Title 1A.CHEMICAL BONDING FINAL (DN & NLR).pdf ✅

NISHITH Multimedia India (Pvt.) Ltd., 1 JEE MAINS - CW - VOL - I CHEMICAL BONDING NISHITH Multimedia India (Pvt.) Ltd., JEE ADVANCED - VOL-II SYNOPSIS * Introduction Chemical Bond is the physical process responsible for the attractive interactions between atoms and molecules, and that which confers stability to diatomic and polyatomic chemical compounds. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics. In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. The study on the “nature of forces that hold or bind atoms together to form a molecule” is required to gain knowledge of the following i) To know about how atoms of same element form different compounds combining with different elements. ii) To know why particular shapes are adopted by molecules. iii) To understand the specific properties of molecules or ions and the relation between the specific type of bonding in the molecules. * KOSSEL-LEWIS APPROACH TO CHEMICAL BONDING In 1916, W.Kossel and G.N.Lewis, separately developed theories of chemical bonding in order to understand why atoms combined to form molecules. Lewis introduced simple symbols to denote the electron present in the outer orbit of atom , these electrons are known as valence electrons. These symbols are known as electron dot symbols and the structure of compound is known as Lewis dot structure. The electron dot structure of Methane is given as H H H H C The number of dots around the symbol is equal to number of electrons. * IONIC BOND OR ELECTROVALENT BOND An ion is an atom or group of atoms which has acquired charge due to the loss or gain of one or more electrons. When an atom gains an electron to form a negative ion (anion), it results increase in size. On the other hand, when an atom loses an electron to give positive ion (cation), it will contract. The electron lost or gained is always from the outermost shell. When two atoms, one of which can lose one or more electrons to attain a noble gas configuration and the other can receive these electrons and thereby acquire a noble gas configuration, they are said to be bonded by an ionic bond. Since the loss or gain of electrons by atoms results in the formation of ions, ionic bond is formed when two ions interact with each other and are thus held together by electrostatic attraction. The formation of potassium chloride (KCl), is illustrated below. K(1s 2s 2p 3s 3p 4s ) K(1s 2s 2p 3s 3p ) 2 2 6 2 6 1electron 2 2 6 2 6 1 loses    2 2 6 2 5 2 2 6 2 6 gains 1 electron (Ar configuration) Cl (1s 2s 2p 3s 3p ) Cl (1s 2s 2p 3s 3p )   From the above illustrations, it is clear that the formation of an ionic compound is obviously related to the ease of formation of the cations and anions from the neutral atom, which depends on two main factors: i) Ionization energy: Lower the value of ionization energy of an atom, greater will be the ease of formation of the cation from it. ii) Electron affinity: Higher the electron affinity of an atom, greater the ease of formation of the anion from it. * Lattice Energy When one mole of an ionic solid is formed from its constituent gaseous ions, the energy released is called the lattice energy. * Energetics of Formation of Ionic Substances: The energy included in the formation of an ionic compound from its constituent elements may be considered as shown by the Born-Haber Cycle for the formation of one mole of sodium chloride from sodium and chlorine. CHEMICAL BONDING
CHEMICAL BONDING 2 NISHITH Multimedia India (Pvt.) Ltd., JEE ADVANCED - VOL-II NISHITH Multimedia India (Pvt.) Ltd.,     Na Na Na  e (g) I (g) S Sublimation (s) - A Dissociation Addition of e 2(g) (g) (g) 1/ 2D* E 1 Cl Cl Cl 2      Crystal formation (g) (g) (S) U Na Cl NaCl      Where S=heat of sublimation of sodium metal I =ionization energy of sodium D=heat of dissociation of molecular chlorine EA =electron affinity of chlorine, and U=lattice energy of sodium chloride The amount of heat liberated in the overall reaction is the heat of formation of sodium chloride. From the above H = S + I + 2 1 D – EA – U The most important of these energy terms are I, EA and U, since these are considerably greater than the remaining terms S and D. More the negative value of the heat of formation, greater would be the stability of the ionic compound produced. Thus on the basis of the above equation, formation of an ionic compound is favoured by a) low ionization energy (I) of the metal. b) high electron affinity (EA ) of the other element. c) higher lattice energy (U) of the resulting compound. * Lattice energy of an ionic compound is also determined with following Born-Lande equation and Kapustinskii Equation. Born-Lande equation: 2 0 0 . . . . 1 L 1 N A Z Z e E r n            If values are taken in S.I. units then 2 0 0 0 . . . . 1 1 4 . L N A Z Z e E  r n            Where 0   permitivity of free space 12 1 (8.85 10 . ) F m    N0 = Avogadro number A = Madelung constant, which depends upon the geometry of crystal. r = inter ionic distance e = charge of electron n = Born exponent(Mostly its value is 9) Z + & Z - = charge on cation & anion respectively Kapustinskii Equation: When it is not possible to know the value of Madelung constant, lattice energy can be calculated byKapustinskii Equation. 1 2 . 345 L 121,000 1 z z E d d                 Where = number of ions per formula unit 1 2 z z = Actual charge on ions with sign d = ( ) r r    i.e., sum of ionic radii in pm. EL = Lattice energy in KJ/mole Q) What is the lattice enthalpy of KNO3 . 3 ( 138 & 189 ) K NO r pm r pm      Solution: Using Kapustinskii equation 1 2 . 345 L 121,000 1 z z E d d                 2 ( 1) ( 1) 34.5 121,000 1 327 327 EL                    = 661.98KJ/mol. * Formation of Ions with Higher Charges: Formation of a cation with unit positive charge is easy if the first ionization energy is low as in the case of alkali metals. Alkaline earth metals ionizes in two successive steps. Mg  Mg+ + e– Mg+  Mg2+ + e– But energy needed to ionize alkaline earth metals are higher than alkali metals. However, dipositive ions like Mg2+, Ca2+, Sr2+ and Ba2+ are quite common. Formation of a tripositive ion like Al3+ requires much more energy (= 5138 kJ) which is not available ordinarily. Successive ionization energies of aluminium are:   Al Al  e E1 E1 = 577kJ E2 2 Al Al e      E2 = 1816 kJ
NISHITH Multimedia India (Pvt.) Ltd., 3 JEE MAINS - CW - VOL - I CHEMICAL BONDING NISHITH Multimedia India (Pvt.) Ltd., JEE ADVANCED - VOL-II Al Al e 2 3      E3 E3 = 2745kJ It is on this account that most of aluminium compounds are covalent. In solution, however, aluminium is known to give hydrated ions [Al.6H2O]3+. This is possible because of the high heat of hydration of Al3+. The energy liberated during hydration of ions is sufficient for ionization. Similarly, anions with unit negative charge (e.g. Cl– , Br– , I– ) are very common. This is because the electron affinity of these atoms is positive and quite high. Formation of anions carrying two units of negative charge (e.g. S2–, O2–) is not so easy as their electron affinities are negative i.e., energy is needed to add second electron. Formation of anions carrying three units of negative charge (e.g. N3–, P3–) is almost rare. * Characteristics of Electrovalent Compounds Melting and Boiling Point: Due to the strong electrostatic force between the ions in a crystal of an electrovalent compound the energy needed to overcome these forces and break down the crystal lattice is more. Hence such compounds possess high melting and boiling points. * Electrical Conductivity: When an electrovalent compound is molten or dissolved in a solvent of high dielectric constant e.g., water, the binding forces in the crystal lattice disappear and the component ions become mobile. Under the influence of applied electrical field, the ions get charged and thus act as charge carrier of the current. Hence their molten forms or solutions conduct electricity. * Solubility: Ionic compounds are soluble in polar solvents like water because of molecules of the polar solvent interact strongly with the ions of the crystal and the solvation energy is sufficient to overcome the attraction between the ions in the crystal lattice. Dissolution is also favoured by the high dielectric constant of the solvents such as water, since this weakens the interionic attractions in the resulting solutions. Nonpolar solvents like benzene and carbon tetrachloride do not solvate the ions as their dielectric constants are low. Ionic compounds are, therefore insoluble in nonpolar solvents. Ionic compounds like sulphates and phosphates of barium and strontium are insoluble in water (because lattice energy is greater than hydration energy). This can be attributed to the high lattice energies of these compounds due to polyvalent nature of both the cation and the anion. In these cases, hydration of ions fails to liberate sufficient energy to offset the lattice energy. Illustration - 1 : Which one is correct for ionic bond (a) it is non directional (b) it is formed by the elements with same electronegativity (c) it is formed by overlapping of orbitals (d) Both (a) and (c) are correct Ans : a Solution: Ionic bond is non directional Illustration - 2 : which one is having highest hydration energy (a) Na+ (b) Li+ (c) Cs + (d) K+ Ans : (b) Li+ ion has highest hydration energy due to its small size * Bond Strength and Bond Length The average distance between two nuclei of the bonded atoms is measured in angstroms. Bond energy or bond strength is the amount of energy required to break one mole of covalent bonds in gas phase. Greater the bond energy of a bond more stronger the bond and more stable the bond. Bond length  1 / Bond strength Bond energy of  bond is greater than bond energy of  bond. Bond energy increases as the no. of bonds increase between the atoms. Bond Energies C C C C C C      Bond Lengths C C C C C C      Bond Energies C N C N C N      Bond Lengths C N C N C N      sp - sp < sp2 - sp2 < sp3 - sp3 (Bond Lengths) sp - sp > sp2 - sp2 > sp3 - sp3 (Bond Energies) Bond angle The angle between two bonds sharing a common atom is known as the bond angle. Bond angle depends on the geometry of the molecule. For
CHEMICAL BONDING 4 NISHITH Multimedia India (Pvt.) Ltd., JEE ADVANCED - VOL-II NISHITH Multimedia India (Pvt.) Ltd., example the bond angle in tetrahedral geometry is 109028’ while bond angle in planar geometry is 1200 . Bond angle also depends on the electronegativity of atoms bonded. For example the geometry of NH3 and NF3 is same but bond angle of NH3 is greater than NF3 because F has higher electronegativity than H. The electron pair is attracted more towards F in NF3 i.e. the bond pairs of electrons are away from N. The distance between bond pairs is more. Hence repulsion between bond pairs in NF3 is less than NH3 . As a result, the bond angle decreases to 102.4°. Whereas in NH3 it decreases to 107.3° only. N H H H N F F F 107.3° 102.4° Explain why bond angle of PH3 is less than that of PF3 . Solution : PH3 and PF3 are pyramidal in shape with one lone pair on P. But PF3 has greater bond angle than PH3 (opposite to NH3 and NF3 ). This is due to resonance in PF3 , leading to partial double bond character as shown below P F F F + As result repulsions between P – F bonds are large and hence the bond angle is large. There is no possibility of formation of double bonds in PH3 . Covalent Bond (By Mutual Sharing of Electrons) The covalent bond is formed when two atoms achieve stability by the sharing of an electron pair, each contributing one electron to the electron pair. The arrangement of electrons in a covalent molecule is often shown by a Lewis structure in which only valence shells (outer shells) are depicted. For sake of clarity, the electrons on different atoms are denoted by dots and crosses. * Polarity of Bonds: A covalent bond is set up by sharing of electrons between two atoms. It is further classified as polar or nonpolar depending upon the fact whether the electron pair is shared unequally between the atoms or shared equally. For example, the covalent bonds in H2 and Cl2 are called nonpolar as the electron pair is equally shared between the two atoms. H H Cl Cl Hydrogen molecule Chlorine molecule (Both formed by equal sharing of electrons between the atoms, i.e., by non-polar bonds) H F   d In the case of hydrogen fluoride the bond is polar as the electron pair is unequally shared. Fluorine has a greater attraction for electrons or has higher electronegativity than hydrogen and the shared pair of electrons is nearer to the fluorine atom than hydrogen atom. The hydrogen end of the molecule, therefore, appears positive with respect to fluorine. Bond polarities affect both physical and chemical properties of compounds containing polar bond. The polarity of a bond determines the kind of reaction that can take place at that bond and even affects the reactivity at nearby bonds. The polarity of bonds can lead to polarity of molecules and affect melting point, boiling point and solubility * POLARISATION AND FAJAN’S RULES: Consider an ionic bond formed between two oppositely charged ions of unequal size of A+ and B - . In this bond cation and anion remains bonded by electrostatic force of attraction bu the valence electrons of larger anion (B- ) are attracted by small cation(A+ ) and so the shape of valence shell is deformed and electron cloud of valence shell of anion remains no longer symmetrical. This process is called Polarisation of anion by cation. The ability of cation to polarise anion is called polarising power and the tendency of anion to get polasied, is called polarisability. A + B - A + B - (a) Normal Ionic bond (b) Polarised anion deform ed valence shell