CHM 213 Inorganic Chemistry I Lecture Notes PDF

CHM 213/STC 213: Inorganic Chemistry I LECTURER’S NAME: Mr. D. S. Olayanju Learning Outcome By the end of this course, the students should be able to; 1. list the first-row transition elements; 2. give the properties of transition elements; 3. explain basic terms used in coordination chemistry;...

CHM 213/STC 213: Inorganic Chemistry I LECTURER’S NAME: Mr. D. S. Olayanju Learning Outcome By the end of this course, the students should be able to; 1. list the first-row transition elements; 2. give the properties of transition elements; 3. explain basic terms used in coordination chemistry; 4. give the IUPAC name of given metal complexes 5. explain crystal field theory (CFT) 6. draw the diagram to illustrate with examples of coordination compounds; 7. state the advantages of CFT over other bonding theories; 8. discuss the comparative Chemistry of the following elements; (I) Ga, In, Tl; (II). Ge, Sn, Pb; (III). As, Sb, Bi and (IV). Se, Te, Po; Course Contents Chemistry of first row transition metals. Properties of transition elements. Basic terms in coordination chemistry. IUPAC nomenclature of Coordination compounds. Introduction to coordination chemistry including elementary treatment of crystal field theory. Comparative Chemistry of the following elements: (a) Ga, In, TI, (b) Ge, Sn, Pb, (c) As, Sb, Bi (d) Se, Te, Po. FIRST ROW TRANSITION ELEMENTS A transition element is one for which an atom has an incomplete d-subshell, or which gives rise to a cation with an incomplete d-subshell. These elements are found in the d-block of the periodic table between Groups 2 and 3. The first transition series is from scandium to zinc. This series of ten d – block metals occur in period 4. Potassium and calcium are the s- block metal in the period. They form a transition in the properties between the highly electropositive s- block metals and the electronegative elements of the p – block. The electronic configuration of the first transition series After calcium, (electronic configuration: 1S2 2S22P63S23P64S2), there is a change in the filling of the subsidiary energy level and the inner 3d shell (which contain a maximum of ten electrons) is filled with electrons before beginning to fill the 4S subshell. This means that the 4s orbital is filled in preference to the five 3d orbitals. This indicates that the 4s orbital has lower energy than the 3d orbital. The electronic configuration of chromium and copper have one electron each in their 4s orbital. This is because half – filled or completely filled sub -shells are more stable than partially filled sub – shells Properties of Transition Elements Physical properties of transition elements Transition elements are metals and hence, show the following physical properties; 1. They hard metals compared to the s – block metals. 2. They have high tensile strength. 3. They are ductile. 4. Transition elements have high melting and boiling points. 5. They have strong metallic bonding. 6. They are good conductors of heat and electricity. Chemical properties Chemical Properties 1. Chemical Reactivity: There is usually a variation in the chemical properties of the elements in the same period in s and p- blocks from left to right. However, this does not happen with the transition metals because electrons are added progressively to the inner d – orbitals, not the outermost orbitals, as in the s – block and p - block elements. The nuclei of the transition metals exerts a greater attraction on their electrons than the nuclei of s – block metals. So, the s -block elements have lower ionization energy and are more reactive than the transition metals. Generally, transition metals are moderately reactive. The reactivity decreases across the period due to corresponding increase in value of ionization energies. 2. Variable oxidation states: transition metals have variable oxidation states because the electrons in 3d and 4s orbitals are available for bond formation. For instance, for manganese; Mn – 1s22s22p63s23p63d54s2 can lose 1, 2, 3, 4, 5, 6 and 7 electron(s) to attain oxidation states of +1, +2, +3, +4, +5, +6 and +7. Table below shows the various oxidation states of first row transition metals with the most stable oxidation states painted blue while oxidation state in square bracket is rare. In addition, the differences in the successive ionization energies of transition metals are small. This allows the formation of stable ions at each stage. 3. Complex ion formation: a complex ion has a central positive ion linked to several other atoms, ions or molecules called ligands. The bonding between the metal ion and the ligands could be electrovalent or coordinate covalent. An example of complex ion is blue tetraaminecopper(II) ion, [Cu(NH3)4]2+, in which the central copper (II) ion is linked by coordinate bonding to four ammonia molecules. The electrons donated by the ligands usually fill up the incomplete d orbitals, and the 4s and 4p orbitals in the central transition metal ion. All transition metals ions tend to form complex ions with water. These hydrated ions are usually unstable. 4. Formation of coloured compounds: transition metallic ions are usually coloured which is a guide in identifying a compound. The colours are as a result of the partially filled 3d orbitals (i.e. 3d1 to 3d9) zinc and scandium ions colourless since they do not have partially filled 3d orbitals. In the complex ions, the nature of the ligands also contributes to the colour. 5. Catalytic Activities: transition elements and their compounds are very effective as catalysts. The ease with which the ions of transition elements change their oxidation state and the partially filled 3d orbitals which allows the exchange of electrons to and from molecules enable them act as catalyst. 6. Paramagnetism: this is the property of substance to attract to a magnet. The occurrence of paramagnetic compounds of d-block metals is common and arises from the presence of unpaired electrons. This phenomenon can be investigated using electron spin resonance (ESR) spectroscopy. It also leads to signal broadening and anomalous chemical shift values in NMR spectra. Coordination Chemistry Ligand: This is a molecule or ion carrying suitable donor groups capable of binding (or coordinating covalently) to a central atom. The range of molecules that can bind to metal ions as ligands includes inorganic atoms, ions and molecules as well as organic molecules and ions. Most ligands are neutral or anionic substances but cationic ones, such as the tropylium cation, are also known. Some ligands can acts as both donor and acceptor of electrons. e.g. CO can donate electrons to the central metal ion and can as well receive electron from it. Types of Ligands 1. Based on the nature of the bond between the ligand and the central atom; a. Anionic Ligands: CN-, Br-, Cl-, SCN-, C5H5- etc. b. Cationic Ligands: NO+, N2H5+ etc. c. Neutral Ligands: CO, H2O, NH3 etc. 2. Based on the denticity or number of binding sites: Denticity of Ligands is the number of Co-ordinating or donor groups present in a single ligand that can coordinate to the central metal atom or ion. Based on denticity, ligands are classified into; a. Monodentate ligands: ligands that bind to the central atom through one atom only; Common monodentate ligands are; Common Name IUPAC Name Formula Fluoro Fluoro F- Chloro Chloro Cl- Bromo Bromo Br- Iodo Iodo I- Azido Azido N3- Cyano Cyano CN- Thiocyano Thiocyanato-S (S-bonded) SCN- Isothiocyano Thiocanato-N (N-bonded) NCS- hydroxo hydroxo OH- Aqua Aqua H2O Carbonyl Carbonyl CO Thiocarbonyl Thiocarbonyl CS Nitrosyl Nitrosyl NO+ Nitro Nitrito-O (O-bonded) ONO- Methylisocyanide Methylisocyanide CH3NC Pyridine Pyridine Py phosphine phosphane PR3 ammine Ammine NH3 methylamine methylamine MeNH2 amido amido NH2- b. Bidentate Ligands: Also called the didentate ligands are molecules or ions that bind to the central atom via two co-ordinate covalent bonds i.e. Ligands having two binding sites. Examples include; Name of ligand symbol Formula/Structure ethylenediamine en NH2CH2CH2NH2 oxalato ox C2O42- Dimethylgly oximato DMG HONCC(CH3)C(CH3)NOH (butanediene dioxime) c. Tridentate and Polydentate Ligands: ligands with three lone pairs of electrons to donate to the central metal atom or ion. The one with four donaor atoms/binding site is tetradentate, with five, six etc  referred to as pentadentate, hexadentate etc, respectively. Generally, ligands with more than two binding sites/donor atoms are refers to as polydentate ligands. Examples; Types Common name IUPAC name Abbrev Structure of iation ligand tridentat diethylenetriami 2,2’- dien NH2CH2NHCH2CH2NH2 e ne diaminediethylamine tetraden triethylenetetraa 1,4,7,10- trien NH2(CH2)2NH(CH2)2NH(CH2)2NH2 tate mine tetraazadecane pentade tetraethylenepen 1,4,7,10,13- NH2(CH2)2NH(CH2)2NH(CH2)2NH(C ntate tamine pentaazatridecane H2)2NH2 hexaden Ethylenediamine 1,2- EDTA (-OOCCH2)2N(CH2)2N(CH2COO-)2 tate tetraacetato ethenediyl(dinitrito)te traacetate d. Ambidentate ligands: these are ligands that can bind with the central atom or ion via two different donor atoms in two separate ways. Examples include; Name Structure Nitrito-N (nitro) M←NO2- Nitrito-O M←O-N=O Cyano M←C≡N isocyano M←N≡C Thiocyanato M←S-C≡N isothiocyanato M←N=C=S e. Bridging Ligands: These are ligand that connect two or more central atoms or ions are known as bridging ligands. µ, is used as prefix to their naming. Examples include; 1. CO in Fe2(CO)9 2. Hydride (H-) in B2H6 Chelating ligands and chelate: Chelating ligands are ligands with two or more points of attachment to central metal atoms or ions. The compound formed by a chelating ligand with the central metal atom or ion is known as a chelate. The primary difference between a chelating ligand and a polydentate ligand is that a chelating ligand bind more than once to the same central or ion, but a normal polydentate ligand does not necessarily ligate to the same metal. Chelate Effect: this is defined as the amplified affinity of a chelating ligand for a central metal ion compared to its monodentate non-chelating ligands e.g. a complex containing chelate rings e.g. [Co(en)3]3+ is more stable than [Co(NH3)6]3+ where en is a chelating ligand. Coordination Number: The number of atoms, ions bonded to a central metal is the coordination number. It could also be defined as the total number of sigma bonds through which the ligands are bound to the coordination centre. It is usually determined by; i. size of the central metal ii. the number of d-electrons iii. steric effects arising from the ligands. Complexes with co-ordination number of 2 to 9 are known but 4 to 6 coordination complexes are the most stable electronically and geometrically and the most numerous. Table below shows complexes with respective coordination numbers, the possible structure and hybridization. Complexes Coordination number Possible hybridisation Structure/geometry [Au(NH3)2]+ or 2 sp Linear [AgCl2]- [HgI3]- 3 sp2 Trigonal planar [FeBr4]2- 4 sp3 Tetrahedral [Ni(CN)4]2- 4 sp2d Square planar [CuCl5]3- 5 sp3d Trigonal bipyramidal [Ni(CN)5]3- 5 sp3d Square-based pyramidal [Co(NH3)6]3+ 6 sp3d2 Octahedral [ZrMe6]2- 6 sd5 or sp3d2 Trigonal prismatic [V(CN)7]4- 7 sp3d3 Pentagonal bipyramidal [NbF7]2- 7 sp3d3 Monocapped trigonal prismatic [P9F8]3- 8 sp3d3f Cubic [Mo(CN)8]4- 8 sp3d4 dodecahedral [TaF8]3- 8 sp3d4 Square antiprismatic [ReH9]2- 9 sp3d5 Tricapped trigonal prismatic Metal Complexes/Coordination Compounds: These are chemical compounds consisting of a central metal atom or ion by co-ordinate covalent bonds to one or more ligands, which are ions or molecules that contain one or more pairs of electrons that can be shared with the metal. The central metal in the metal complexes is usually transition element and is called the coordination centre. Metal complexes, if charged, are called complex ions. Metal complexes/coordination complexes could be classified as; i. Cationic complexes: In this the coordination sphere is a cation e.g. [Co(NH3)6]Cl3 ii. Anionic complexes: the coordination sphere is anion e.g. K4[Fe(OH)6] iii. Neutral complexes: the coordination sphere is neutral i.e. neither cationic nor anionic e.g. [Ni(CO)4] iv. Homoleptic complexes: consisting of similar type of ligands e.g. K4[Fe(CN)6] v. Heteroleptic complexes: these consists of different types of ligands bonded to the central metal atoms e.g. [Co(NH3)5Cl]SO4 vi. Mononuclear Complexes: the co-ordination sphere has single transition metal ion e.g. K4[Fe(CN)6] vii. Polynuclear complexes: consist of more than one metal ion e.g. [ReCl4][ReCl4]  Figure 1: Example of a metal complex Co-ordination Sphere: this is the non-ionizable part of a complex compound which consists of central transition metal ion surrounded by neighbouring atoms or groups enclosed in square bracket. The coordination sphere comprises of the coordination Centre, the ligands attached to the coordination centre, and the net-charge of the chemical compound as a whole.it is usually accompanied by a counter ion (the ionizable groups that attach to charged coordination complexes) e.g. [Co(NH3)6]3+.3Cl- IUPAC NOMENCLATURE OF COORDINATION COMPOUNDS IUPAC rules on naming coordination compounds include; 1. Positive ions are named first in complexes before the negative ions, e.g K3[Fe(CN)6] - Potassium hexacyanoferrate (III). 2. The ligands are named first in the coordination sphere before the metal, e.g: [Cu(NH3)4] SO4 : Tetraammine copper (II) sulphate. 3. a. The number of ligands is indicated by the following prefixes: 2- di, 3- tri, 4- tetra, 5- penta, 6- hexa, 7- hepta, 8- octa, 9- nona, 10- deca, etc. Example: [Co(NH3)6] Cl3 – hexaammine cobalt(III) chloride b. If the name of the ligand has already include a prefix, e.g ethylenediammine, (prefix, di, already included), or the name is complicated, the name is set of in parenthesis and the second set of prefixes indicating number of ligand ends in – is/kis, e.g bis- 2, tris- 3, tetrakis- 4, pentakis- 5, hexakis- 6, heptakis- 7, octakis- 8, nonakis- 9, decakis- 10 etc. e.g. [Co(H2NCH2CH2NH2)2Cl2]+ - Dichlorobis(ethylenediammine)cobalt(III) ion. [Fe(NH4C5-C5H4N)3]2+ - Tris(bipyridine) Iron(II)ion. 4. In a case of more than one ligand in coordination sphere, the ligands should name alphabetically regardless of the prefixes, e.g [Co(NH3)4Cl2]+ - Tetraammine dichloro cobalt (III) ion; [Pt(NH3)Br(CH3NH2)Cl] - Amminebromochloromethylamine platinum (II). 5. Anionic ligands are end with suffix, - o - ,e.g chloro, bromo, iodo, etc. coordinated water molecule is named, aqua, coordinated ammonia molecule, ammine (double, m) in case of uncoordinated ammonia, single , m is used, e.g methylamine. 6. If the coordination sphere is having overall negative charge, the name of the metal in the sphere is modified to end in –ate, e.g  [PtCl6]2- - hexachloroplatinate (IV). 7. Oxidation state of metals are indicated by Roman numeral in parenthesis according to Stock system. 8. Oxidation state of metals are indicated by putting overall charges on the coordination sphere in parenthesis after the name of the metal, [Pt(NH3)4]2+ - tetraammine platinum(+2), this is based on Ewing-Bassett system. 9. To indicate geometrical adjacent and opposite locations, the prefixes; cis and trans are applied, e.g cis- diamminedichloroplatinum(II), cisplatin 10. Bridging ligands between two metals ions have the prefix- μ. E.g μ.-amido- μ.-hydroxobis(tetraamminecobaltIV). 11. Naming of some metals takes the origin with modification and ends with the suffix ---ate , provided the complex is negatively charged ,e.g, Gold(Au)- aurate, Silver(Ag)-argentate, Iron(Fe)- Ferrate, etc. e.g. [Au(CN)2]- - Dicyanoaurate(I) ion [FeCl4]- - Tetrachloroferrate(III) ion. CRYSTAL FIELD THEORY Crystal field theory is an electrostatic model which predicts that the d orbitals in a metal complex are not degenerate. The pattern of splitting of the d orbitals depends on the crystal field, this being determined by the arrangement and type of ligands. It is one of the approaches to the explanation of bonding in complexes of the d-block metals. Ligands are considered as point charges and there are no metal–ligand covalent interactions. The octahedral crystal field Octahedral complexes such as given by Figure 2a below has six ligands placed on the Cartesian axes at the vertices of an octahedron. Each ligand is treated as a negative point charge and there is an electrostatic attraction between the metal ion and ligands. In a similar manner, repulsive interaction between electrons in the d orbitals and the ligand point charges also exists. If the electrostatic field (the crystal field) were spherical, then the energies of the five 3d orbitals would be raised (destabilized) by the same amount. However, since the dz2 and dx2-y2 atomic orbitals point directly at the ligands while the dxy, dyz and dxz atomic orbitals point between them, the dz2 and dx2-y2 atomic orbitals are destabilized to a greater extent than the dxy, dyz and dxz atomic orbitals (Figure 3.0). Figure 2: a. octahedral complex structure, b. Shapes of d-orbital orbitals The tetrahedral crystal field Figure 3: Octahedral Crystal Field Splitting Thus, with respect to their energy in a spherical field (the barycentre, a kind of ‘centre of gravity’), the dz2 and dx2-y2 atomic orbitals are destabilized while the dxy, dyz and dxz atomic orbitals are stabilized. The dz2 and dx2-y2 orbitals then have eg symmetry, while the dxy, dyz and dxz orbitals possess t2g symmetry. The energy separation between them is Δoct or 10Dq. The overall stabilization of the t2g orbitals equals the overall destabilization of the eg set. Thus, orbitals in the eg set are raised by 0.6Δoct with respect to the barycentre while those in the t2g set are lowered by 0.4 Δoct (Figure 4.0) also shows these energy differences in terms of 10Dq. Figure 4: Splitting of the d orbitals in an octahedral crystal field, with the energy changes measured with respect to the barycentre. Tetrahedral Complexes These are complexes having four ligands on the apices of a tetrahedron around the central metal. [CoX 4]2- (X = Cl,Br, I), Ni(CO)4, etc. are all examples of tetrahedral complexes. When a metal is placed on the origin of the Cartesian axes for a tetrahedral complexes, as in the octahedral complexes, e orbitals (dx2-y2, dz2) are distant from the ligands while t2 orbitals (dxy, dyz, dxz) become nearer to the ligands. As a result, the electronic repulsion on t2 orbitals becomes larger than that on e orbitals and the t2 orbitals are relatively destabilized compared to those of the e orbitals. The five-fold degenerate orbitals of the central metal is then split into two-fold degenerate e and three-fold degenerate t2 sets (Figure 5). Hence, for a regular tetrahedron, the splitting of the d orbitals is inverted compared with that for a regular octahedral structure, and the energy difference (∆t) is smaller than that of octahedral. The relative splitting ∆oct and ∆t are related by; 4 ∆t = 9 ∆oct ≈ 1/2∆oct Since ∆t is significantly smaller than ∆oct, tetrahedral complexes are high-spin. Also, smaller amounts of energy are needed for t2←e transition (tetrahedral) than for eg←t2g transition in octahedral complexes. This account for the differences encountered in the colour of the corresponding octahedral and tetrahedral complexes.  Figure 5: Crystal Field Splitting in Tetrahedral Complexes Jahn - Teller distortion A regular octahedral environment is the most stable one for a spherically symmetrical metal ion surrounded by six donor atoms. For metal ions with certain d electron configurations which are not spherically symmetric, the regular octahedral configuration is not the most stable. This situation is expressed in Jahn – Teller theorem: Any non-linear molecule that is in an electronically degenerate state will undergo distortion to lower the symmetry, remove the degeneracy, and lower the energy. Jahn–Teller effects are seen in d9 tetrahedral complexes such as [CuCl4]2-, and high-spin d4 complexes such as [FeO4]4-. Much weaker distortion also occurs if the t2g level is not uniformly occupied. In the case of eg, any odd electron can occupy either the dz2 or the dx2-y2 orbital. If, however, the complex undergoes distortion the eg level is split and the electron can occupy the lower of the two orbitals (the dz2 orbital in the case of tetragonal elongation, or the dx2-y2 orbital in the case of tetragonal compression). In tetragonal elongation the ligands on the z axis move out and therefore interact less with those orbitals which have a z component; i.e. the dz2, dxz, and dyz, and these orbitals attain lower energy (i.e. stabilized). Those orbitals without a z component, i.e. dx2-y2 and dxy, will be raised a corresponding amount.  Elongation: dz2 attains lower energy Figure 6: Jahn-Teller Splitting in a Cu2+ Ion The square planar crystal field A square planar arrangement of ligands can be formally derived from an octahedral array by removal of two trans-ligands. If we remove the ligands lying along the z axis, then the dz2 orbital is greatly stabilized. The energies of the dyz and dxz orbitals are also lowered, although to a smaller extent. The resultant ordering of the metal d orbitals is given in the Figure 7 below. The fact that square planar d8 complexes such as [Ni(CN)4]2- are diamagnetic is a consequence of the relatively large energy difference between the dxy and dx2-y2 orbitals.  Figure 7: Crystal Field Splitting in Square Planar Complex Advantages of Crystal Field Theory over bonding theories Crystal field accounts for the following properties of metal complexes which bonding theories failed to: 1. Splitting of d-orbitals of metal 2. Colour observed in complexes 3. Magnetic Properties of complexes 4. Spectral properties of metal complexes 5. Distorted octahedral and tetrahedral geometry. Limitations of Crystal Field Theory Crystal field theory brings together structures, magnetic properties and electronic properties. With crystal field theory, we can interpret the contrasting magnetic properties of high- and low-spin octahedral complexes on the basis of the positions of weak- and strong-field ligands in the spectrochemical series. However, the theory provides no explanation as to why particular ligands are placed where they are in spectrochemical series. Comparative Chemistry of Ga, In and Tl Occurrence The three elements occur in trace amounts as sulfides in various minerals. Extraction Increase in the demand for gallium arsenide (GaAs) in components for electronic equipment, makes its production be on the increase. The main source of Ga include: i. crude bauxite, in which Ga is associated with Al. ii. Residues from the Zn-processing industry. Indium occurs in the zinc sulfide ore sphalerite (also called zinc blende) where, being a similar size to Zn, it substitutes for some of the Zn. The extraction of zinc from ZnS therefore provides indium as a by-product. Thallium is obtained as a by-product of the smelting of Cu, Zn and Pb ores, although demand for the element is low. Physical properties Electronic configurations These elements has outer electronic configuration of ns2np1 as given below: 31 Ga - [Ar]3d104s24p1 49 In - [Kr]4d105s25p1 81 Tl – [Xe]4f145d106s26p1 Oxidation states Although an oxidation state of +3 (and for Ga and Tl, +1) is characteristic of these elements, Ga could also form compounds in which a formal oxidation state of +2 is suggested, e.g. GaCl2 which is the mixed oxidation state species Ga[GaCl4]. Appearance Gallium is a silver-coloured metal with a particularly long liquid range (303–2477K). Indium and thallium are soft metals, and In has the unusual property of emitting a high-pitched ‘cry’ when the metal is bent. Melting points The melting point of Ga is 303 K, In is 430 K and that of Tl is 576.5 K. The melting are found to increase down the group due to increase in the metallic character as one descends the group. Reactivity 1. Gallium, indium and thallium dissolve in most acids to give salts of Ga(III), In(III) or Tl(I), but only Ga liberates H2 from aqueous alkali. 2. All three metals react with halogens at, or just above, 298K to give MX3 product with the exceptions of the two reactions given below; 2Tl + 2Br2 → Tl[TlBr4] 3Tl + 2I2 → Tl3I4 3. Formation of Hydrides: Digallane, Ga2H6 exist and structurally similar to B2H6. Digallane, Ga2H6, is prepared by the reaction; The product condenses at low temperature as a white solid (mp 223K) but decomposes above 253K. Gallaborane GaBH6, is another hydride of Ga synthesized from the reaction of H2Ga(µ-Cl)2GaH2 with Li[BH4] at 250K in the absence of air and moisture. Structure of Gallaborane A number of adducts of InH3 containing phosphine donors have also been isolated, e.g. These products are stable in the solid state at 298K, but decompose in solution. Comparison between digallane and borane i. Ga2H6 is unlike B2H6 in that Ga2H6 rapidly decomposes to its constituent elements. ii. Ga2H6 and B2H6 both react with HCl, but in the case of the borane, substitution of a terminal H by Cl is observed, whereas both terminal and bridging H atoms can be replaced in Ga2H6. iii. Ga2H6 is like B2H6 in that it reacts with Lewis bases. 4. Halides and complex halides: The trifluorides of Ga, In and Tl are non-volatile solids, best prepared by fluorination of the metal (or one of its simple compounds) with F2. Each trifluoride is high melting and has an infinite lattice structure. Gallium and indium trichlorides and tribromides form adducts of 4, 5 or 6 coordination number such as: [MCl6]3-, [MBr6]3-, [MCl5]2-, [MCl4]- and [MBr4]- (M = Ga or In) and L.GaX3 or L3.InX3 (L = neutral Lewis base). 5. Oxides of Ga, In and Tl: Gallium, like Al, forms more than one polymorph of Ga2O3, GaO(OH) and Ga(OH)3, and the compounds are amphoteric. In2O3, InO(OH) and In(OH)3 are basic. Thallium is unique among the group in exhibiting an oxide for the M(I) state such as 2TlOH formed by reacting Tl 2O with water. Tl2O + H2O → 2TlOH Thallium (III) forms the oxide Tl2O3, but no simple hydroxide. 6. Nitrides: most of their nitrides are more reactive than that of B or Al. 7. Coordination complex formation: Increasing numbers of coordination complexes of the ions of these elements are known with Octahedral coordination been the most common, e.g. in [M(acac)3] (M = Ga, In), [M(ox)3]3- (M = Ga, In) and mer-[Ga(N3)3(py)3]. The complexes [M(acac)3] are structurally related to [Fe(acac)3]. Note: Thallium shows similarities to elements outside those in group 13, and can be compared to the alkali metals, Ag, Hg and Pb. Uses Gallium and indium phosphides, arsenides and antimonides have important applications in the semiconductor industry. They are used as transistor materials and in light-emitting diodes (LEDs) in pocket calculators; the colour of the light emitted depends on the band gap. GaAs, apart from its use in LEDs, it equally found application in laser diodes, photodetectors, solar cells and in integrated circuits, e.g. in high- performance computers. Indium is used in thin-film coatings, e.g. liquid-crystal displays and electroluminescent lamps. Indium is also used in lead-free solders, in semiconductors, for producing seals between glass, ceramics and metals owing to its bond to non-wettable materials. Also used in fabricating special mirrors which reduce headlight glare. Indium–tin oxide (ITO) is used as a coating material for flat-panel computer displays, for coating architectural glass panels, and in electrochromic devices. Coating motor vehicle and aircraft windscreens and motor vehicle rear windows allows them to be electrically heated for de-icing purposes. A thin film of ITO on the cockpit canopy of an aircraft such as the stealth plane renders this part of the plane radar-silent, contributing to the sophisticated design that allows the stealth plane to go undetected by radar. Thallium sulfate used as insecticides and rodenticides though extremely toxic and must be treated with caution. It is used in semiconducting materials, in selenium rectifiers, in Tl-activated NaCl and NaI crystals, in ‫ץ‬-radiation detectors, and in IR radiation detection and transmission equipment. The radioisotope 201Tl (t1/2 = 12.2 d) is used for cardiovascular imaging. Comparative Chemistry of Arsenic (As), Antimony (Sb), and Bismuth (Bi) Arsenic, antimony and bismuth are the heavier pnictogen (Group 15) elements. They adopt the ground state electron configuration ns2np3. Electronic Configuration As - 33: [Ar]3d104s24p3 Sb – 51: [Kr]4d105s25p3 Bi – 83: [Xe]4f145d106s26p3 Physical Properties Arsenic and antimony are considered metalloids, while bismuth is metallic. Arsenic and bismuth are monoisotopic, while antimony has two substantially abundant naturally occurring isotopes. The atomic, covalent, and ionic radii increase from arsenic to bismuth, consistent with their increasing atomic mass and number of electron shells.  General properties of Arsenic (As), Antimony (Sb), and Bismuth (Bi) Oxidation States Arsenic and bismuth exhibit a relatively unstable +5 oxidation state, while antimony tends to form extended molecular structures, which distinguishes its chemistry from that of arsenic and bismuth. The pnictogen elements can access oxidation states ranging from -3 to +5, but they thermodynamically favor the elemental form over positive oxidation states. Allotropes Arsenic primarily exists in the form of β-arsenic (grey arsenic), while antimony and bismuth are most stable in their α forms, which are rhombohedral and typically grey in appearance. The allotropes of these elements reflect their distinct structural properties, with the α-allotropic forms being analogous to black phosphorus, composed of layers of hexagonally connected sheets. Arsenic is observed to exist in two (yellow and black), additional allotropic forms, while antimony adopts five allotropes and bismuth adopts at least three allotropes. Most of these alternate allotropes are only nominally stable or require high temperature or pressure conditions. Structural Characteristics The coordination environments preferred by these elements lead to significant structural features. Arsenic and antimony can adopt extended structures similar to phosphate, while bismuth forms a variety of cationic element clusters. The tendency of antimony and bismuth to form extended molecular structures further distinguishes their chemistry from arsenic. Bond Energies They can engage in p-bonding with neighbouring atoms, although this is thermodynamically disfavored compared to s-bonding. Arsenic, antimony, and bismuth form stable covalent bonds with most elements. Bond energies generally decrease from arsenic to bismuth. For example, the Pn-H bonds in AsH₃ and SbH₃ are 319.2 kJ mol⁻¹ and 288.3 kJ mol⁻¹, respectively. Bonds involving lighter elements are generally stronger, as seen in the Bi-X bonds in BiF₃ and BiBr₃. Relativistic Effects and Orbital contraction Relativistic effects and orbital contraction are significant for these elements. Bismuth experiences relativistic effects, while arsenic shows orbital contraction due to a relatively high effective nuclear charge. Coordination Chemistry The coordination environment preferred by these elements lead to significant structural features. They can adopt various coordination numbers, with a common geometry involving three covalent bonds and one lone pair of electrons. Biological Activity Arsenic and antimony are known for their toxicity or negative bioactivity, whereas bismuth is recognized for providing therapeutic responses or demonstrating positive bioactivity e.g. the use of bismuth compounds such as peptobismol and DeNol in the treatment of many medical disorders. This difference in biological activity highlights the varying implications of these elements in health and environmental contexts. Comparative Chemistry of Selenium, Tellurium, and Polonium Selenium, tellurium, and polonium are Group 16 elements (chalcogens) in the periodic table. They exhibit unique transitions in properties, moving from semi-metallic (selenium, tellurium) to metallic (polonium). Their increasing metallic character impacts their reactivity and applications. Electronic Configuration Selenium (Se): [Ar] 3d¹⁰ 4s² 4p⁴ Tellurium (Te): [Kr] 4d¹⁰ 5s² 5p⁴ Polonium (Po): [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p⁴ Key Trends: - All have six valence electrons. - Oxidation states range from -2 to +4 and +6, with metallic tendencies increasing down the group. Occurrence Selenium: Found in sulfide ores (e.g., pyrite) and obtained during copper refining. Tellurium: Found in tellurides of gold and silver, and extracted as a byproduct of copper refining. Polonium: A rare radioactive element, found in uranium ores or synthesized by bombarding bismuth-209 with neutrons. Physical Properties Selenium exist as either Gray or red solid with density 4.81 g/cm³ and melting point of 221°C. Tellurium is silver-white solid, density of 6.24 g/cm³ and melting point of 450°C. While Polonium is metallic silver solid with density of 9.32 g/cm³ and melting point of 254°C. Ionization energies trend down the group The first ionization of these elements are 941, 869 and 812 kJ/mol for selenium, Tellurium and polonium, respectively. The Ionization energy for these elements decreases down the group due to increasing atomic radius and shielding effect. Chemical Properties 1. Reaction with Hydrogen: Se + H₂ → H₂Se (acidic gas) Te + H₂ → H₂Te Po + H₂ → H₂Po (unstable). 2. Reaction with Oxygen: Se + O₂ → SeO₂ Te + O₂ → TeO₂ Po + O₂ → PoO₂ (stable). 3. Reaction with Acids Selenium reacts with nitric acid, Tellurium reacts similarly. While polonium reacts with hydrochloric acid. Se + 4HNO₃ → SeO₂ + 2NO₂ + 2H₂O Te + 4HNO₃ → TeO₂ + 2NO₂ + 2H₂O Po + 4HCl → PoCl₄ + 2H₂. 4. Complex Formation Selenium and tellurium form complexes with ligands containing oxygen or halides. Polonium, being metallic, forms simpler coordination compounds. Coordination Numbers:  Selenium (4, 6): [SeO₂Cl₂]²⁻ (tetrahedral)  Tellurium (4, 6): [TeCl₄]²⁻ (square planar or octahedral)  Polonium (4): [PoCl₄]²⁻ (tetrahedral) Example of Reactions leading to complex formation:  SeO₂ + H₂O → [Se(OH)₄]  TeCl₄ + 2H₂O → [Te(OH)₄] + 4HCl. Applications Selenium: Photovoltaic cells, photocopiers, and as a micronutrient. Tellurium: Used in thermoelectric devices and as an alloy additive in steel. Polonium: Serves as a heat source in space probes and in neutron sources. Assignment: Discuss the comparative Chemistry of Germanium (Ge), Tin (Sn) and Lead (Pb)