CY 211 - Basic Inorganic Chemistry - Chapter 4 PDF

CY 211 - Basic Inorganic Chemistry and Thermodynamics II Prof. K. Muralidharan School of Chemistry University of Hyderabad Chapter 4 – Selected topics in chemistry of s & p block elements What is main group elemen...

CY 211 - Basic Inorganic Chemistry and Thermodynamics II Prof. K. Muralidharan School of Chemistry University of Hyderabad Chapter 4 – Selected topics in chemistry of s & p block elements What is main group elements? Why? International Union of Pure and Applied Chemistry (iupac.org) What is main group elements? Why? The main group elements are the most abundant elements on Earth, in the Solar System. They comprise 80% of the Earth’s crust. For this reason, the main group elements are also called the representative elements. Top ten most abundant element on earth, eight of them are in human body. These elements are critical for supporting life. Biological molecules require main group elements, particularly carbon, oxygen, hydrogen, nitrogen, sulfur, and phosphorus. They are fundamental constituents of ordinary matter with a rich and interesting chemistry. The main group elements and their compounds are among the most economically important elements. The majority of manufactured products contain these elements. Hydrogen Bonding Introduced by: Maurice Loyal Huggins (1919), Published by: Wendell Latimer and Worth Rodebush in 1920. Chemical Laboratory of the University of California, Berkeley. H-bond usually occurs when proton donor is sufficiently electronegative to enhance the acidic nature of H and where the acceptor has regions of high electron density (lone pair electron) which can interact strongly with acidic hydrogen Bond energy = 10-60 kJ/mol XHY where atom Y may or may not be the same as X. It is not necessary for the electronegative atom X to be highly electronegative for there to be a meaningful hydrogen-bonded interaction. Thus, in addition to hydrogen bonds of the type F─H∙∙∙∙∙F, O─H∙∙∙∙∙F, N─H∙∙∙∙∙F, O─H∙∙∙∙∙O, N─H∙∙∙∙∙O, O─H∙∙∙∙∙N and N─H∙∙∙∙∙N, it is now well recognized that weaker hydrogen bonds, in particular C─H∙∙∙∙∙O interactions, play an important role in the solid state structures of small molecules and biological systems. The H bond in A─H∙∙∙∙∙B can be either linear as in schematic structure (1) or significantly non linear as in structures (lb) and (IC). H-bonds can also join three adjacent atoms (bifurcated) as in structure (2) or even four atoms (trifurcated) as in structure (3).  The nature of H bond in A─H∙∙∙∙∙B linear as in schematic structure (1) non linear as in structures 1a and 1b H-bonds can also join three adjacent atoms (bifurcated) as in structure (2) Join our atoms (trifurcated) as in structure (3). strengths of hydrogen bonds can be obtained experimentally comes from the dissociation of a carboxylic acid dimer in the vapour state (equation 9.25). The position of equilibrium 9.25 is temperature-dependent, and ΔHo for the reaction can be obtained from the variation of Kp with temperature: For formic acid (methanoic acid), ΔHo for the dissociation in equation 9.25 is found to be +60 kJ mol-1, or the value can be expressed as +30 kJ per mole of hydrogen bonds. This quantity is often referred to as the hydrogen-bond energy, but this is not strictly correct since other bonds change slightly when hydrogen bonds are broken In general, H-bonds of energy < 25 kJmol-1 are classified as weak; those in the range 25-40 kJmol-1 are medium; and those in the range > 40 kJmol-1 are strong. Hydrogen bonding may be symmetrical or unsymmetrical. In unsymmetrical hydrogen bonding, the H atom is not midway between the two nuclei, even when the heavier linked atoms are identical. For example, the [ClHCl] ion is linear but the H atom is not midway between the Cl atoms (Fig. 10.9).  the H atom is symmetrically positioned, e.g. in [HF2]-or [H5O2]+ in the bifluoride ion, [FHF], the H atom lies midway between the F atoms; the FꟷF separation (226 pm) is significantly less than twice the van der Waals radius of the F atom (2 × 135 pm). In the formation of [HF2]- appreciable stretching of the original covalent H-F bond takes place, to give two equivalent H……F interactions. HF + F − → [HF2 ]− The bonding in symmetrical X….H….X interactions is best considered in terms of a 3c-2e interaction, i.e. as a delocalized interaction such as was described for B2H6 Each H….F bond is relatively strong, with the bond dissociation enthalpy being of a similar magnitude to that of the F-F bond in F2 (158 kJ mol-1).  Intermolecular and Intramolecular Hydrogen Bonding Intermolecular H-bonding only occurs if the two molecules are in the vicinity (Closer to each other). The ortho isomer will have intramolecular H-bond whereas the para isomer will have intermolecular H-bond.  Effect of hydrogen bonding The hydrogen bonds are usually much weaker than conventional bonds The collective action of H-bonding is responsible for stabilizing complex structures such as the open network structure of ice The H-boning has important effects on the solid state structures of many compounds  Ice possesses an infinite lattice The hydrogen-bonded network may be described in terms of a wurtzite lattice in which each O atom in a tetrahedral environment with respect to other O atoms. Each O atom is involved in four hydrogen bonds, through the use of two lone pairs and two H atoms. The hydrogen bonds are asymmetrical (OH distances = 101pm and 175 pm) and non-linear; each H atom lies slightly off the O····O line, so that the intramolecular HꟷOꟷH bond angle is 105°. The strength of a hydrogen bond in ice or water is 25 kJ mol-1. When water acts as a solvent, hydrogen bonds between water molecules are destroyed as water–solute interactions form; 92A.F. WELLS, Water and hydrates, Chap. 15 in Structural Inorganic Chemisty, 5th edn., pp. 653-98, Oxford University Press, Oxford, 1984.  Density of water ice ice has a relatively low density (0.92 g cm3). On melting (273 K), the lattice partly collapses, allowing some of the lattice cavities to be occupied by H2O molecules. Consequently, the density increases, reaching a maximum at 277 K; between 277 and 373 K, thermal expansion is the dominant effect, causing the density to decrease (Fig. 6.2). Even at the boiling point (373 K), much of the hydrogen bonding remains and is responsible for water having high values of the enthalpy and entropy of vaporization. within the bulk liquid, intermolecular bonds are continually being formed and broken (thus transferring a proton between species) and the lifetime of a given H2O molecule is only ~10-12 s. When water acts as a solvent, hydrogen bonds between water molecules are destroyed as water–solute interactions form;  Intermolecular hydrogen bonding in solid sate: Carboxylic acids form dimeric units are carboxylic acids and amides Solvents of crystallization may be involved in the packing motifs, e.g. H2O or MeOH: (Hydrates?) Hydrogen bonding between difunctional carboxylic acids can result in the formation of chains: If three or more carboxylic acid functionalities are present in a molecule, 2-dimensional sheets or 3- dimensional networks may assemble depending upon the spatial arrangement of the COOH groups. Consider the following pyridine ligands: If the ligand coordinates to a metal ion through the N-donor, the peripheral carboxylic acid units in the metal complex can associate with one another through hydrogen bonding. An example is trans-[PdCl2L2] (L=pyridine-3-carboxylic acid) which forms infinite chains:  Structure of B(OH)3 The three oxygen atoms form a trigonal planar geometry around the boron. Collective hydrogen bonding interactions play a large part in maintaining the structure of protein molecules. They are also responsible for the recognition between specific DNA bases, adenine/thymine and guanine/cytosine, that underlies gene replication.  Water of crystallization, aquo complexes and solid hydrates Many salts crystallize from aqueous solution not as the anhydrous compound but as a well-defined hydrate. Still other solid phases have variable quantities of water associated with them, and there is an almost continuous gradation in the degree of association or “bonding” between the molecules of water and the other components of the crystal. It is convenient to recognise five limiting types of interaction though the boundaries between them are vague and undefined and many compounds incorporate more than one type. (a) H20 coordinated in a cationic complex. The most familiar class can be exemplified by complexes such as [Be(OH2)4]S04, [Mg(OH2)6]Cl2, [Ni(OH2)6](NO3)2, etc.; the metal ion is frequently in the +2 or +3 oxidation state and tends to be small and with high coordination power. Sometimes there is further interaction via H bonding between the aquocation and the anion, particularly if this derives from an oxoacid, e.g. the alums {[M(OH2)6]+[Al(OH2)6]3+[S04]22- related salts of Cr3+, Fe3+, etc. (b) H20 coordinated by H bonding to oxounions. This mode is relatively uncommon but occurs in the classic case of CuS04·5H20 and probably also in ZnS04·7H20. Thus, in hydrated copper sulfate, 1 of the H20 molecules is held much more tenaciously than the other 4 (which can all be removed over P4O10 or by warming under reduced pressure); the fifth can only be removed by heating the compound above 350°C (or to 250” in vacuo). The crystal structure shows that each Cu atom is coordinated by 4H20 and 2S04 groups in a trans octahedral configuration (Fig. 14.10) and that the fifth H20 molecule is not bound to Cu but forms H (donor) bonds to 2 SO4 groups on neighbouring Cu atoms and 2 further H (acceptor) bonds with cis-H20 molecules on 1 of the Cu atoms. (c) Lattice water. Sometimes hydration of either the cation or the anion is required to improve the size compatibility of the units comprising the lattice, and sometimes voids in the lattice so formed can be filled by additional molecules of water. Although LiF and NaF are anhydrous, the larger alkali metal fluorides can form definite hydrates 1 2 1 MF·nH20 (n = 2 and 4 for K; 1 for Rb; and 1 for Cs). 1 3 2 (d) Zeolitic water. The large cavities of the framework silicates can readily accommodate water molecules, and the lack of specific strong interactions enables the “degree of hydration” to vary continuously over very wide ranges. The swelling of ion-exchange resins and clay minerals are further examples of non-specific hydrates of variable composition. (e) Clathrate hydrates. The structure motif of zeolite “hosts” accommodating “guest” molecules of water can be inverted in an intriguing way: just as the various forms of ice are formally related to those of silica so (H20), can be induced to generate various cage-like structures with large cavities thereby enabling the water structure itself to act as host to various guest molecules. Thus polyhedral frameworks, sometimes with cavities of more than one size, can be generated from unit cells containing 12H20, 46H20, 136H20, etc.  Clathrates A clathrate is a host–guest compound, a molecular assembly in which the guest molecules occupy cavities in the lattice of the host species. Molecules of gas (G) are physically trapped in these cavities, there being only weak van der Waals interactions between “guest” and “host” molecules. The clathrates are therefore nonstoichiometric but have an “ideal” or “limiting” composition of [G{C6H4(OH)2}3 Similar clathrates are obtained with numerous other gases of comparable size, such as 02, N2, CO and SO2 (the first clathrate to be fully characterized, by H. M. Powell in 1947) but not He or Ne, which are too small or insufficiently polarizable to be retained. Clathrates provide a means of storing noble gases and of handling the various radioactive isotopes of Kr and Xe which are produced in nuclear reactors.  Noble gas Clathrates The most familiar of all clathrates are those formed by Ar, Kr and Xe with quinol, 1,4-C6H4(OH)2, and with H2O. The noble gas atoms are guests within hydrogen-bonded host lattices. Quinol clathrates are obtained by crystallizing quinol from aqueous or other convenient solution in the presence of the noble gas at a pressure of 10-40 atm. The quinol crystallizes in the less common -form, the lattice of which is held together by hydrogen bonds in such a way as to produce cavities in the ratio 1 cavity: 3 molecules of quinol. Noble gas hydrates are formed similarly when water is frozen under a high pressure of gas. They have the ideal composition, [G(H20)46], and again are formed by Ar, Kr and Xe but not by He or Ne. (limiting composition Ar·6H2O, Kr·6H2O and Xe·6H2O). Other noble gas containing clathrates include 3:5Xe·8CCl4·136D2O and 0:866Xe·3[1,4-(OH)2C6H4] (Figure 17.1). Although this type of system is well established, it must be stressed that no chemical change has occurred  Halogen Clathrates Dichlorine, dibromine and diiodine are sparingly soluble in water. By freezing aqueous solutions of Cl2 and Br2, solid hydrates of approximate composition X2·8H2O may be obtained. These crystalline solids (known as clathrates) consist of hydrogen-bonded structures with X2 molecules occupying cavities in the lattice. An example is 1,3,5-(HOOC)3C6H3 ·0.16Br2; the hydrogen-bonded lattice of pure 1,3,5-(HOOC)3C6H3 was described in  donor name Crown ethers Crowns: The crown ethers are cyclic ethers. O Crown ethers N (NH, NR) Aza crown ethers Crowns are generally described as macrocyclic compounds with hetero S Thio crown ethers atoms such as O, N, S, P or Se as the donor atoms in their ring structures and having property of incorporating cations into their cavities. O, N, S Aza thio crown ethers Only N Aza crowns Sometimes termed as “multidentate macrocyclic compounds” or Only S Thio crowns “macroheterocyclics” nomenclature gives the total number (C + O) and number of O atoms in the ring 12-Crown-4 15-Crown-5 18-Crown-6  Characteristics of crown ethers Crown ethers bound to cationic portion of alkali, alkali earth metal salts, ammonia salts, ionic or other compounds such a thiourea, semicarbazide, diazonium salts; crown ether = host metal ions = guest The complex of crown ether with a guest is formed by ion-dipole interaction between negatively charged donor atoms in the ring structure of cyclic polyethers and cations. Selectivity of crown ethers for a given cations is dependent on following; (a) Relative size of cavity of the crown ring and diameter of the cation (b) Number of donor atoms in the crown ring and the topological effect (c) The relationship between the “hardness” of the cation and that of the donor atom (d) Charge number of cation  Characteristics of crown ethers Different crown ethers have different cavity sizes, The cavity size is not a fixed property because of the ability of the ligand to change conformation. Thus, the radii of the holes in 18-crown-6, 15-crown-5 and 12-crown-4 can be taken to be roughly 140, 90 and 60 pm respectively. e.g., 15-Crown-5 can capture Na+, 18-Crown-6 can capture K+, 21-Crown-7 can capture Cs+. Property Li Na K Rb Cs Metallic radius, rmetal/pm 152 186 227 248 265 Ionic radius, rmetal/pm 76 102 138 149 170 Donald J Crams, Jean-Marie Lehn and Charles J. Pedersen shared Nobel Prize in 1987 for their development and use of molecules with structure specific interactions of high selectivity  Selectivity of crown ethers The radius of the cavity inside the 18-crown-6 ring is 140 pm, and this compares with values of rion for the alkali metal ions ranging from 76 pm for Li+ to 170 pm for Cs+. The radius of the K+ ion (138 pm) is well matched to that of the macrocycle, The stability constants for the formation of [M(18-crown-6)]+ in acetone follow the sequence K+ > Rb+ > Cs+ Na+ > Li+. Property Li Na K Rb Cs Metallic radius, rmetal/pm 152 186 227 248 265 Ionic radius, rmetal/pm 76 102 138 149 170 It has been pointed out that the stability constants for [KL]+ complexes are often higher than for corresponding [ML]+ complexes where M = Li, Na, Rb or Cs, even when hole-matching is clearly not the all- important factor. An alternative explanation focuses on the fact that, when a crown ether binds M+, the chelate rings that are formed are all 5-membered, and that the size of the K+ ion is ideally suited to 5-membered chelate ring formation. A complex may form in which M+ sits above the plane containing the donor atoms, e.g. [Li(12-crown-4)Cl] Alternatively a 1 :2 complex [ML2]+ may result in which the metal ion is sandwiched between two ligands, e.g. [Li(12-crown-4)2]+. Crown ethers have been used to stabilize [H OH2 n ]+ ions, the stabilizing factor being the formation of hydrogen bonds between the O atoms of the macrocyclic ligand and the H atoms of the [H OH2 n ]+ ion. Hydrogen bonding between the [H5 O2 ]+ and [H7 O3 ]+ ions and crown ethers is shown by hashed lines; hydrogen atoms in the crown ethers are omitted for clarity. the association of a chain structure involving alternating crown ether and [𝐇𝟕 𝐎𝟑 ]+ ions. the bond lengths determined by neutron diffraction show two asymmetrical Colour code: C, grey; O, red; H, white. hydrogen bonds and this is consistent with [𝐇𝟕 𝐎𝟑 ]+ being considered in terms of [𝐇𝟑 𝐎]+ · 𝟐𝐇𝟐 𝐎. the encapsulation of an [𝐇𝟓 𝐎𝟐 ]+ ion within a single crown ether Figure: The stabilization in the solid state of [H5 O2 ]+ and [H7 O3 ]+ by hydrogen bonding to crown ethers: (a) the structure of dibenzo-24-crown-8; (b) the structure of [(H5O2)(dibenzo-24-crown-8)]+ determined for the [AuCl4] salt by X-ray diffraction (c) the structure of 15-crown-5; and (d) part of the chain structure of [(H7O3)(15-crown-5)]+ determined for the [AuCl4] salt by neutron diffraction. Synthesis in 1967, Charles J Pedersen, USA, Du pond  Application of crown ethers 1. Crystalline complexes 2. Solubilize the inorganic slats in organic solvent 3. Phase transfer catalysis 4. Ion carrier inside the living cell to transport ions across cell membranes and thus maintain the balance between Na+ & K+ inside and outside cells whereas KMnO4 is water-soluble but insoluble in benzene, [K(18-crown-6)][MnO4] is soluble in benzene; mixing benzene with aqueous KMnO4 leads to the purple colour being transferred from the aqueous to the benzene layer. This phenomenon is very useful in preparative organic chemistry, the anions being little solvated and, therefore, highly reactive. F- ion should be strong base and nucleophile. However, it is not, because ions are paired in solution and not free to react. 18-Crown-6/KF/benzene, solubility increases 10 times, increases nucleophilicity. 1 PS + Et3N·3HF 144 h Yield = 63% 2 Bu4P+·HF2- 2h Yield = 89% 3 KF(18-Crown-6) 1 h Yield = quantitative Another example of this fit between cation and ligand cavity is thought to be responsible for the transport of Na and K ions across cell membranes. The ions cross the hydrophobic cell membrane by means of cavities lined with donor atoms. The donor atoms are arranged to form a cavity, the size of which determines whether Na or K is bound. Such ion channels modulate the Na/K concentration differential across the cell membrane that is essential for particular functions of the cell. Valinomycin is a cyclic polypeptide. Valinomycin is present in certain microorganisms. The naturally occurring molecule valinomycin is an antibiotic. Selectively coordinates K: the resulting hydrophobic 1:1 complex transports K through a bacterial cell membrane, depolarizing the ion differential and resulting in cell death. Figure illustrates that the valinomycin ligand uses six of its carbonyl groups to octahedrally coordinate K+. The [K(valinomycin)]+ ion has a hydrophobic exterior which makes it lipid-soluble, and the complex ion can therefore be transported across the lipid bilayer of a cell membrane.  Cryptands A cryptand is a polycyclic ligand containing a cavity; when the ligand coordinates to a metal ion, the complex ion is called a cryptate. These molecules are three dimensional analogues to crown ethers. A cryptand is a polycyclic ligand containing a cavity. When the ligand coordinates to a metal ion, the complex ion is called a cryptate. Figure shows the structure of the cryptand ligand 4,7,13,16,21,24-hexaoxa-1,10-diazabicyclo[8.8.8]hexacosane, commonly called cryptand-222 or crypt-222, where the 222 notation gives the number of O-donor atoms in each of the three chains. Cryptand-222 is an example of a bicyclic ligand which can encapsulate an alkali metal ion. Cryptands protect the complexed metal cation even more effectively than do crown ethers. They show selective coordination behaviour; cryptands-211, -221 and -222 with cavity radii of 80, 110 and 140 pm, respectively, form their most stable alkali metal complexes with Liþ+, Na+ and K+ respectively  Cryptands (1) Cryptands protect the complexed metal cation even more effectively than do crown ethers. (2) Better selectivity and strength of binding than crown ether (3) Bind to salts so as to solubilize in organic solvents (4) Phase transfer catalysts (5) Synthesis of alkalide and electride (6) Used in the crystallization of Zintle ions such as Sn92-  Complexes of Cryptands Cooling a solution of Na in ethylamine with cryptand 2.2.2 will give [Na(crypt−222)]+ Na−. The isolated product is the diamagnetic, golden-yellow and stable below -10°C. Formed by the disproportionation of Na into Na+ and Na−. Electron transfer occurred between two sodium atoms forming Na+ and Na−. The solid state structure indicates that the effective radius of the sodide ion is 230 pm, i.e. Na− is similar in size to I−. The large crypt ligand completely shield Na− ion and prevents recombining with Na+ The replacement of the O atoms in crypt-222 by NMe groups generates ligand 10.2, ideally suited to encapsulate K+. Its use in place of crypt-222 has aided the study of alkalide complexes by increasing their thermal stability. [K(10.2]+ Na− and [K(10.2]+ K − are stable at 298 K. The first hydrogen sodide 𝐇+ 𝐍𝐚− was prepared using ligand 10.4 to encapsulate H+, Alkalides have also been prepared containing Rb− and Cs −. In these reactions, the cryptand:metal molar ratio is 1: 2. If the reaction is carried out using a greater proportion of ligand, paramagnetic black electrides can be isolated, in which the electron is trapped in a cavity of radius 240 pm.