of various molecular orbitals for O2 and F2 is : • For molecules such as B2, C2, N2, etc. the increasing order of energies is
  • 123.
    for O2 andF2 For B2, C2, N2
  • 124.
    Electronic configuration ofthe molecule • The distribution of electrons among various molecular orbitals is called the electronic configuration of the molecule Stability of Molecules: • If Nb is the number of electrons occupying bonding orbitals and Na the number occupying the antibonding orbitals, then • (i) the molecule is stable if Nb is greater than Na, and • (ii) the molecule is unstable if Nb is less than Na.
  • 125.
    Bond order • Bondorder (b.o.) is defined as one half the difference between the number of electrons present in the bonding and the antibonding orbitals i.e., Bond order (b.o.) = ½ (Nb–Na) • Nb > Na Positive bond order means a stable molecule • Nb< Na Negative bond order , unsatble molecule • Nb = Na Zero bond order , unstable molecule. • Integral bond order values of 1, 2 or 3 correspond to single, double or triple bonds respectively • The bond length decreases as bond order increases.
  • 126.
    Magnetic nature • Ifall the molecular orbitals in a molecule are doubly occupied, the substance is diamagnetic (repelled by magnetic field). No unpaired electron • However if one or more molecular orbitals are singly occupied it is paramagnetic (attracted by magnetic field), e.g., O2molecule. Presence of unpaired electron
  • 127.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • Bonding In Some Homonuclear Diatomic Molecules as per MOT By-Vijay Kumar Sethi
  • 128.
    Hydrogen molecule (H2) electronic configuration of H2:- (σ1s)2 Magnetic property:- Diamagnetic ( no unpaired electron)
  • 129.
    Helium molecule (He2): • electronic configuration of He2:- (σ1s)2 (σ*1s)2 • He2 molecule does not exist. • Similarly, it can be shown that Be2 molecule also does not exist. (σ1s)2 (σ*1s)2 (σ2s)2 (σ*2s)2 Bond order = 0
  • 130.
    Lithium molecule (Li2): • The electronic configuration of lithium is 1s2 2s1 . • There are six electrons in Li2. • The electronic configuration of Li2 molecule is • Li2 : (σ1s)2 (σ*1s)2 (σ2s)2 • The above configuration is also written as KK(σ2s)2 where KK represents the closed K shell structure (σ1s)2 (σ*1s)2 . • Li2 Diamagnetic and stable. • Li2 molecules are known to exist in the vapour phase.
  • 131.
    Carbon molecule (C2): • The electronic configuration of carbon is 1s2 2s2 2p2 . • There are twelve electrons in C2. • The electronic configuration of C2 molecule is C2 diamagnetic, detected in vapour phase. Double bond in C2 consists of both pi bonds because of the presence of four electrons in two pi molecular orbitals.
  • 133.
    Oxygen molecule (O2): • The electronic configuration of oxygen atom is 1s2 2s2 2p4 . each atom 8 electrons • There are 16 electrons in O2molecule.. • The electronic configuration of O2 molecule is
  • 134.
    Oxygen molecule (O2 ): itcontains two unpaired electrons in p * 2px and p * 2py molecular orbitals, therefore it is paramagnetic
  • 135.
    Class XI Chemistry UNIT4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Topics:- • HYDROGEN BONDING By-Vijay Kumar Sethi
  • 136.
    HYDROGEN BONDING • Nitrogen,oxygen and fluorine are the highly electronegative elements. • When they are attached to a hydrogen atom to form covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. • This partially positively charged hydrogen atom of one molecule forms a bond with the more electronegative atom of other molecule. This bond is known as hydrogen bond and is weaker than the covalent bond. • For example HF • Hydrogen bond is represented by a dotted line (– – –) while a solid line represents the covalent bond. • Thus, hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule.
  • 137.
    Cause of Formationof Hydrogen Bond • An electrostatic force of attraction between hydrogen atom of one molecule and electronegative atom of second molecule. Conditions for the Formation of Hydrogen Bond • (1) Hydrogen atom should be bonded to highly electronegative atom say F, O orN. (2) The electronegative atoms should be small in size. • For example in HF and HCl , H atom is bonded to more electronegative atom F and Cl respectively but H-bond exist in HF not in HCl because size of Cl is larger
  • 138.
    Types of H-Bonds Thereare two types of H-bonds (i) Intermolecular hydrogen bond (ii) Intramolecular hydrogen bond (1) Intermolecular hydrogen bond : • It is formed between two different molecules of the same or different compounds. • For example, H-bond in case of HF molecule, alcohol or water molecules, etc.
  • 139.
    (2) Intramolecular hydrogenbond : • It is formed when hydrogen atom is in between the two highly electronegative (F, O, N) atoms present within the same molecule. • For example, in o-nitrophenol the hydrogen is in between the two oxygen atoms.