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BASIC CONCEPTS OF INORGANIC CHEMISTRY (Second Edition) D N Singh (M.Sc., Ph.D) Formerly, Reader in Chemistry P.G.M.S College, Motihari (B.R Ambedkar Bihar University Muzaffarpur) &KDQGLJDUK'HOKL&KHQQDL The aim of this publication is to supply information taken from sources believed to be valid and reliable This is not an attempt to render any type of professional advice or analysis, nor is it to be treated as such While much care has been taken to ensure the veracity and currency of the information presented within, neither the publisher nor its authors bear any responsibility for any damage arising from inadvertent omissions, negligence or inaccuracies (typographical or factual) that may have found their way into this book Copyright © 2012 Dorling Kindersley (India) Pvt Ltd Licensees of Pearson Education in South Asia No part of this eBook may be used or reproduced in any manner whatsoever without the publisher’s prior written consent This eBook may or may not include all assets that were part of the print version The publisher reserves the right to remove any material present in this eBook at any time ISBN 9788131768617 eISBN 9788131798683 Head Office: A-8(A), Sector 62, Knowledge Boulevard, 7th Floor, NOIDA 201 309, India Registered Office: 11 Local Shopping Centre, Panchsheel Park, New Delhi 110 017, India DEDICATED TO “MY GURUS” This page is intentionally left blank CONTENTS Preface Periodic Table and Periodicity of Properties xi Mendeleev’s periodic law Modern Periodic Table IUPAC table Periodic table and Aufbau principle Determination of group and period of an element Classification of elements Atomicity and elements of the periodic table Effective nuclear charge Atomic and ionic size 10 Ionization energy 17 Electron affinity 19 Electronegativity 21 Physical properties and the periodic table 23 Artificial elements 29 Some records of periodic table 30 Chemical Bonding and Molecular Structure Different types of chemical bonds Summery of bond types Ionic bond Born – Haber cycle Cations of stable electron configuration Polarization and its effects Lattice energy General properties of ionic compounds Covalent bonding σ and π bonds Comparison of σ and π bond Electronegativity Coordinate covalent bond Electronegativity and Dipole moment Lewis structure (or Dot structure) Resonance structure Hybridization Resonance Resonance effect Resonance energy 33 33 34 34 37 38 38 40 42 42 43 44 46 47 48 52 54 55 61 63 63 Bond length Bond energy VSEPR model Structure and shape Shape of Molecules Bond angle Odd electron molecules Molecular orbital model Shape and symmetry of molecular orbitals Formation of π bonds Diatomics of the first period elements Homonuclear Diatomics of Second Period Elements NO molecule CO molecule Hydrogen bond Van der Waals, forces Metallic bond Metal structures Electron gas model Band model Cohesive energy Acids and Bases Bronsted–Lowry theory Strength of acids and bases PH Buffer solution Henderson equation Amphoterism Acid strength Strength of hydra acids Strength of Oxy acids Base strength Strength of hydra bases Strength of hydroxide bases Lewis Acid – Base theory Hard and soft acids and bases Chemical Reaction Types of chemical reactions Hydrolysis 64 66 68 68 68 72 76 76 78 79 80 81 87 88 89 94 95 95 96 97 98 102 102 103 105 106 106 107 109 109 110 113 114 115 116 118 122 123 126 Oxidation–reduction reactions Oxidation number Oxidizing agents Reducing agents Oxidizing and reducing agents Strength of oxidants and reductants Equivalent weights of oxidizing and reducing agents Balancing of redox reactions Oxidation number method Ion–electron method Some important half reactions Transition Elements General properties Atomic and ionic radii Bonding in Transition metals and its effect on properties Electrode potential Oxidation states Paramagnetic nature of transition metal compounds Complex compound formation Colour of transition metal compounds The d – d transition Colour and charge transfer Hydrolysis of transition metal compounds Catalytic property Lanthanides and actinides Coordination Chemistry Coordination number Types of ligands Chelates Conditions for complex formation Werner’s coordination theory Nomenclature Valence bond model for complexes Crystal field model of bonding Stability and CF model Magnetic properties and CF model Colour and CF model Isomerism in coordination compounds Geometrical isomerism Optical isomerism Preparations of complex compounds Stability of complex compound in a solution Applicability of complex compounds 129 130 132 132 133 134 136 137 137 137 141 144 146 146 147 148 148 151 152 152 152 154 154 155 160 166 166 166 167 167 168 171 172 176 178 180 181 183 185 188 191 191 192 Organometallic compounds Preparation of organometallies Bonding in organometallic compounds Bonding in alkene complexes Abundance and metallurgy Chemical elements in the Earth’s crust Cosmic abundance of elements Abundance in oceans Occurrence of metals Metallurgy Terms used in metallurgical process Concentration of ore Leaching Isolation of metal from concentrated ore Thermal (or chemical reduction) Auto reduction Electrolytic method of reduction Displacement of one metal by the other Purification of isolated metals Sodium Magnesium Calcium Aluminium Iron Ashoka Pillar at Delhi Pig iron Grey cast iron White cast iron Wrought iron Steel Special steel Conversion of iron into steel Steel from wrought iron Comparison Pig iron, Wrought and Steel Heat treatments of steel Chemically pure iron Brief chemistry of iron Corrosion of iron Important compounds of iron Copper Alloys of copper Brief chemistry of Cu Important compounds and complexes Zinc Brief chemistry of Zn Important compounds Mercury Brief chemistry of Hg 192 194 195 196 201 201 201 202 202 203 203 205 207 207 208 208 209 209 209 211 212 214 215 218 218 220 220 220 220 221 221 221 223 224 224 225 225 228 228 230 232 233 235 236 237 239 239 240 Important compounds Tin Allotropic forms Chemical reaction Lead Physical properties Important compounds Important Chemical Compounds Metal compounds NaOH Na2Co3 K2Cr2O7 KMnO4 Non-metal compounds O3 H2O2 NA2S2O3 H2S Hydrogen and Its Chemistry 242 243 244 245 246 247 248 252 252 255 258 261 265 266 266 268 273 274 278 Position in the periodic table Isotopes of H Oxidation states and bonding Preparation of H2 Uyeno’s reaction Bosch’s process Lane’s process Laboratory Preparation of H2 Nascent Hydrogen Compounds of protium (H) Water Zeolite water Water clathrates Hard and soft water Temporary hardness Permanent hardness Inorganic exchangers Organic exchangers Bad effects of hard water Structure of water and ice Density of water and ice Density of water at 4oC Heavy water 278 278 280 281 282 283 283 284 284 285 286 289 289 289 290 290 291 291 291 292 293 293 293 10 Group – 1(IA) The alkali metals 299 Chemical reactions OXO salts Halides Flame colour 301 302 303 305 Alkali metals and liquid NH3 Anomalous behaviour of Li 305 306 11 Group – 2(IIA) [Be, Mg, Ca, Sr, Ba, Ra] 309 Properties which decrease down the group Properties which increase down the gorup Oxidation states and nature of bond Hydrides Halides Oxides and hydroxides Oxo salts Flame colouration 12 Group – 11(IIB) Cu, Ag, Au Metallic bond strength Sublimation energy Atomic and ionic radii Ionization energy Noble metal nature Malleability, thermal and electrical conductivities Oxidation states Magnetic properties Colour of compounds Solubility of Silver–Hlides Chemistry of photography 13 Group – 12(IIB) Zn, Cd, Hg 310 311 311 311 312 314 316 318 321 322 322 322 322 322 323 323 324 325 326 326 329 Ionization energy Oxidation states Nature of bonds Electrode potential Magnetic properties Colour of compounds Some useful compounds Biochemistry of Zn, Cd and Hg 330 330 331 331 332 332 333 333 14 Group – 13 (IIIA) B, Al, Ga, In, Td 336 Oxidation states and nature of bond Hydrides Diborane Structure of B2H6 Borazole Boric acid Halides Lewis acid strength of BX3 Alums Isolation of B Crystalline B 337 339 339 340 342 345 346 347 347 348 348 15 Group – 14(IVA) C, Si, Ge, Sn, Pb Catenation Allotropy and structure Graphite Diamond Fullerenes Semiconductor property of Si and Ge Physical properties of group – 14 elements Oxidation states and bonding Carbides Oxides Cyanogens HCN Cyanides Halides Hydrides Silicones Silicates Isolation of Si 16 Group – 15(VA)N, P, As, Sb, Bi Allotropes of P Oxidation state and nature of bond Hydrides NH3 PH3 Oxides on N and P N2O NO N2O3 NO2 N2O5 P4O6 and P4O10 HNO2 HNO3 Aquaregia H3PO2 H3PO3 Phosphoric acids Acid strength of H3PO2, H3PO3 and P3PO4 Halides Isolation of N and P Fertilizers 351 351 352 352 353 353 354 355 355 356 357 360 360 361 362 364 364 366 368 372 373 374 375 377 378 378 379 379 380 380 381 381 382 383 385 386 386 387 389 389 392 393 17 Group–16(VIA) O, S, Se, Te, Po 399 Physical state of the elements Allotropy of O and S Effect of heat on S 399 400 401 Viscosity of liquid S and temperature Oxidation state and nature of bond Hydrides H2O2 Strength of H2O2 Acid strength of H2O2 and H2O Structure of H2O2 Halides SOCl2 Oxides Oxo acids 18 Group – 17 (VIIIA) Halogens F, Cl, Br, I, and At Physical state Special properties of F Oxidation state and bonding Formation of X2 Manufacture of Cl2 Manufacture of Br2 Manufacture of I2 Reactions of X2 Hydrogen halides HF HCl HBr Hl Halides Preparation of anhydrous halides Halogen oxides Oxo acids Acid Strength Oxidizing power CIO–n anions Halic acids Perhalic acids Interhalogen compounds Pseudohalogens and pseudohalides 401 401 403 404 405 405 406 406 406 408 414 420 420 422 422 423 424 425 425 426 427 428 429 430 430 431 433 433 435 436 436 437 439 440 441 443 19 Group – 18 The Noble Gases 447 Atomicity Radii Water solubility Special properties of He Uses of noble gas Clathrate Xe compounds Structure of Xe – Compounds 447 448 448 448 449 449 449 452 20 Analytical Chemistry Carbonates Sulphite Nitrite Chlorides Bromides Iodide Nitrate 454 455 456 459 460 461 462 463 Sulphate Tests for basic radicals Flame test Borax bead test Solution test for basic radicals Test of NH+4 ion 464 464 466 466 468 469 21 Problems on Inorganic Reactions 484 Additional Practice Questions 491 1RGDOSODQH ± VS ] S] V Figure 2.37 (c) Combination of p and p - orbitals The p – orbitals can combine in two different ways (i) At an axis and (ii) At parallel axes which are perpendicular on the bond axis It results into the formation of π bonds When two p orbitals combine at an axis (say on z – axis) two MOs are formed One of these is bonding σ(p – p)MO which has cylindrical symmetry around the bond axis The antibonding, σ *( p - p ) has no cylindrical symmetry around the bond axis It has a nodal plane perpendicular to the bond axis S ] S ] S ] S ] S] S ] V S±S S ] V S±S S ] 1RGDOSODQH Figure 2.38 Formation of π bonds When two p–orbitals (px and px or py and py) at parallel axes combine a π-bond is formed The π bonds are always formed in pairs (i.e., π p x and π p y ) because two set of p – orbitals (px – px and py – py) always combine Thus, two equivalent (degenerate) π bonding MOs and two equivalent π*, anti-bonding MOs are formed 80 Basic Concepts of Inorganic Chemistry S [ S [ RU S \ S \ S [±S [ ± ± ± 1RGDOD[LV QRGDOSODQH ± RU S \ ±S \ S [RUS \ ± S [ RUS \ ± ± 1RGDOD[LV QRGDOSODQH Figure 2.39 The π bonding MO has a nodal plane This plane lies in the plane of the molecule dividing entire π MO in two parts i.e., above and below the nodal plane For example, in C2H4 the nodal plane lies in the molecular plane and π MO is above and below this plane + + & & 1RGDO D[LV ± + + Figure 2.40 The π* MO also has a nodal plane but it is perpendicular at the internuclear axis (or bond axis) indicated in the diagram of π* MO Homonuclear Diatomic Molecules Diatomic molecules of the same two atoms are homonuclear (Homo–‘the same’) diatomics For example, H2, N2, C2, O2, F2 etc Diatomics of the first period elements The first period elements are H and He and their diatomics are H2, H +2 , He2, He +2 etc The valence orbital for these elements is the 1s orbital only Therefore, a diatomic molecule will involve only two 1s atomic orbitals for the formation of a diatomic molecule The 1s atomic orbitals centred at two atoms combine to form one σ MO and the other σ* MO V V H2 V + VV + Figure 2.41 V + Chemical Bonding and Molecular Structure 81 Figure 2.41 is the MO diagram of H2 The H2 molecule can also be represented as H1s1 + H1s1 H2σ 12s The σ 12s presentation of H2 is its molecular orbital configuration Thus, bond order in H2, B.O = 2−0 =1 As there is no unpaired electron in H2, the molecule is diamagnetic Table 2.43 H2, H+2 , He2 and He +2 molecules Property Molecules H +2 Molecular orbitals H2 He +2 He2 σ 1*s σ1s Bond order 0.5 1.0 0.5 Bond length (Å) Bond energy (KJmol−1) 1.06 256 0.74 435 1.08 230 Magnetic property, µ(BM) Paramag Diamag paramag Stability order H2 > H +2 > He +2 3 Note: • • • • The He2 molecule is not possible because bond order is zero Molecular cation is formed, like atom, by the loss of electron from the highest occupied molecular orbital (in He +2 from σ 1*s MO) The diatomic system in the case of noble gas (e.g., He +2 ) is possible only in the excited state Such species are often called exonomers The bond energy in He +2 is smaller than H +2 due to the presence of electron in the anti-bonding (σ*) molecular orbital in He +2 Homonuclear Diatomics of Second Period Elements The second period elements are Li, Be, B, C, N, O, F and Ne Their diatomics are Li2, Be2, B2 etc The valence orbitals in second period elements is 2s and 2p i.e., a total of four orbitals, 2s, 2px, 2py and 2pz The 2s and 2p orbitals differ in energy but the three p–orbitals (2px, 2py and 2pz) are degenerate orbitals Let us first assume that z – axis is bond axis A combination of 2s atomic orbitals form two σ molecular orbitals, one bonding σ2s and the other antibonding σ *2s If z – axis is taken as bond axis then 2pz orbitals will overlap to form two σ - type molecular orbitals, one bonding, σ p z and the other anti-bonding σ *2 p z Now the overlap of two 2px or two 2py will give π - type molecular orbitals 82 Basic Concepts of Inorganic Chemistry A overlap of 2px and 2px will give π p bonding and π *2 p x anti-bonding MOs Similarly 2py and 2py x overlap will form bonding π p y and anti-bonding π *2 p y MOs As the type of overlap is similar and 2px and 2py orbitals are equal in energy, the π p x and π p y molecular orbitals are degenerate (i.e, of equal energy) so also the π *2 p x and π *2 p y MOs Spectroscopic data show that MO energy pattern is not similar for all the homonuclear diatomics of the second period (This difference is due to difference in the energy gap between 2s and 2p atomic orbitals from elements to elements) It is found that: (i) Energy sequence for diatomics from Li2 to N2 is, σ2s < σ *2s < π p x = π p y < σ p z < π *2 p x = π *2 p y < σ *2 p z and (ii) For diatomics from O2 to F2 the energy order is σ2s < σ *2s < σ p z < π p x = π p y < π *2 p x = π *2 p y < σ *2 p z The inner 1s orbitals overlap to form non – bonding MOs and so are not considered MO diagram for Li2 to N2 V 3 ] S S[ S S\ S S V S] ( S S[ S S\ V V V V $2 $2 V V 02 Figure 2.42 Chemical Bonding and Molecular Structure 83 MO diagram for O2 and F2 V 3 S S[ ] S S\ S S S S[ S S\ ( V S] V V V V $2 $2 V V 02 Figure 2.43 Li2 Molecule Li has valence orbitals configuration 2s1 A Li2 molecule will be formed by the combination of two 2s orbitals Li2 KK σ 22s Li(K)2s1 + Li(K)2s1 /L V /L V /L V V V V V V ± %2 V V Figure 2.44 It has no unpaired electron Li2 is thus a diamagnetic molecule Li2 is found in vapour state (The 1s orbitals will constitute non-bonding MOs It is represented in the MO configuration of Li2 as K and K) 84 Basic Concepts of Inorganic Chemistry Be2 molecule V V The valence orbital electron configuration of Be is 2s2 Be2 KKσ 22s σ *22s Be(K)2s2 + Be (K)2s2 B.O = 2−2 =0 V %H Therefore, diatomic Be2 is not possible The MO diagram is V %H V V The σ 2s and σ *2s are non-bonding MOs Figure 2.45 The 1s orbitals will form non-bonding MOs (as in Li2) B2 , C2 , N2 molecules Table 2.44 Property B2 C2 N2 σ *2 p z π *2 p x π *2 p y π *2 p z π 2p π 2p x y σ *2s σ 2s Property B2 C2 N2 Bond order Bond length (Å) 1.59 1.31 1.1 Bond energy (KJmol−1) 288 627 940 No of unpaired electrons – – Magnetism Paramag Diamag diamag 0 µ(BM) Note: • • • • Bond order increases (B2 to C2 to N2) because number of electrons in bonding molecular increases Bond energy increases when bond order increases Bond length decreases when bond order increases s and p mixing takes place as energy gap between 2s and 2p is small Molecular ions N +2 , N -2 , (N2) Molecular cations are formed when electrons are removed from highest energy occupied molecular orbitals (HOMO) Chemical Bonding and Molecular Structure 85 Molecular anions are formed when electrons are added in lowest energy unoccupied molecular orbitals Table 2.45 Property N2 N +2 N2– σ *2 p z π *2 p x π *2 p y π *2 p z π 2p π 2p x y σ *2s σ 2s Bond order 2.5 1.1 1.12 Bond energy (KJmol ) 940 828 No of unpaired electrons 1 Magnetism Diamag Paramag Paramag µ(BM) Stability order N2 > N +2 > N -2 Bond length (Å) −1 2.5 3 O2, F2 and Ne2 Molecules Table 2.46 Property O2 F2 Ne2 σ *2 p z π *2 p x π *2 p y π 2p π 2p x y s 2px σ *2s σ 2s Bond order – (Continued) 86 Basic Concepts of Inorganic Chemistry Property O2 F2 Ne2 Bond length (Å) Bond energy (KJmol−1) 1.21 494 1.42 159 – – No of unpaired electrons – Magnetism Paramag Diamag à(BM) Note ã ã ã Diatomic Neon is not possible as bond order is zero Noble gases are monoatomic bond energy sharply decreasses from O2 to F2 It is because of the presence of more antibonding electrons in F2 No s and p mixing takes place Table 2.47 O +2 O2 O -2 O =2 σ 2s Bond order 2.5 2.0 1.5 1.0 Bond length (Å) 1.12 1.21 1.26 1.49 Bond energy (KJmol−1) 642 494 394 210 No of unpaired electrons Magnetism µ(BM) Paramag Paramag Paramag Diamag Stability order + =2 O > O2 > O > O Property σ *2 p z π *2 p x π *2 p y π p π 2p y x 62pz σ *2s Note • For O +2 electron is removed from π* • For O -2 and O =2 electrons are added in π* Chemical Bonding and Molecular Structure 87 • The ion O -2 is called superoxide ion (K, Rb and Cs form superoxides, MO2 They are paramagnetic and coloured • The O =2 is known as peroxide ion Heteronuclear Diatomics $QWLERQGLQJ02 In a heteronuclear diatomic (Hetero–means different) two different atoms are bonded together (AB), e.g., HCl, CO, NO Such species can be treated using LCAO – MO concept similar to homonuclear diatomics (A2) But since the atoms are different – $ ( (i) The energy of the atomic orbitals are different (ii) Their relative contributions to the molecular orbitals are also different (iii) The bonding MOs are closer in energy to the more electronegative atom (iv) The antibonding MOs are closer in energy to the less electronegtive atom % %RQJ02 Figure 2.46 MO diagram The orbitals of more electronegative atom are more stable They are, therefore, kept at lower level when sketching MO diagram The qualitiative diagram may represented as (where χB > χA) The value of b, i.e., the difference in the electronegativity of A and B determines the polarity of the bond V S HCl Molecule In HCl, the 1s orbital of H combines with 3pz orbitals of Cl to form σ - type bonding and anti-bonding MOs The Cl, 3s, 3px and 3py orbitals remain nonbonding (no orbital at H to combine with) • • • E QE V &O QE V Figure 2.47 The bonding MO is concentrated near Cl, so the bond is polar The bond order is one, as there is no elelctron in σ* MO The HCl molecule is diamagnetic as there is no unpaired electron NO molecule Nitrogen and Oxygen belong to the second period and not differ widely in electronegativity Therefore, MO energy levels may be taken either similar to N2 (plus on electron) or O2 (minus one electron) N K 2s 2p1x 2p1y 2p1z O K 2s 2p 2x 2p1y 2p1z ... H3BO3 H2CO3 (not stable) HNO2 HNO3 O2 and HF Nature of oxide Basic (alkaline) Amphoteric Acidic 28 Basic Concepts of Inorganic Chemistry Oxides of third period elements Table 1.32 Formula Na2O... secondary forces, 34 Basic Concepts of Inorganic Chemistry I H – bond and II Van der Waals’ force D Multi-centre covalent bonding Summery of bond types Table 2.1 Types of bond Mechanism of formation... Physical state of the elements Allotropy of O and S Effect of heat on S 399 400 401 Viscosity of liquid S and temperature Oxidation state and nature of bond Hydrides H2O2 Strength of H2O2 Acid