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Unit 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Scientists are constantly discovering new c...

Unit 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE Scientists are constantly discovering new compounds, orderly arranging the facts about them, trying to explain with the existing knowledge, organising to modify the earlier views or evolve theories for explaining the newly After studying this Unit, you will be observed facts. able to understand Kössel-Lewis approach to chemical bonding; explain the octet rule and its Matter is made up of one or different type of elements. limitations, draw Lewis structures Under normal conditions no other element exists as an of simple molecules; independent atom in nature, except noble gases. However, explain the formation of different a group of atoms is found to exist together as one species types of bonds; having characteristic properties. Such a group of atoms is called a molecule. Obviously there must be some force describe the VSEPR theory and which holds these constituent atoms together in the predict the geometry of simple molecules; molecules. The attractive force which holds various constituents (atoms, ions, etc.) together in different explain the valence bond chemical species is called a chemical bond. Since the approach for the formation of formation of chemical compounds takes place as a result of covalent bonds; combination of atoms of various elements in different ways, predict the directional properties it raises many questions. Why do atoms combine? Why are of covalent bonds; only certain combinations possible? Why do some atoms explain the different types of combine while certain others do not? Why do molecules hybridisation involving s, p and possess definite shapes? To answer such questions different d orbitals and draw shapes of theories and concepts have been put forward from time simple covalent molecules; to time. These are Kössel-Lewis approach, Valence Shell Electron Pair Repulsion (VSEPR) Theory, Valence Bond (VB) describe the molecular orbital theory of homonuclear diatomic Theory and Molecular Orbital (MO) Theory. The evolution molecules; of various theories of valence and the interpretation of the nature of chemical bonds have closely been related to explain the concept of hydrogen the developments in the understanding of the structure bond. of atom, the electronic configuration of elements and the periodic table. Every system tends to be more stable and bonding is nature’s way of lowering the energy of the system to attain stability. 2024-25 Unit 4.indd 100 9/12/2022 9:36:09 AM  Chemical Bonding And Molecular Structure 101 4.1 KÖ ssel - L e w i s A p p roach t o the number of valence electrons. This number Chemical Bonding of valence electrons helps to calculate the In order to explain the formation of chemical common or group valence of the element. bond in terms of electrons, a number of The group valence of the elements is generally attempts were made, but it was only in either equal to the number of dots in Lewis 1916 when Kössel and Lewis succeeded symbols or 8 minus the number of dots or independently in giving a satisfactory valence electrons. explanation. They were the first to provide Kössel, in relation to chemical bonding, some logical explanation of valence which was drew attention to the following facts: based on the inertness of noble gases. In the periodic table, the highly Lewis pictured the atom in terms of a electronegative halogens and the highly positively charged ‘Kernel’ (the nucleus plus electropositive alkali metals are separated the inner electrons) and the outer shell that by the noble gases; could accommodate a maximum of eight The formation of a negative ion from a electrons. He, further assumed that these halogen atom and a positive ion from eight electrons occupy the corners of a cube an alkali metal atom is associated with which surround the ‘Kernel’. Thus the single the gain and loss of an electron by the outer shell electron of sodium would occupy respective atoms; one corner of the cube, while in the case of The negative and positive ions thus a noble gas all the eight corners would be formed attain stable noble gas electronic occupied. This octet of electrons, represents configurations. The noble gases (with the a particularly stable electronic arrangement. exception of helium which has a duplet Lewis postulated that atoms achieve of electrons) have a particularly stable the stable octet when they are linked by outer shell configuration of eight (octet) chemical bonds. In the case of sodium and electrons, ns2np6. chlorine, this can happen by the transfer of The negative and positive ions are stabilized an electron from sodium to chlorine thereby by electrostatic attraction. giving the Na+ and Cl– ions. In the case of other molecules like Cl2, H2, F2, etc., the bond For example, the formation of NaCl from is formed by the sharing of a pair of electrons sodium and chlorine, according to the above scheme, can be explained as: between the atoms. In the process each atom attains a stable outer octet of electrons. Na → Na+ + e– Lewis Symbols: In the formation of a [Ne] 3s1 [Ne] molecule, only the outer shell electrons take Cl + e– → Cl– part in chemical combination and they are [Ne] 3s 3p [Ne] 3s2 3p6 or [Ar] 2 5 known as valence electrons. The inner shell Na+ + Cl– → NaCl or Na+Cl– electrons are well protected and are generally not involved in the combination process. Similarly the formation of CaF2 may be G.N. Lewis, an American chemist introduced shown as: simple notations to represent valence electrons Ca → Ca2+ + 2e– in an atom. These notations are called Lewis [Ar]4s2 [Ar] symbols. For example, the Lewis symbols for F + e– → F– the elements of second period are as under: [He] 2s2 2p5 [He] 2s2 2p6 or [Ne] Ca2+ + 2F– → CaF2 or Ca2+(F– )2 The bond formed, as a result of the Significance of Lewis Symbols : The electrostatic attraction between the number of dots around the symbol represents positive and negative ions was termed as 2024-25 Unit 4.indd 101 9/12/2022 9:36:09 AM  102 chemistry the electrovalent bond. The electrovalence chlorine atoms attain the outer shell octet of is thus equal to the number of unit charge(s) the nearest noble gas (i.e., argon). on the ion. Thus, calcium is assigned a The dots represent electrons. Such positive electrovalence of two, while chlorine structures are referred to as Lewis dot a negative electrovalence of one. structures. Kössel’s postulations provide the basis for The Lewis dot structures can be written for the modern concepts regarding ion-formation other molecules also, in which the combining by electron transfer and the formation of ionic atoms may be identical or different. The crystalline compounds. His views have proved important conditions being that: to be of great value in the understanding and Each bond is formed as a result of sharing systematisation of the ionic compounds. At of an electron pair between the atoms. the same time he did recognise the fact that a large number of compounds did not fit into Each combining atom contributes at least these concepts. one electron to the shared pair. The combining atoms attain the outer- 4.1.1 Octet Rule shell noble gas configurations as a result Kössel and Lewis in 1916 developed an of the sharing of electrons. important theory of chemical combination Thus in water and carbon tetrachloride between atoms known as electronic theory molecules, formation of covalent bonds of chemical bonding. According to this, can be represented as: 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. 4.1.2 Covalent Bond Langmuir (1919) refined the Lewis postulations by abandoning the idea of the stationary cubical arrangement of the octet, and by introducing the term covalent bond. The Lewis-Langmuir theory can be Thus, when two atoms share one understood by considering the formation of electron pair they are said to be joined by the chlorine molecule, Cl2. The Cl atom with a single covalent bond. In many compounds electronic configuration, [Ne]3s2 3p5, is one we have multiple bonds between atoms. The electron short of the argon configuration. formation of multiple bonds envisages sharing The formation of the Cl­2 molecule can be of more than one electron pair between two understood in terms of the sharing of a pair atoms. If two atoms share two pairs of of electrons between the two chlorine atoms, electrons, the covalent bond between them each chlorine atom contributing one electron is called a double bond. For example, in the to the shared pair. In the process both carbon dioxide molecule, we have two double bonds between the carbon and oxygen atoms. Similarly in ethene molecule the two carbon atoms are joined by a double bond. or Cl – Cl Covalent bond between two Cl atoms Double bonds in CO2 molecule 2024-25 Unit 4.indd 102 9/12/2022 9:36:09 AM  Chemical Bonding And Molecular Structure 103 number of valence electrons. For example, for the CO32– ion, the two negative charges indicate that there are two additional electrons than those provided by the neutral atoms. For NH +4 ion, one positive charge indicates the loss of one electron from the group of neutral atoms. C2H4 molecule Knowing the chemical symbols of the When combining atoms share three combining atoms and having knowledge electron pairs as in the case of two nitrogen of the skeletal structure of the compound atoms in the N2 molecule and the two (known or guessed intelligently), it is easy carbon atoms in the ethyne molecule, a to distribute the total number of electrons triple bond is formed. as bonding shared pairs between the atoms in proportion to the total bonds. In general the least electronegative atom occupies the central position in the molecule/ion. For example in the NF3 and N2 molecule CO32–, nitrogen and carbon are the central atoms whereas fluorine and oxygen occupy the terminal positions. After accounting for the shared pairs of electrons for single bonds, the remaining C2H2 molecule electron pairs are either utilized for multiple bonding or remain as the lone 4.1.3 Lewis Representation of Simple pairs. The basic requirement being Molecules (the Lewis Structures) that each bonded atom gets an octet of electrons. The Lewis dot structures provide a picture of bonding in molecules and ions in terms of Lewis representations of a few molecules/ the shared pairs of electrons and the octet ions are given in Table 4.1. rule. While such a picture may not explain the bonding and behaviour of a molecule Table 4.1 The Lewis Representation of completely, it does help in understanding the Some Molecules formation and properties of a molecule to a large extent. Writing of Lewis dot structures of molecules is, therefore, very useful. The Lewis dot structures can be written by adopting the following steps: The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms. For example, in the CH4 molecule there are eight valence electrons available for bonding (4 from carbon and 4 from the four hydrogen atoms). For anions, each negative charge would mean addition of one electron. For cations, each positive charge would result in * Each H atom attains the configuration of helium subtraction of one electron from the total (a duplet of electrons) 2024-25 Unit 4.indd 103 9/12/2022 9:36:10 AM  104 chemistry Problem 4.1 each of the oxygen atoms completing the octets on oxygen atoms. This, however, Write the Lewis dot structure of CO does not complete the octet on nitrogen molecule. if the remaining two electrons constitute Solution lone pair on it. Step 1. Count the total number of valence electrons of carbon and oxygen atoms. The outer (valence) shell configurations of carbon and oxygen atoms are: 2s2 2p2 Hence we have to resort to multiple and 2s2 2p4, respectively. The valence bonding between nitrogen and one of electrons available are 4 + 6 =10. the oxygen atoms (in this case a double bond). This leads to the following Lewis Step 2. The skeletal structure of CO is dot structures. 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 (in this case a triple 4.1.4 Formal Charge bond) between C and O atoms. This satisfies the octet rule condition for both Lewis dot structures, in general, do not atoms. represent the actual shapes of the molecules. 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 Problem 4.2 difference between the number of valence Write the Lewis structure of the nitrite electrons of that atom in an isolated or free ion, NO2–. state and the number of electrons assigned to that atom in the Lewis structure. It is Solution expressed as : Step 1. Count the total number of valence electrons of the nitrogen atom, Formal charge (F.C.) the oxygen atoms and the additional one on an atom in a Lewis = negative charge (equal to one electron). structure N(2s2 2p3), O (2s2 2p4) 5 + (2 × 6) +1 = 18 electrons total number of valence total number of non electrons in the free — bonding (lone pair) Step 2. The skeletal structure of NO2– is atom electrons written as : O N O total number of Step 3. Draw a single bond (one shared — (1/2) bonding (shared) electrons electron pair) between the nitrogen and 2024-25 Unit 4.indd 104 9/12/2022 9:36:10 AM  Chemical Bonding And Molecular Structure 105 The counting is based on the assumption 4.1.5 Limitations of the Octet Rule that the atom in the molecule owns one The octet rule, though useful, is not universal. electron of each shared pair and both the It is quite useful for understanding the electrons of a lone pair. structures of most of the organic compounds Let us consider the ozone molecule (O3). and it applies mainly to the second period The Lewis structure of O3 may be drawn as: elements of the periodic table. There are three types of exceptions to the octet rule. The incomplete octet of the central atom In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. The atoms have been numbered as 1, 2 Examples are LiCl, BeH2 and BCl3. and 3. The formal charge on: The central O atom marked 1 1 Li, Be and B have 1, 2 and 3 valence electrons =6–2– (6) = +1 only. Some other such compounds are AlCl3 2 and BF3. The end O atom marked 2 Odd-electron molecules 1 =6–4– (4) = 0 In molecules with an odd number of electrons 2 like nitric oxide, NO and nitrogen dioxide, The end O atom marked 3 NO2, the octet rule is not satisfied for all the 1 atoms =6–6– (2) = –1 2 Hence, we represent O3 along with the formal charges as follows: 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 We must understand that formal charges around the central atom. This is termed as do not indicate real charge separation within the expanded octet. Obviously the octet rule the molecule. Indicating the charges on the does not apply in such cases. atoms in the Lewis structure only helps in Some of the examples of such compounds keeping track of the valence electrons in are: PF 5 , SF 6 , H 2 SO 4 and a number of the molecule. Formal charges help in the coordination compounds. 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. The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighbouring atoms. 2024-25 Unit 4.indd 105 9/12/2022 9:36:11 AM  106 chemistry Interestingly, sulphur also forms many affinity, is the negative of the energy change compounds in which the octet rule is obeyed. accompanying electron gain. In sulphur dichloride, the S atom has an octet Obviously ionic bonds will be formed of electrons around it. more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy. Other drawbacks of the octet theory Most ionic compounds have cations It is clear that octet rule is based upon derived from metallic elements and anions the chemical inertness of noble gases. from non-metallic elements. The ammonium However, some noble gases (for example ion, NH 4+ (made up of two non-metallic xenon and krypton) also combine with elements) is an exception. It forms the cation oxygen and fluorine to form a number of of a number of ionic compounds. compounds like XeF2, KrF2, XeOF2 etc. Ionic compounds in the crystalline This theory does not account for the shape state consist of orderly three-dimensional of molecules. arrangements of cations and anions held It does not explain the relative stability of together by coulombic interaction energies. the molecules being totally silent about These compounds crystallise in different the energy of a molecule. crystal structures determined by the size of the ions, their packing arrangements and 4.2 Ionic or Electrovalent Bond other factors. The crystal structure of sodium From the Kössel and Lewis treatment of the chloride, NaCl (rock salt), for example is formation of an ionic bond, it follows that the shown below. formation of ionic compounds would primarily 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. The formation of a positive ion involves ionization, i.e., removal of electron(s) from the neutral atom and that of the negative ion involves the addition of electron(s) to the Rock salt structure neutral atom. In ionic solids, the sum of the electron gain M(g) → M+(g) + e– ; enthalpy and the ionization enthalpy may be Ionization enthalpy positive but still the crystal structure gets X(g) + e– → X – (g) ; stabilized due to the energy released in the Electron gain enthalpy formation of the crystal lattice. For example: the ionization enthalpy for Na+(g) formation M+(g) + X –(g) → MX(s) from Na(g) is 495.8 kJ mol–1 ; while the electron The electron gain enthalpy, ∆eg H, is the gain enthalpy for the change Cl(g) + e–→ enthalpy change (Unit 3), when a gas phase Cl– (g) is, – 348.7 kJ mol–1 only. The sum of the atom in its ground state gains an electron. two, 147.1 kJ mol-1 is more than compensated The electron gain process may be exothermic for by the enthalpy of lattice formation of or endothermic. The ionization, on the other NaCl(s) (–788 kJ mol–1). Therefore, the energy hand, is always endothermic. Electron released in the processes is more than the 2024-25 Unit 4.indd 106 9/12/2022 9:36:11 AM  Chemical Bonding And Molecular Structure 107 energy absorbed. Thus a qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving octet of electrons around the ionic species in gaseous state. Since lattice enthalpy plays a key role in the formation of ionic compounds, it is important that we learn more about it. 4.2.1 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 Fig. 4.1 The bond length in a covalent kJ of energy is required to separate one mole molecule AB. of solid NaCl into one mole of Na+ (g) and one R = rA + rB (R is the bond length and rA and rB are mole of Cl– (g) to an infinite distance. the covalent radii of atoms A and B respectively) This process involves both the attractive forces between ions of opposite charges in the same molecule. The van der Waals and the repulsive forces between ions of radius represents the overall size of the like charge. The solid crystal being three- atom which includes its valence shell in a dimensional; it is not possible to calculate nonbonded situation. Further, the van der lattice enthalpy directly from the interaction Waals radius is half of the distance between of forces of attraction and repulsion only. two similar atoms in separate molecules in Factors associated with the crystal geometry a solid. Covalent and van der Waals radii of have to be included. chlorine are depicted in Fig. 4.2. 4.3 Bond Parameters rc = 99 pm 19 4.3.1 Bond Length 8 pm 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 about which you will learn in higher classes. Each atom of the bonded pair contributes to the r vd w = bond length (Fig. 4.1). In the case of a covalent 18 0 bond, the contribution from each atom is pm called the covalent radius of that atom. pm 0 36 The covalent radius is measured approximately as the radius of an atom’s Fig. 4.2 Covalent and van der Waals radii in core which is in contact with the core of an a chlorine molecule. The inner circles adjacent atom in a bonded situation. The correspond to the size of the chlorine covalent radius is half of the distance between atom (rvdw and rc are van der Waals and two similar atoms joined by a covalent bond covalent radii respectively). 2024-25 Unit 4.indd 107 9/12/2022 9:36:11 AM  108 chemistry Some typical average bond lengths for Table 4.2 Average Bond Lengths for Some single, double and triple bonds are shown in Single, Double and Triple Bonds Table 4.2. Bond lengths for some common Covalent Bond molecules are given in Table 4.3. Bond Type Length (pm) The covalent radii of some common O–H 96 elements are listed in Table 4.4. C–H 107 4.3.2 Bond Angle N–O 136 C–O 143 It is defined as the angle between the orbitals C–N 143 containing bonding electron pairs around the C–C 154 central atom in a molecule/complex ion. Bond C=O 121 angle is expressed in degree which can be N=O 122 experimentally determined by spectroscopic C=C 133 methods. It gives some idea regarding the C=N 138 distribution of orbitals around the central C≡N 116 atom in a molecule/complex ion and hence it C≡C 120 helps us in determining its shape. For Table 4.3 Bond Lengths in Some Common example H–O–H bond angle in water can be Molecules represented as under : Molecule Bond Length (pm) H2 (H – H) 74 F2 (F – F) 144 4.3.3 Bond Enthalpy Cl2 (Cl – Cl) 199 It is defined as the amount of energy required Br2 (Br – Br) 228 to break one mole of bonds of a particular I2 (I – I) 267 type between two atoms in a gaseous state. N2 (N ≡ N) 109 The unit of bond enthalpy is kJ mol–1. For O2 (O = O) 121 example, the H – H bond enthalpy in hydrogen HF (H – F) 92 molecule is 435.8 kJ mol–1. HCl (H – Cl) 127 HBr (H – Br) 141 H2(g) → H(g) + H(g); ∆aH = 435.8 kJ mol–1 HI (H – I) 160 Similarly the bond enthalpy for molecules containing multiple bonds, for example O2 and Table 4.4 Covalent Radii, *rcov/(pm) N2 will be as under : O2 (O = O) (g) → O(g) + O(g); ∆aH = 498 kJ mol–1 N2 (N ≡ N) (g) → N(g) + N(g); ∆aH = 946.0 kJ mol–1 It is important that larger the bond dissociation enthalpy, stronger will be the bond in the molecule. For a heteronuclear diatomic molecules like HCl, we have HCl (g) → H(g) + Cl (g); ∆aH = 431.0 kJ mol–1 In case of polyatomic molecules, the measurement of bond strength is more complicated. For example in case  o f H 2O * The values cited are for single bonds, except where molecule, the enthalpy needed to break the otherwise indicated in parenthesis. (See also Unit 3 for two O – H bonds is not the same. periodic trends). 2024-25 Unit 4.indd 108 9/12/2022 9:36:11 AM  Chemical Bonding And Molecular Structure 109 H2O(g) → H(g) + OH(g); ∆aH1 = 502 kJ mol–1 OH(g) → H(g) + O(g); ∆aH2 = 427 kJ mol–1 The difference in the ∆aH value shows that the second O – H bond undergoes some change because of changed chemical environment. This is the reason for some difference in energy of the same O – H bond in different molecules like C2H5OH (ethanol) and water. 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 as explained below in case of water molecule, 502 + 427 Fig. 4.3 Resonance in the O3 molecule Average bond enthalpy = 2 (structures I and II represent the two canonical = 464.5 kJ mol–1 forms while the structure III is the resonance hybrid) 4.3.4 Bond Order In the Lewis description of covalent bond, In both structures we have a O–O single the Bond Order is given by the number bond and a O=O double bond. The normal of bonds between the two atoms in a O–O and O=O bond lengths are 148 pm molecule. The bond order, for example in and 121 pm respectively. Experimentally H2 (with a single shared electron pair), in O2 determined oxygen-oxygen bond lengths in (with two shared electron pairs) and in N2 the O3 molecule are same (128 pm). Thus the (with three shared electron pairs) is 1,2,3 oxygen-oxygen bonds in the O3 molecule are respectively. Similarly in CO (three shared intermediate between a double and a single electron pairs between C and O) the bond bond. Obviously, this cannot be represented order is 3. For N2, bond order is 3 and its by either of the two Lewis structures shown is 946 kJ mol–1; being one of the highest above. for a diatomic molecule. The concept of resonance was introduced Isoelectronic molecules and ions have to deal with the type of difficulty experienced identical bond orders; for example, F2 and in the depiction of accurate structures of O22– have bond order 1. N2, CO and NO+ have molecules like O3. According to the concept bond order 3. of resonance, whenever a single Lewis A general correlation useful for structure cannot describe a molecule understanding the stablities of molecules accurately, a number of structures with is that: with increase in bond order, similar energy, positions of nuclei, bonding bond enthalpy increases and bond length and non-bonding pairs of electrons are decreases. taken as the canonical structures of the hybrid which describes the molecule 4.3.5 Resonance Structures accurately. Thus for O3, the two structures It is often observed that a single Lewis structure shown above constitute the canonical is inadequate for the representation of a structures or resonance structures and molecule in conformity with its experimentally their hybrid i.e., the III structure represents determined parameters. For example, the the structure of O3 more accurately. This is ozone, O3 molecule can be equally represented also called resonance hybrid. Resonance is by the structures I and II shown below: represented by a double headed arrow. 2024-25 Unit 4.indd 109 9/12/2022 9:36:11 AM  110 chemistry Some of the other examples of resonance structures are provided by the carbonate ion and the carbon dioxide molecule. Problem 4.3 Fig. 4.5 Resonance in CO 2 molecule, I, II and III represent the three canonical Explain the structure of CO32– ion in forms. terms of resonance. Solution In general, it may be stated that The single Lewis structure based on Resonance stabilizes the molecule as the the presence of two single bonds and energy of the resonance hybrid is less one double bond between carbon than the energy of any single cannonical and oxygen atoms is inadequate to structure; and, represent the molecule accurately as it represents unequal bonds. According Resonance averages the bond characteristics as a whole. to the experimental findings, all carbon to oxygen bonds in CO32– are equivalent. Thus the energy of the O 3 resonance Therefore the carbonate ion is best hybrid is lower than either of the two described as a resonance hybrid of the cannonical froms I and II (Fig. 4.3). canonical forms I, II, and III shown below. Many misconceptions are associated with resonance and the same need to be dispelled. You should remember that : The cannonical forms have no real existence. The molecule does not exist for a certain fraction of time in one cannonical form and for other fractions of time in other Fig. 4.4 Resonance in CO32–, I, II and III cannonical forms. represent the three canonical There is no such equilibrium between forms. the cannonical forms as we have Problem 4.4 between tautomeric forms (keto and Explain the structure of CO2 molecule. enol) in tautomerism. The molecule as such has a single Solution structure which is the resonance The experimentally determined carbon hybrid of the cannonical forms and to oxygen bond length in CO 2 is which cannot as such be depicted by 115 pm. The lengths of a normal a single Lewis structure. carbon to oxygen double bond (C=O) and carbon to oxygen triple bond (C≡O) 4.3.6 Polarity of Bonds are 121 pm and 110 pm respectively. The carbon-oxygen bond lengths in The existence of a hundred percent ionic or CO2 (115 pm) lie between the values covalent bond represents an ideal situation. for C=O and C≡O. Obviously, a single In reality no bond or a compound is either Lewis structure cannot depict this completely covalent or ionic. Even in case of position and it becomes necessary to covalent bond between two hydrogen atoms, write more than one Lewis structures there is some ionic character. and to consider that the structure of CO2 is best described as a hybrid of When covalent bond is formed between the canonical or resonance forms I, II two similar atoms, for example in H2, O2, and III. Cl2, N2 or F2, the shared pair of electrons is equally attracted by the two atoms. As a result 2024-25 Unit 4.indd 110 9/12/2022 9:36:12 AM  Chemical Bonding And Molecular Structure 111 electron pair is situated exactly between the In case of polyatomic molecules the dipole two identical nuclei. The bond so formed is moment not only depend upon the individual called nonpolar covalent bond. Contrary to dipole moments of bonds known as bond this in case of a heteronuclear molecule like dipoles but also on the spatial arrangement HF, the shared electron pair between the two of various bonds in the molecule. In such atoms gets displaced more towards fluorine case, the dipole moment of a molecule is the since the electronegativity of fluorine (Unit 3) vector sum of the dipole moments of various is far greater than that of hydrogen. The bonds. For example in H2O molecule, which resultant covalent bond is a polar covalent has a bent structure, the two O–H bonds are bond. oriented at an angle of 104.50. Net dipole As a result of polarisation, the molecule moment of 6.17 × 10–30 C m (1D = 3.33564 possesses the dipole moment (depicted × 10–30 C m) is the resultant of the dipole below) which can be defined as the product of moments of two O–H bonds. the magnitude of the charge and the distance between the centres of positive and negative charge. It is usually designated by a Greek letter ‘µ’. Mathematically, it is expressed as follows : Dipole moment (µ) = charge (Q) × distance of separation (r) Dipole moment is usually expressed in Net Dipole moment, µ = 1.85 D Debye units (D). The conversion factor is = 1.85 × 3.33564 × 10–30 C m = 6.17 ×10–30 C m 1 D = 3.33564 × 10–30 C m The dipole moment in case of BeF2 is zero. where C is coulomb and m is meter. This is because the two equal bond dipoles Further dipole moment is a vector quantity point in opposite directions and cancel the and by convention it is depicted by a small effect of each other. 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 In tetra-atomic molecule, for example in cross is on positive end and arrow head is on BF3, the dipole moment is zero although the negative end. For example the dipole moment B – F bonds are oriented at an angle of 120o of HF may be represented as : to one another, the three bond moments give a net sum of zero as the resultant of any two H F is equal and opposite to the third. This arrow symbolises the direction of the shift of electron density in the molecule. Note that the direction of crossed arrow is opposite to the conventional direction of dipole moment vector. Peter Debye, the Dutch chemist received Nobel prize in 1936 for Let us study an interesting case of NH3 his work on X-ray diffraction and dipole moments. The magnitude and NF3 molecule. Both the molecules have of the dipole moment is given in pyramidal shape with a lone pair of electrons Debye units in order to honour him. on nitrogen atom. Although fluorine is more electronegative than nitrogen, the resultant 2024-25 Unit 4.indd 111 9/12/2022 9:36:12 AM  112 chemistry dipole moment of NH3 (4.90 × 10–30 C m) is The smaller the size of the cation and the greater than that of NF3 (0.8 × 10–30 C m). This larger the size of the anion, the greater is because, in case of NH3 the orbital dipole the covalent character of an ionic bond. due to lone pair is in the same direction as The greater the charge on the cation, the the resultant dipole moment of the N – H greater the covalent character of the ionic bonds, whereas in NF3 the orbital dipole is in bond. the direction opposite to the resultant dipole For cations of the same size and charge, moment of the three N–F bonds. The orbital the one, with electronic configuration dipole because of lone pair decreases the effect (n-1)dnnso, typical of transition metals, is of the resultant N – F bond moments, which more polarising than the one with a noble results in the low dipole moment of NF3 as gas configuration, ns2 np6, typical of alkali represented below : and alkaline earth metal cations. The cation polarises the anion, pulling the electronic charge toward itself and thereby increasing the electronic charge between the two. This is precisely what happens in a covalent bond, i.e., buildup of electron charge density between the nuclei. The polarising power of the cation, the polarisability of the anion and the extent of distortion (polarisation) of anion are the factors, which determine the per Dipole moments of some molecules are cent covalent character of the ionic bond. shown in Table 4.5. Just as all the covalent bonds have 4.4 The Valence Shell Electron some partial ionic character, the ionic Pair Repulsion (VSEPR) Theory bonds also have partial covalent character. As already explained, Lewis concept is unable The partial covalent character of ionic to explain the shapes of molecules. This bonds was discussed by Fajans in terms of theory provides a simple procedure to predict the following rules: the shapes of covalent molecules. Sidgwick Table 4.5 Dipole Moments of Selected Molecules Dipole Type of Molecule Example Geometry Moment, µ(D) Molecule (AB) HF 1.78 linear HCl 1.07 linear HBr 0.79 linear Hl 0.38 linear H2 0 linear Molecule (AB2) H2O 1.85 bent H2S 0.95 bent CO2 0 linear Molecule (AB3) NH3 1.47 trigonal-pyramidal NF3 0.23 trigonal-pyramidal BF3 0 trigonal-planar Molecule (AB4) CH4 0 tetrahedral CHCl3 1.04 tetrahedral CCl4 0 tetrahedral 2024-25 Unit 4.indd 112 9/12/2022 9:36:12 AM  Chemical Bonding And Molecular Structure 113 and Powell in 1940, proposed a simple theory result in deviations from idealised shapes and based on the repulsive interactions of the alterations in bond angles in molecules. electron pairs in the valence shell of the For the prediction of geometrical shapes atoms. It was further developed and redefined of molecules with the help of VSEPR theory, by Nyholm and Gillespie (1957). it is convenient to divide molecules into The main postulates of VSEPR theory are two categories as (i) molecules in which as follows: the central atom has no lone pair and The shape of a molecule depends upon (ii) molecules in which the central atom the number of valence shell electron pairs has one or more lone pairs. (bonded or nonbonded) around the central Table 4.6 (page114) shows the atom. arrangement of electron pairs about a Pairs of electrons in the valence shell repel central atom A (without any lone pairs) and one another since their electron clouds are geometries of some molecules/ions of the type negatively charged. AB. Table 4.7 (page 115) shows shapes of some simple molecules and ions in which the central These pairs of electrons tend to occupy atom has one or more lone pairs. Table 4.8 such positions in space that minimise (page 116) explains the reasons for the repulsion and thus maximise distance distortions in the geometry of the molecule. between them. The valence shell is taken as a sphere As depicted in Table 4.6, in the with the electron pairs localising on the compounds of AB2, AB3, AB4, AB5 and AB6, spherical surface at maximum distance the arrangement of electron pairs and the from one another. B atoms around the central atom A are : linear, trigonal planar, tetrahedral, A multiple bond is treated as if it is a single trigonal-bipyramidal and octahedral, electron pair and the two or three electron respectively. Such arrangement can be seen pairs of a multiple bond are treated as a in the molecules like BF3 (AB3), CH4 (AB4) and single super pair. PCl5 (AB5) as depicted below by their ball and Where two or more resonance structures stick models. can represent a molecule, the VSEPR model is applicable to any such structure. The repulsive interaction of electron pairs decrease in the order: Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp) Fig. 4.6 The shapes of molecules in which central atom has no lone pair Nyholm and Gillespie (1957) refined the VSEPR model by explaining the important The VSEPR Theory is able to predict difference between the lone pairs and bonding geometry of a large number of molecules, pairs of electrons. While the lone pairs are especially the compounds of p-block elements localised on the central atom, each bonded accurately. It is also quite successful in pair is shared between two atoms. As a result, determining the geometry quite-accurately the lone pair electrons in a molecule occupy even when the energy difference between more space as compared to the bonding pairs possible structures is very small. The of electrons. This results in greater repulsion theoretical basis of the VSEPR theory regarding between lone pairs of electrons as compared the effects of electron pair repulsions on to the lone pair - bond pair and bond pair - molecular shapes is not clear and continues bond pair repulsions. These repulsion effects to be a subject of doubt and discussion. 2024-25 Unit 4.indd 113 9/12/2022 9:36:12 AM  114 chemistry Table 4.6 Geometry of Molecules in which the Central Atom has No Lone Pair of Electrons 2024-25 Unit 4.indd 114 9/12/2022 9:36:13 AM  Chemical Bonding And Molecular Structure 115 Table 4.7 Shape (geometry) of Some Simple Molecules/Ions with Central Ions having One or More Lone Pairs of Electrons(E). 2024-25 Unit 4.indd 115 9/12/2022 9:36:13 AM  116 chemistry Table 4.8 Shapes of Molecules containing Bond Pair and Lone Pair Molecule No. of No. of Arrangement Shape Reason for the type bonding lone of electrons shape acquired pairs pairs AB2E 4 1 Bent Theoretically the shape should have been triangular planar but actually it is found to be bent or v-shaped. The reason being the 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°. AB3E 3 1 Trigonal Had there been a bp in place of lp the shape would have pyramidal been tetrahedral but one lone pair is present and due to the repulsion between lp-bp (which is more than bp-bp repulsion) the angle between bond pairs is reduced to 107° from 109.5°. Bent The shape should have been AB2E2 2 2 tetrahedral if there were all bp but two lp are present so the shape is distorted tetrahedral or angular. The reason is lp-lp repulsion is more than lp-bp repulsion which is more than bp-bp repulsion. Thus, the angle is reduced to 104.5° from 109.5°. AB4E 4 1 See- In (a) the lp is present at axial saw 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 (More stable) or a see-saw. 2024-25 Unit 4.indd 116 9/12/2022 9:36:13 AM  Chemical Bonding And Molecular Structure 117 Molecule No. of No. of Arrangement Shape Reason for the type bonding lone of electrons shape acquired pairs pairs AB3E2 3 2 T-shape 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). 4.5 Valence Bond Theory of the valence bond theory is based on the As we know that Lewis approach helps in knowledge of atomic orbitals, electronic writing the structure of molecules but it configurations of elements (Units 2), the fails to explain the formation of chemical overlap criteria of atomic orbitals, the bond. It also does not give any reason for the hybridization of atomic orbitals and the principles of variation and superposition. A difference in bond dissociation enthalpies and rigorous treatment of the VB theory in terms bond lengths in molecules like H2 (435.8 kJ of these aspects is beyond the scope of this mol-1, 74 pm) and F2 (155 kJ mol-1, 144 pm), book. Therefore, for the sake of convenience, although in both the cases a single covalent valence bond theory has been discussed in bond is formed by the sharing of an electron terms of qualitative and non-mathematical pair between the respective atoms. It also treatment only. To start with, let us consider gives no idea about the shapes of polyatomic the formation of hydrogen molecule which is molecules. the simplest of all molecules. Similarly the VSEPR theory gives the Consider two hydrogen atoms A and B geometry of simple molecules but theoretically, approaching each other having nuclei NA it does not explain them and also it has limited and NB and electrons present in them are applications. To overcome these limitations represented by eA and eB. When the two atoms the two important theories based on quantum are at large distance from each other, there mechanical principles are introduced. These is no interaction between them. As these two are valence bond (VB) theory and molecular atoms approach each other, new attractive orbital (MO) theory. and repulsive forces begin to operate. Valence bond theory was introduced Attractive forces arise between: by Heitler and London (1927) and developed (i) nucleus of one atom and its own electron further by Pauling and others. A discussion that is NA – eA and NB– eB. 2024-25 Unit 4.indd 117 9/12/2022 9:36:13 AM  118 chemistry (ii) nucleus of one atom and electron of together to form a stable molecule having the other atom i.e., NA– eB, NB– eA. bond length of 74 pm. Similarly repulsive forces arise between Since the energy gets released when the (i) electrons of two atoms like eA – eB, bond is formed between two hydrogen atoms, the hydrogen molecule is more stable than (ii) nuclei of two atoms NA – NB. that of isolated hydrogen atoms. The energy Attractive forces tend to bring the two so released is called as bond enthalpy, which atoms close to each other whereas repulsive is corresponding to minimum in the curve forces tend to push them apart (Fig. 4.7). depicted in Fig. 4.8. Conversely, 435.8 kJ of energy is required to dissociate one mole of H2 molecule. H2(g) + 435.8 kJ mol–1 → H(g) + H(g) Fig. 4.8 The potential energy curve for the formation of H2 molecule as a function of internuclear distance of the H atoms. The minimum in the curve corresponds to the most stable state of H2. 4.5.1 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 Fig. 4.7 Forces of attraction and repulsion orbitals undergo partial interpenetration. This during the formation of H2 molecule partial merging of atomic orbitals is called overlapping of atomic orbitals which results in Experimentally it has been found that the pairing of electrons. The extent of overlap the magnitude of new attractive force is decides the strength of a covalent bond. In more than the new repulsive forces. As a result, two atoms approach each other and general, greater the overlap the stronger is the potential energy decreases. Ultimately a stage bond formed between two atoms. Therefore, is reached where the net force of attraction according to orbital overlap concept, the balances the force of repulsion and system formation of a covalent bond between two acquires minimum energy. At this stage atoms results by pairing of electrons present two hydrogen atoms are said to be bonded in the valence shell having opposite spins. 2024-25 Unit 4.indd 118 9/12/2022 9:36:13 AM  Chemical Bonding And Molecular Structure 119 4.5.2 Directional Properties of Bonds As we have already seen, the covalent bond is formed by overlapping of atomic orbitals. The molecule of hydrogen is formed due to the overlap of 1s-orbitals of two H atoms. In case of polyatomic molecules like CH4, NH3 and H2O, the geometry of the molecules is also important in addition to the bond formation. For example why is it so that CH4 molecule has tetrahedral shape and HCH bond angles are 109.5°? Why is the shape of NH3 molecule pyramidal ? The valence bond theory explains the shape, the formation and directional properties of bonds in polyatomic molecules like CH4, NH3 and H2O, etc. in terms of overlap and hybridisation of atomic orbitals. 4.5.3 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 (Fig. 4.9). Positive and negative sign on boundary surface diagrams in the Fig. 4.9 show the sign (phase) of orbital wave function and are not related to charge. Orbitals forming bond should have same sign Fig.4.9 Positive, negative and zero overlaps of s and p atomic orbitals (phase) and orientation in space. This is called positive overlap. Various overlaps of s and p hydrogen. The four atomic orbitals of carbon, orbitals are depicted in Fig. 4.9. each with an unpaired electron can overlap with the 1s orbitals of the four H atoms which The criterion of overlap, as the main factor are also singly occupied. This will result in the for the formation of covalent bonds applies formation of four C-H bonds. It will, however, uniformly to the homonuclear/heteronuclear be observed that while the three p orbitals of diatomic molecules and polyatomic molecules. carbon are at 90° to one another, the HCH We know that the shapes of CH4, NH3, and angle for these will also be 90°. That is three H2O molecules are tetrahedral, pyramidal C-H bonds will be oriented at 90° to one and bent respectively. It would be therefore another. The 2s orbital of carbon and the 1s interesting to use VB theory to find out if these orbital of H are spherically symmetrical and geometrical shapes can be explained in terms they can overlap in any direction. Therefore of the orbital overlaps. the direction of the fourth C-H bond cannot Let us first consider the CH4 (methane) be ascertained. This description does not fit molecule. The electronic configuration of in with the tetrahedral HCH angles of 109.5°. carbon in its ground state is [He]2s2 2p2 which Clearly, it follows that simple atomic orbital in the excited state becomes [He] 2s1 2px1 2py1 overlap does not account for the directional 2pz1. The energy required for this excitation is characteristics of bonds in CH4. Using similar compensated by the release of energy due to procedure and arguments, it can be seen that in overlap between the orbitals of carbon and the the case of NH3 and H2O molecules, the HNH 2024-25 Unit 4.indd 119 9/12/2022 9:36:14 AM  120 chemistry and HOH angles should be 90°. This is in above and below the plane of the disagreement with the actual bond angles of participating atoms. 107° and 104.5° in the NH3 and H2O molecules respectively. 4.5.4 Types of Overlapping and 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 4.5.5 Strength of Sigma and pi Bonds bond is formed by the end to end (head- Basically the strength of a bond depends on) overlap of bonding orbitals along the upon the extent of overlapping. In case of internuclear axis. This is called as head sigma bond, the overlapping of orbitals takes on overlap or axial overlap. This can be place to a larger extent. Hence, it is stronger formed by any one of the following types as compared to the pi bond where the extent of combinations of atomic orbitals. of overlapping occurs to a smaller extent. s-s overlapping : In this case, there is Further, it is important to note that in the overlap of two half filled s-orbitals along formation of multiple bonds between two the internuclear axis as shown below : atoms of a molecule, pi bond(s) is formed in addition to a sigma bond. 4.6 Hybridisation In order to explain the characteristic geometrical shapes of polyatomic molecules s-p overlapping: This type of overlap like CH4, NH3 and H2O etc., Pauling introduced occurs between half filled s-orbitals of one the concept of hybridisation. According to him atom and half filled p-orbitals of another the atomic orbitals combine to form new set of atom. equivalent orbitals known as hybrid orbitals. Unlike pure orbitals, the hybrid orbitals are used in bond formation. The phenomenon is known as hybridisation which can be defined as the process of intermixing of the orbitals of p–p overlapping : This type of overlap slightly different energies so as to redistribute takes place between half filled p-orbitals their energies, resulting in the formation of of the two approaching atoms. new set of orbitals of equivalent energies and shape. For example when one 2s and three 2p-orbitals of carbon hybridise, there is the formation of four new sp3 hybrid orbitals. Salient features of hybridisation: The main (ii) pi( ) bond : In the formation of π bond features of hybridisation are as under : the atomic orbitals overlap in such a way that their axes remain parallel to 1. The number of hybrid orbitals is equal to each other and perpendicular to the the number of the atomic orbitals that get internuclear axis. The orbitals formed hybridised. due to sidewise overlapping consists 2. The hybridised orbitals are always of two saucer type charged clouds equivalent in energy and shape. 2024-25 Unit 4.indd 120 9/12/2022 9:36:14 AM  Chemical Bonding And Molecular Structure