Title PMT_Chem_Chemical_Bonding_and_Molecular_Refresher.pdf ✅

4 CHAPTER Chemical Bonding and Molecular Structure ■ Valency ■ Chemical Bond ■ Ionic or Kernel Bond ■ Covalent Bond ■ Coordinate or Semi-polar Bond ■ Hydrogen Bond ■ Metallic Bonding ■ Resonance ■ Hybridization ■ VSEPR (Valence Shell Electron Pair Repulsion Theory) ■ Molecular Orbital Theory Chapter Outline Valency Valency is a property of atoms whereby they form chemical bond among themselves. The term valency was introduced by Frankland and it means ‘power to combine.’ Hence, it is the power of an atom to combine with another atom. Atoms do so by either giving up or accepting electrons in their outermost shell. Modern or electronic concept of valency was given by Kossel and Lewis; it was completed by Langmuir. Valency (V) = No. of valence electrons For instance, the electronic configuration for the group IA element sodium (Na), is 2, 8, 1. Here, the number of valence electron is 1 and hence its valency is 1. If the number of valence electrons is more than 4, then we use the following relationship to determine the valency: V = V e– – 8 (number of valence electrons minus 8) For example, the configuration of nitrogen (N) is 2, 5. According to the above relationship, its valency will be given as V = 5 – 8 = –3 (negative sign signifies the tendency to accept electrons) Chemical Bond Chemical bond is the force of attraction that binds two at- oms together. A chemical bond balances the force of attrac- tion and force of repulsion at a particular distance. A chemical bond is formed to: • Attain the octet state • Minimize energy • Gain stability • Decrease reactivity When two atoms come close to each other, forces of at- traction and repulsion operate between them. The distance at which the attractive forces overcome repulsive forces is called bond distance. Here, potential energy for the system is lowest, hence the bond is formed. Types of Bonds Following are the six types of chemical bonds. Here, they are listed in a decreasing order of their respective bond strengths. 1. Ionic bond 2. Covalent bond 3. Coordinate bond
4.70 ■ Chapter 4 4. Metallic bond 5. Hydrogen bond 6. Van der Waals bond Metallic bond, hydrogen bond and van der Waals bond are interactions. Octet Rule It was introduced by Lewis and Kossel. According to this rule, each atom tries to obtain the octet state, that is, a state with eight valence electrons. Exceptions to the octet rule • Transition metal ions like Cr+3, Mn+2, Fe+2. • Pseudo inert gas configuration cations like Zn2+, Cd2+. Contraction of octet state • Here central atom is electron deficient or does not have an octet state. For example, BeX2 4 BX3 6 AX3 6 Ge(CH3 ) 3 6e– Expansion of octet state • Here central atom has more than 8 e– due to empty d orbital. For example, PCl5, SF6, OsF8, ICl3 • Odd electronic species like NO, NO2, ClO2 • Inter halogens compounds like IF7, BrF3 • Compounds of xenon such as XeF2, XeF4, XeF6 Ionic or Kernel Bond Ionic bond is formed by the complete transfer of valence electrons from a metal to a non-metal. This was first studied by Kossel. For example, Na + Cl Na+ Cl– (2, 8, 1) (2, 8, 7) (2, 8) (2, 8) Mg + O Mg+2 O–2 (2, 8, 2) (2, 6) (2, 8) (2, 8) Al + N A+3 N–3 (2, 8, 3) (2, 5) (2, 8) (2, 8) • Number of electrons transferred is equal to electrova- lency. • Maximum number of electrons transferred by a metal to non-metal is three, as is the case of AlF3 , (Al metal transfers three electrons to F). • During electron transfer, the outermost orbit of metal is destroyed and the remaining portion is called core or kernel, so this bond is also called kernel bond. • Nature of ionic bond is electrostatic or coloumbic force of attraction. • It is a non-directional bond. Conditions for the Formation of an Ionic Bond The process of bond formation must be exothermic ( H = –ve) and for it the essential conditions are • Metal must have low ionization energy. • Non-metal must have high electron affinity. • Ions must have high lattice energy. • Cation should be large with low electronegativity. • Anion must be small with high electronegativity. Born–Haber Cycle The formation of an ionic compound in terms of energy can be shown by Born–Haber cycle. It is also used to find lattice energy, ionization energy and electron affinity. For example, Hf = S + 1 2 D + I – E – U Here, S = Heat of sublimation D = Heat of dissociation I = Ionization enthalpy E = Electron gain enthalpy or electron affinity U = Lattice energy • For the formation of an ionic solid, energy must be released during its formation, that is, H must be nega- tive for it. –E – U > S + 1⁄2 D + I
Chemical Bonding and Molecular Structure ■ 4.71 Properties of Ionic Compounds 1. Ionic compounds have solid crystalline structures (flat surfaces), with definite geometry, due to strong electrostatic force of attraction as constituents are arranged in a definite pattern. 2. These compounds are hard in nature. Hardness ∝ Electrostatic force of attraction ∝ Charge on ion ∝ 1 Ionic radius 3. Ionic compounds have high value of boiling point, melt- ing point and density due to strong electrostatic force of attraction. Boiling point, melting point ∝ Electrostatic force of attraction Volatile nature ∝ 1 Electorstatic force of attraction 4. Ionic compounds show isomorphism, that is, they have same crystalline structure. For example, all alums, NaF and MgO. 5. These are conductors in fused, molten or aqueous state due to presence of free ions. In solid state, these are non- conductors as no free ions are present. 6. They show fast ionic reactions as activation energy is zero for ions. 7. They do not show space isomerism due to nondirectional nature of ionic bond. 8. Lattice energy (U) is released during the formation of an ionic solid molecule from its constituent ions. Lattice Energy Lattice energy is also the energy needed to break an ionic solid molecule into its constitutent ions. It is denoted by U. U ∝ Charge on ion ∝ 1 Size of ion Hence, lattice energy for the following compounds increases in the order shown below: NaCl < MgCl2 < AlCl3 < SiCl4 As charge on a metal atom increases, its size decreases. In case of univalent and bivalent ionic compounds, lattice energy decreases as follows: Bi-bi > Uni-bi or Bi-uni > Uni-uni For example, MgO > MgCl2 > NaCl. 9. Ionic compounds are soluble in polar solvents like water due to the high dielectric constant of these solvents, therefore, force of attraction between ions are destroyed and they dissolve in the solvent. Facts Related to Solubility l If ΔH (hydration) > Lattice energy then ionic com- pound is soluble. l If ΔH (hydration) < Lattice energy then ionic com- pound is insoluble l If ΔH (hydration) = Lattice energy then the com- pound is at equilibrium state Some Solubility Orders a. LiX < NaX < KX < RbX < CsX b. LiOH < NaOH < KOH < RbOH < CsOH c. BeX2 < MgX2 < CaX2 < BaX2 d. Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Ba(OH)2 e. BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4 f. AIF3 > AlCl3 > AlBr3 > AII3 l Crystals of high ionic charges are less soluble. For example, compounds of CO3 –2, SO4 –2, PO4 –3 are less soluble. Compounds Ba+2, Pb+2 are insoluble as lattice energy > ΔHhy Compounds of Ag (salt) are insoluble as lattice energy > ΔHhy l Presence of common ions decrease solubility. For example, solubility of AgCl decreases in presence of AgNO3 or KCl, due to presence of common ions that is, Ag+ and Cl– respectively. Covalent Bond A covalent bond is formed by equal sharing of electrons between two similar or different atoms.