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CHEMICAL BONDING AND MOLECULAR STRUCTURE, Chemical bond: the attractive force which holds various constituents (atoms, Ions, etc) together in different chemical, species is called a chemical bond., • Octet Rule: Atoms of different elements take part in chemical combination in order to complete their octet or to, attain the noble gas configuration., • Valence Electrons: It is the outermost shell electron which takes part in chemical combination., • Facts Stated by Kossel in Relation to Chemical Bonding:, In the periodic table, the highly electronegative halogens and the highly electro-positive alkali metals are, separated by noble gases., Formation of an anion and cation by the halogens and alkali metals are formed by gain of electron and loss of, electron respectively., Both the negative and positive ions acquire the noble gas configuration., The negative and positive ions are stabilized by electrostatic attraction., Example:, ,  Modes of Chemical Combination, — By the transfer of electrons: The chemical bond which formed by the complete transfer of one or more, electrons from one atom to another is termed as electrovalent bond or ionic bond., — By sharing of electrons: The bond which is formed by the equal sharing of electrons between one or two atoms, is called covalent bond. In these bonds electrons are contributed by both., — Co-ordinate bond: When the electrons are contributed by one atom and shared by both, the bond is formed, and it is known as dative bond or co-ordinate bond., Covalent Bond—Lewis-Langmuir Concept,  When the bond is formed between two or more atoms by mutual contribution and sharing of electrons, it is, known as covalent bond.,  If the combining atoms are same the covalent molecule is known as homoatomic. If they are different, they are, known as heteroatomic molecule.,  For Example:

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Lewis Representation of Simple Molecules (the Lewis Structures):, The Lewis dot Structure can be written through the following steps:, (i) Calculate the total number of valence electrons of the combining atoms., (ii) Each anion means addition of one electron and each cation means removal of one electron. This gives the total, number of electrons to be distributed., (iii) By knowing the chemical symbols of the combining atoms., (iv) After placing shared pairs of electrons for single bond, the remaining electrons may account for either multiple, bonds or as lone pairs. It is to be noted that octet of each atom should be completed., For example: The oxygen atom which has six electrons in its valence shell completes its octet by sharing its two, electrons with two hydrogen atoms to form a water molecule., , Lewis Structure of Water Molecule, Types of Covalent Bonds:, Depending upon the number of shared electron pairs, the covalent bond can be classified into:, , Single Covalent Bond, , Double Covalent Bond, , Triple Covalent Bond, o Single Bonds, A single bond is formed when only one pair of the electron is shared between the two participating atoms. It is, represented by one dash (-). Although this form of covalent bond has a smaller density and is weaker than a double, and triple bond, it is the most stable., For Example, HCL molecule has one Hydrogen atom with one valence electron and one Chlorine atom with seven, valence electrons. In this case, a single bond is formed between hydrogen and chlorine by sharing one electron., , o Double Bonds, A double bond is formed when two pairs of electrons are shared between the two participating atoms. It is, represented by two dashes (=). Double covalent bonds are much stronger than a single bond, but they are less stable., Example: Carbon dioxide molecule has one carbon atom with six valence electrons and two oxygen atom with four, valence electrons., To complete its octet, carbon shares two of its valence electrons with one oxygen atom and two with another oxygen, atom. Each oxygen atom shares its two electrons with carbon and therefore there are two double bonds in CO2., , CO2 Molecule with Double Covalent bond, Oxygen-Molecule: In the formation of the oxygen molecule, each oxygen atom has six electrons in their valence shell., Each atom requires two more electrons to complete their octet. Therefore the atoms share two electrons each to, form the oxygen molecule. Since two electron pairs are shared there is a double bond between the two oxygen, atoms., , O2 Molecule with Double Covalent bond, Ethylene Molecule: In ethylene, each carbon atom shares two of its valence electron with two hydrogen atoms and, remaining two electrons with the other carbon atom. So there is a double bond between the carbon atoms.

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Double Bond in Ethylene Molecule, o Triple Bond, A triple bond is formed when three pairs of electrons are shared between the two participating atoms. Triple covalent, bonds are represented by three dashes (≡) and are the least stable types of covalent bonds., For Example:, In the formation of a nitrogen molecule, each nitrogen atoms having five valence electrons provides three electrons, to form three electron pairs for sharing. Thus, a triple bond is formed between the two nitrogen atoms., , Nitrogen Molecule with Triple Bond, • Formal Charge:, In polyatomic ions, the net charge is the charge on the ion as a whole and not by particular atom. However, charges, can be assigned to individual atoms or ions. These are called formal charges., It can be expressed as, , • Limitations of the Octet Rule:, (i) The incomplete octet of the central atoms: In some covalent compounds central atom has less than eight, electrons, i.e., it has an incomplete octet. For example,, , Li, Be and B have 1, 2, and 3 valence electrons only., (ii) Odd-electron molecules: There are certain molecules which have odd number of electrons the octet rule is not, applied for all the atoms.

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(iii) The expanded Octet: In many compounds there are more than eight valence electrons around the central atom., It is termed as expanded octet. For Example,, , • Other Drawbacks of Octet Theory, (i) Some noble gases, also combine with oxygen and fluorine to form a number of compounds like XeF2 , XeOF2 etc., (ii) This theory does not account for the shape of the molecule., (iii) It does not give any idea about the energy of the molecule and relative stability., Ionic or Electrovalent Bond:,  Ionic or Electrovalent bond is formed by the complete transfer of electrons from one atom to another. Generally, it, is formed between metals and non-metals. The electrostatic force of attraction holds the oppositely charged ions, together.,  The compounds which is formed by ionic or electrovalent bond is known as electrovalent compounds. For, Example,, (i) NaCl is an electrovalent/ Ionic compound. Formation of NaCl is given below:, , Na+ ion has the configuration of Ne while Cl– ion represents the configuration of Ar., (ii) Formation of magnesium oxide from magnesium and oxygen., MgO, Electrovalency: Electrovalency is the number of electrons lost or gained during the formation of an ionic bond or, electrovalent bond., Factors Affecting the Formation of Ionic Bond:, (i) Ionization enthalpy: Ionization enthalpy of any element is the amount of energy required to remove an electron, from outermost shell of an isolated gaseous atom to convert it into cation., Hence, lesser the ionization enthalpy, easier will be the formation of a cation and have greater chance to form an, ionic bond. Due to this reason alkali metals have more tendency to form an ionic bond., For example, in formation of Na+ ion I.E = 496 kJ/mole, While in case of magnesium, it is 743 kJ/mole. That’s why the formation of positive ion for sodium is easier than that, of magnesium., Lower the ionization enthalpy, greater the chances of ionic bond formation., (ii) Electron gain enthalpy (Electron affinities): It is defined as the energy released when an isolated gaseous atom, takes up an electron to form anion. Greater the negative electron gain enthalpy, easier will be the formation of, anion. Consequently, the probability of formation of ionic bond increases., For example. Halogens possess high electron affinity. So, the formation of anion is very common in halogens., , Greater the negative electron gain enthalpy, greater the chances of ionic bond formation., (iii) Lattice energy or enthalpy: It is defined as the amount of energy required to separate 1 mole of ionic compound, into separate oppositely charged ions., Lattice energy of an ionic compound depends upon following factors:

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(i) Size of the ions: Smaller the size, greater will be the lattice energy., (ii) Charge on the ions: Greater the magnitude of charge, greater the interionic attraction and hence higher the, lattice energy., Bond Parameters, Covalent bonds can be characterized on the basis of several bond parameters such as bond length, bond angle, bond, order, and bond energy (also known as bond enthalpy). These bond parameters offer insight into the stability of a, chemical compound and the strength of the chemical bonds holding its atoms together., Bond Length, It is defined as the equilibrium distance between the centres of the nuclei of the two bonded atoms covalent bonds,, the bond length is inversely proportional to the bond order – higher bond orders result in stronger bonds, which are, accompanied by stronger forces of attraction holding the atoms together., Covalent radius: The covalent radius is the radius of an atom in a covalent bond; it is measured in bonded state., Van der waals radius: The Van der waals radius is the radius of the atom including the valence shell in non bonded, state., Bond Angle, Bond angle can be defined as the angle formed between two covalent bonds that originate from the same atom., Bond Enthalpy, It is defined as the amount of energy required to break one mole of bonds of a particular type to separate them into, gaseous atoms. Bond Enthalpy is also known as bond dissociation enthalpy or simple bond enthalpy. Unit of bond, enthalpy = kJ mol-1, Greater the bond enthalpy, stronger is the bond. For e.g., the H—H bond enthalpy in hydrogen is 435.8 kJ mol-1., The magnitude of bond enthalpy is also related to bond multiplicity. Greater the bond multiplicity, more will be the, bond enthalpy. For e.g., bond enthalpy of C —C bond is 347 kJ mol-1 while that of C = C bond is 610 kJ mol-1., In polyatomic molecules, the term mean or average bond enthalpy is used., Bond Order, The bond order is given by the number of bonds between two atoms in a molecule. For example,, Bond order of H2 (H —H) =1, Bond order of 02 (O = O) =2, Bond order of N2 (N = N) =3, Isoelectronic molecules and ions have identical bond orders., For example, F2 and O22- have bond order = 1, N2, CO and NO+ have bond order = 3., With the increase in bond order, bond enthalpy increases and bond length decreases., For example,, , Resonance Structures, There are many molecules whose behaviour cannot be explained by a single-Lew is structure, For example, Lewis, structure of Ozone represented as follows:

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Thus, according to the concept of resonance, whenever a single Lewis structure cannot explain all the properties of, the molecule, the molecule is then supposed to have many structures with similar energy. Positions of nuclei,, bonding and nonbonding pairs of electrons are taken as the canonical structure of the hybrid which describes the, molecule accurately. For 03, the two structures shown above are canonical structures and the III structure represents, the structure of 03 more accurately. This is also called resonance hybrid., Polarity of Bonds: Ionic Character of covalent bond, Polar and Non-Polar Covalent bonds, Non-Polar Covalent bonds: When the atoms joined by covalent bond are the same like; H2, 02, Cl2, the shared pair of, electrons is equally attracted by two atoms and thus the shared electron pair is equidistant to both of them., Electron pairs lie exactly in the centre of the bonding atoms. As a result, no poles are developed and the bond is, called as non-polar covalent bond. The corresponding molecules are known as non-polar molecules., For Example,, , Polar bond: When covalent bonds formed between different atoms of different electronegativity, shared electron, pair between two atoms gets displaced towards highly electronegative atoms., For Example, in HCl molecule, since electronegativity of chlorine is high as compared to hydrogen thus, electron pair, is displaced more towards chlorine atom, and thus chlorine will acquire a partial negative charge (δ–) and hydrogen, atom have a partial positive charge (δ+) with the magnitude of charge same as on chlorination. Such covalent bond is, called polar covalent bond., , Dipole Moment, Due to polarity, polar molecules are also known as dipole molecules and they possess dipole moment. Dipole, moment is defined as the product of magnitude of the positive or negative charge and the distance between the, charges.

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In NH3, the direction of the orbital dipole moment due to lone pair is in the same direction of the “resultant dipole, moment” of the 3 N-H bonds. Hence, the net dipole moment of NH3 4.90 x 10-30cm., But in NF3, molecule the direction of the orbital dipole moment due to lone pair is in the opposite direction to the, “resultant dipole moment” of the 3 N-F bonds. Hence, its net dipole moment decreases to 0.80 x 10-30cm. Therefore,, dipole moment of NH3 is greater than that of NF3., Applications of Dipole Moment,  For determining the polarity of the molecules.,  In finding the shapes of the molecules.,  For example, the molecules with zero dipole moment will be linear or symmetrical. Those molecules which, have unsymmetrical shapes will be either bent or angular., (e.g., NH3 with μ = 1.47 D).,  In calculating the percentage ionic character of polar bonds., Polarization: Covalent character of ionic bonds, When cations and anions approach each other, the valence shell of anions are pulled towards cation nucleus due to, the coulombic attraction and thus shape of the anion is deformed. This phenomenon of deformation of anion by a, cation is known as polarization and the ability of cation to polarize a nearby anion is called as polarizing power of, cation., , The partial covalent character of ionic bonds follows the following rules:,  The smaller the size of cation and larger the size on anion more will be the covalent character., Example: Li+ > Na+,  The greater the charge on the cation, greater the covalent character. Example: Fe+2 < Fe+3,  The cation with transition metals configuration is more polarizing than those with noble gas configuration., The Valence Shell Electron Pair Repulsion (VSEPR) Theory, Sidgwick and Powell in 1940, proposed a simple theory based on repulsive character of electron pairs in the valence, shell of the atoms. It was further developed by Nyholm and Gillespie (1957)., Main Postulates are the following:, (i) The exact shape of molecule depends upon the number of electron pairs (bonded or non bonded) around the, central atoms., (ii) The electron pairs have a tendency to repel each other since they exist around the central atom and the electron, clouds are negatively charged., (iii) Electron pairs try to take such position which can minimize the rupulsion between them., (iv) The valence shell is taken as a sphere with the electron pairs placed at maximum distance.

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(v) A multiple bond is treated as if it is a single electron pair and the electron pairs which constitute the bond as, single pairs., lp= lone pair, bp= bond pair, Valence Bond Theory, Valence bond theory was introduced by Heitler and London (1927) and developed by Pauling and others. It is based, on the concept of atomic orbitals and the electronic configuration of the atoms., Let us consider the formation of hydrogen molecule based on valence-bond theory., Let two hydrogen atoms A and B having their nuclei NA and NB and electrons present in them are eA and eB ., As these two atoms come closer new attractive and repulsive forces begin to operate., (i) The nucleus of one atom is attracted towards its own electron and the electron of the other and vice versa., (ii) Repulsive forces arise between the electrons of two atoms and nuclei of two atoms. Attractive forces tend to, bring the two atoms closer whereas repulsive forces tend to push them apart., The magnitude of new attractive force is more than the new repulsive forces. As a result, the two atoms approach, each other and potential energy decreases. At one stage the net force of attraction balances the force of repulsion, and total energy will be minimum. At this stage the atoms are said to be bonded and form stable molecule with, minimum energy. Since the energy is released thus the molecule has less energy than the individual atoms., The important postulates of the valence bond theory are listed below:,  Covalent bonds are formed when two valence orbitals (half-filled) belonging to two different atoms overlap on, each other. The electron density in the area between the two bonding atoms increases as a result of this, overlapping, thereby increasing the stability of the resulting molecule.,  The presence of many unpaired electrons in the valence shell of an atom enables it to form multiple bonds with, other atoms. The paired electrons present in the valence shell do not take participate in the formation of, chemical bonds as per the valence bond theory.,  Covalent chemical bonds are directional and are also parallel to the region corresponding to the atomic orbitals, that are overlapping.,  Sigma bonds and pi bonds differ in the pattern that the atomic orbitals overlap in, i.e. pi bonds are formed from, sidewise overlapping whereas the overlapping along the axis containing the nuclei of the two atoms leads to the, formation of sigma bonds., Orbital Overlap Concept, According to orbital overlap concept, covalent bond formed between atoms results in the overlap of orbitals, belonging to the atoms having opposite spins of electrons. Stability of a Molecular orbital depends upon the extent, of the overlap of the atomic orbitals. Formation of hydrogen molecule as a result of overlap of the two atomic, orbitals of hydrogen atoms is shown in the figure that follows:, , Types of Orbital Overlap, Depending upon the type of overlapping, the covalent bonds are of two types, known as sigma (σ ) and pi (π) bonds., (i) Sigma (σ bond): Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the, internuclear axis., The axial overlap involving these orbitals is of three types:

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• s-s overlapping: In this case, there is overlap of two half-filled s-orbitals along the internuclear axis as shown, below:, , • s-p overlapping: This type of overlapping occurs between half-filled s-orbitals of one atom and half filled p-orbitals, of another atoms., , • p-p overlapping: This type of overlapping takes place between half filled p-orbitals of the two approaching atoms., , (ii) pi (π bond): π bond is formed by the atomic orbitals when they overlap in such a way that their axis remain, parallel to each other and perpendicular to the internuclear axis. The orbital formed is due to lateral overlapping or, side wise overlapping., , Strength of Sigma and pi Bonds, Sigma bond (σ bond) is formed by the axial overlapping of the atomic orbitals while the π-bond is formed by side, wise overlapping. Since axial overlapping is greater as compared to side wise. Thus, the sigma bond is said to be, stronger bond in comparison to a π-bond., Distinction between sigma and pi bonds, , Hybridisation, Hybridisation is the process of intermixing of the orbitals 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:, (i) Orbitals with almost equal energy take part in the hybridisation., (ii) Number of hybrid orbitals produced is equal to the number of atomic orbitals mixed,

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(iii) Geometry of a covalent molecule can be indicated by the type of hybridisation., (iv) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals., Conditions necessary for hybridisation:, (i) Orbitals of valence shell take part in the hybridisation., (ii) Orbitals involved in hybridisation should have almost equal energy., (iii) Promotion of electron is not necessary condition prior to hybridisation., (iv) In some cases filled orbitals of valence shell also take part in hybridisation., Types of Hybridisation:, (i) sp hybridisation: When one s and one p-orbital hybridise to form two equivalent orbitals, the orbital is known as, sp hybrid orbital, and the type of hybridisation is called sp hybridisation., Each of the hybrid orbitals formed has 50% s-characer and 50%, p-character. This type of hybridisation is also known, as diagonal hybridisation., , (ii) sp2 hybridisation: In this type, one s and two p-orbitals hybridise to form three equivalent sp2 hybridised orbitals., All the three hybrid orbitals remain in the same plane making an angle of 120°. Example, A few compounds in which, sp2 hybridisation takes place are BF3, BH3, BCl3 carbon compounds containing double bond etc.

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(iii) sp3 hybridisation: In this type, one s and three p-orbitals in the valence shell of an atom get hybridised to form, four equivalent hybrid orbitals. There is 25% s-character and 75% p-character in each sp3 hybrid orbital. The four sp3, orbitals are directed towards four corners of the tetrahedron., , The angle between sp3 hybrid orbitals is 109.5°., A compound in which sp3 hybridisation occurs is CH4. The structures of NH2 and H20 molecules can also be explained, with the help of sp3 hybridisation., Hybridization of elements involving d orbitals:, The elements that are present in the third period comprise d orbitals along with s and p orbitals. The energy of the, 3d orbitals is close to the energy of 3s as well as 3p orbitals. The energy of 3d orbitals is also equivalent to 4s as well, as 4p orbitals. As a result, the hybridization including either 3s, 3p and 3d or 3d, 4s, and 4p is feasible. Due to the, difference in energies of 3p and 4s orbitals, no hybridization including 3p, 3d, and 4s orbitals is possible.,  sp3d Hybridisation (PCl5):, sp3d hybridization involves the mixing of 3p orbitals and 1d orbital to form 5 sp3d hybridized orbitals of equal, energy. They have trigonal bipyramidal geometry., Ground state-, , Excited state-, , Three hybrid orbitals lie in the horizontal plane inclined at an angle of 120° to each other known as the equatorial, orbitals. The remaining two orbitals lie in the vertical plane at 90 degrees plane of the equatorial orbitals known as, axial orbitals. Example: Hybridization in Phosphorus pentachloride (PCl5), ,  sp3d2 Hybridization(SF6):, sp3d2 hybridization has 1s, 3p and 2d orbitals, that undergo intermixing to form 6 identical sp3d2 hybrid orbitals., These 6 orbitals are directed towards the corners of an octahedron. The geometry is octahedral., They are inclined at an angle of 90 degrees to one another.

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The Molecular Orbital Theory, The Molecular Orbital Theory (often abbreviated to MOT) is a theory on chemical bonding developed by F. Hund and, R. S. Mulliken to describe the structure and properties of different molecules., The valence-bond theory failed to explain how certain molecules contain two or more equivalent bonds whose bond, orders lie between that of a single bond and that of a double bond, such as the bonds in resonance-stabilized, molecules., The key features of the molecular orbital theory are listed below:,  The total number of molecular orbitals formed will always be equal to the total number of atomic orbitals offered, by the bonding species.,  There exist different types of molecular orbitals viz; bonding molecular orbitals, anti-bonding molecular orbitals,, and non-bonding molecular orbitals. Of these, anti-bonding molecular orbitals will always have higher energy than, the parent orbitals whereas bonding molecular orbitals will always have lower energy than the parent orbitals.,  The electrons are filled into molecular orbitals in the increasing order of orbital energy (from the orbital with the, lowest energy to the orbital with the highest energy).,  The most effective combinations of atomic orbitals (for the formation of molecular orbitals) occur when the, combining atomic orbitals have similar energies., The molecular orbital theory states that each atom tends to combine together and form molecular orbitals. As a, result of such arrangement, electrons are found in various atomic orbitals and they are usually associated with, different nuclei. In short, an electron in a molecule can be present anywhere in the molecule., One of the main impacts of the molecular orbital theory after its formulation is that it paved a new way to, understand the process of bonding. With this theory, the molecular orbitals are basically considered as linear, combinations of atomic orbitals., Linear Combination of Atomic Orbitals (LCAO):, The formation of molecular orbitals can be explained by the linear combination of atomic orbitals. Combination, takes place either by addition or by subtraction of wave function as shown below:

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The molecular orbital formed by addition of atomic orbitals is called bonding molecular orbital while molecular, orbital formed by subtraction of atomic orbitals is called antibonding molecular orbital., Anti Bonding Molecular Orbitals, The electron density is concentrated behind the nuclei of the two bonding atoms in anti-bonding molecular orbitals., This results in the nuclei of the two atoms being pulled away from each other. These kinds of orbitals weaken the, bond between two atoms., Non-Bonding Molecular Orbitals, In the case of non-bonding molecular orbitals, due to a complete lack of symmetry in the compatibility of two, bonding atomic orbitals, the molecular orbitals formed have no positive or negative interactions with each other., These types of orbitals do not affect the bond between the two atoms., Characteristics of Bonding Molecular Orbitals:, • The probability of finding the electron in the internuclear region of the bonding molecular orbital is greater than, that of combining atomic orbitals., • The electrons present in the bonding molecular orbital result in the attraction between the two atoms., • The bonding molecular orbital has lower energy as a result of attraction and hence has greater stability than that, of the combining atomic orbitals., • They are formed by the additive effect of the atomic orbitals so that the amplitude of the new wave is given by, Φ= ΨA + ΨB, • They are represented by σ, π, and δ., Characteristics of Anti-bonding Molecular Orbitals:, • The probability of finding the electron in the internuclear region decreases in the anti-bonding molecular, orbitals., • The electrons present in the anti-bonding molecular orbital result in the repulsion between the two atoms., • The anti-bonding molecular orbitals have higher energy because of the repulsive forces and lower stability., • They are formed by the subtractive effect of the atomic orbitals. The amplitude of the new wave is given by Φ ´=, ΨA – ΨB, • They are represented by σ∗, π∗, δ∗, The energy levels of bonding molecular orbitals are always lower than those of anti-bonding molecular orbitals. This, is because the electrons in the orbital are attracted by the nuclei in the case of bonding Molecular Orbitals whereas, the nuclei repel each other in the case of the anti-bonding Molecular Orbitals., Conditions for the combination of atomic orbitals:, (1) The combining atomic orbitals must have almost equal energy:

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The atomic orbitals combining to form molecular orbitals should have comparable energy. This means that 2p, orbital of an atom can combine with another 2p orbital of another atom but 1s and 2p cannot combine together as, they have appreciable energy difference., (2) The combining atomic orbitals must have same symmetry about the molecular axis:, The combining atoms should have the same symmetry around the molecular axis for proper combination, otherwise,, the electron density will be sparse. For e.g. all the sub-orbitals of 2p have the same energy but still, 2pz orbital of an, atom can only combine with a 2pz orbital of another atom but cannot combine with 2px and 2py orbital as they have, a different axis of symmetry. In general, the z-axis is considered as the molecular axis of symmetry., (3) The combining atomic orbitals must overlap to the maximum extent., The two atomic orbitals will combine to form molecular orbital if the overlap is proper. Greater the extent of overlap, of orbitals, greater will be the nuclear density between the nuclei of the two atoms., Types of Molecular Orbitals:, Sigma (σ) Molecular Orbitals: They are symmetrical around the bond-axis., pi (π) Molecular Orbitals: They are not symmetrical, because of the presence of positive lobes above and negative, lobes below the molecular plane.,  Electronic configuration and Molecular Behaviour: The distribution of electrons among various molecular orbitals, is called electronic configuration of the molecule., • Stability of Molecules, , • Bond Order, Bond order is defined as half of the difference between the number of electrons present in bonding and antibonding, molecular orbitals., Bond order (B.O.) = 1/2 [Nb-Na],  The bond order may be a whole number, a fraction or even zero.,  It may also be positive or negative.,  Nature of the bond: Integral bond order value for single double and triple bond will be 1, 2 and 3 respectively.,  Bond-Length: Bond order is inversely proportional to bond-length. Thus, greater the bond order, smaller will be, the bond-length.,  Magnetic Nature: If all the molecular orbitals have paired electrons, the substance is diamagnetic. If one or more, molecular orbitals have unpaired electrons, it is paramagnetic e.g., 02 molecule., Bonding in Some Homonuclear (Diatomic) Molecules, Hydrogen molecule (H2): It is formed by the combination of two hydrogen atoms. Each hydrogen atom has one, electron in Is orbital, so, the electronic configuration of hydrogen molecule is, , This indicates that two hydrogen atoms are bonded by a single covalent bond. Bond dissociation energy of hydrogen, has been found = 438 kJ/mole. Bond-Length = 74 pm, No unpaired electron is present therefore,, it is diamagnetic.

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Helium molecule (He2): Each helium atom contains 2 electrons, thus in He2 molecule there would be 4 electrons., The electrons will be accommodated in σ1s and σ*1s molecular orbitals:

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Hydrogen Bonding, When highly electronegative elements like nitrogen, oxygen, flourine are attached to hydrogen to form covalent, bond, the electrons of the covalent bond are shifted towards the more electronegative atom. Thus, partial positive, charge develops on hydrogen atom which forms a bond with the other electronegative atom. This bond is known as, hydrogen bond and it is weaker than the covalent bond. For example, in HF molecule, hydrogen bond exists between, hydrogen atom of one molecule and fluorine atom of another molecule., It can be depicted as, • Types of H-Bonds:, (i) Intermolecular hydrogen bond: It is formed between two different molecules of the same or different, compounds. For Example, in HF molecules, water molecules etc., , (ii) Intramolecular hydrogen bond: In this type, hydrogen atom is in between the two highly electronegative F, N, O, atoms present within the same molecule. For example, in o-nitrophenol, the hydrogen is in between the two oxygen, atoms.