2 Molecular structure.
    1. Lewis structuires: The earliest description of the formation of a covalent bond is that it consists of a shared e- pair. Lewis structures depict schematically how such pairs are shared and give a topological picture of bonding in a molecule.
    2. Formal charge and oxidation number: The formal charge on an atom in a Lewis structure is an indication of the charge it would carry if it shared the e-s in a pair equally. The oxidation number of an element, which independent of the Lewis structure proposed for a species, is an indication of the extent to which e-s transfer to or from an atom in a molecule.
    3. Bond Parameters: Certain properties of bonds, particularly their lengths and their strengths, are approximately transferrable between species.
    4. Molecular orbitals: The principal theoretical description of chemical bonding is in terms of molecular orbitals, which are wavefunctions that spread over two or more atoms. Molecular orbitals are usually approximated by linear combinations of atomic orbitals.
    5. Bonding and antibonding orbitals: A molecular orbital is built by superimposing all the atomic orbitals of the appropriate symmetry. From N atomic orbitals, N molecular orbitals can be formed. Approximatly half will be bonding orbitals, about half will be antibonding orbitals, and the rest are nonbonding orbitals. The greatest bonding and antibonding effects are obtained by overlap of orbitals of similar energies.
    6. Local symmetry classification of orbitals: Molecular orbitals are classified as s, p and d orbitals depending on their symmetry w.r.t. rotation around the internuclear axis.
    7. Electron configurations: The ground state e- configs of molecules are predicted by applying the building - up principle to the molecular orbitals formed by over -  lap of atomic orbitals.
    9. Bond correlations: Certain properties, particularly bond length and bond strength, corralate with the bond order (and with each other).
    10. Electron delocalization: Molecular orbitals in polyatomic molecules are delocalized, and their bonding or antibonding influences are shared over all the atoms in the molecule. Thus, an e- pair can bind more than one pair of atoms together.
    11. Localized descriptions: Although molecular orbitals are delocalized, it is possible to form mathematically equivalent localized descriptions. One way of modeling such localized orbitals is to build them from the overlap of hybrid orbitals on each atom. Such hybrid orbitals have definite geometrical arrangments that depend on their composition.
    12. Isolobality: the concept of localized bonds and hybridization leads to the concept of isolopalitty, which allows certain bonding analogies to be predicted.
    13. Band theory: The concept of molecuklar oorbital formations can be extended to effectivly infinite numbers of atoms in solids, where orbital overlap leads to the formation of bands of orbitals separated by energy gaps.
    14. Electronic conduction: The occupation of bands and the existence of band gaps accounts for the classification of solids as metalic conductors and semiconductors.

3 Molecular Shape and Symmetry.
    1. The VSEPR model: in the valence shell e- pair repulsion (VSEPR) model of molecular shape, it is supposed that e- pairs repel each other and take up positions as far apart as possible. The basic shapes of the theory are modified by allowing for the larger repulsions from lone pairs of e-.
    2. Fluxionality: When different experimental techniques suggest that a molecule may be fluxional on the timescale of the experiment. The resolution of experimental techniques is governed by the lifetimes of the conformations.
    3. Molecular orbitals and molecular shape: The Walsh approach to the explanation of molecular shape is an attempt to identify the origin of shape without the context of delocalized molecular orbitals. In the approach, correlation diagrams for the orbitals are constructed, and the shape that results in the lowest energy is inferred from the variation of energy with bond angle.
    4. Point group assignment: A molecule is assigned to a point group by identifieing the symmetry elements it possesses and working through the flow chart.
    5. Polarity and chirality: Some properties of molecules, particularly their polarity and chirality, can be inferred from theidentity of the point group alone. Other properties (particularly the composition of orbitals and normal modes and selection rules)require a more detaild analysis in terms of charecter tables.
    6. Symmetry - adapted linear combinations: Charecter tables are used to construct symmetry - adapted linear combinations of atomic orbitals as an inital step in the construction of molecular orbitals: only SALCs of the samesymmetry type have nonzero overlap. A4 is a pictorial summery of a number of SALCs.
    7. Vibrational modes: Molecular vibrations are conveniently expressed as normal modes. A nonlinear molecule that consists of N atoms has 3N - 6 modes of vibration: if it is linear, then it has 3N - 5 vibrational modes.
    8. Infrared and Raman activity: A normal mode is infrared active if it corresponds to a change in electric dipole moment of the molecule; it is Raman active if the polarizability changes during vibration. If the molecule has a center of inversion, a mode cannot be both infrared and Raman active.

4 The Structure of Solids.
    1. Unit cells and crystal lattices: The structures of crystalline solids are discussed in terms of a unit cell, the fundamental unit from which the crystal may be regarded as constructed, and the pattern of atoms in a crystal is depicted in terms of the crystal lattice.
    2. The hard - sphere model: The structures of simple solids can sometimes be expressed in terms of a model  in which hard spheres representing ions are stacked together.
    3. Close - packed structures: Many metals hae close - packed structures in which the spheres pack together with least waste of space. Many other substances have structures that can be expressed in terms of less closely packed structures or in terms of the occupation of the tetrahedral and octahedral holes in a close - packed structure.
    4. Alloys: Alloys may be either substitutional or interstitial. Nonmetal may also form interstitial solid solutions in metals.
    5. Intermetallic phases: Some pairs of metals form intermetallic compounds which have definite structures unrelated to the parent compounds. The Zintl phases are a particular case in which a strongly electropositive metal combines with a less electropositive metal.
    6. Typical crystal structures: A number of ionic solids have charecteristic structures that include the rock - -salt structure, the cesium - chloride structure, the sphalerite structure, the fluorite and antifluorite structures, the wurtzite structure, the nickle - arsenide structure, the rutile structure, and the perovskite structure.
    7. Ionic radii: there are several convertions for the definition of ionic radii and valures from different sources should be mixed with care. Large ionic radii tend to favour high coordination numbers.
    8. Structure maps: An empirical rationalisation of structure is in terms of structure maps, in which the axes denote difference in electronegativity and the sizes of the ions.
    9. Lattice enthalpy: A measure of the strength of bonding in a solid is the lattice enthalpy, which is determined by using a born - haber cycle and thermodynamic data. If the observed lattice enthalpy agrees with that calculated in the basis of coulombic interactions between ions, then ionic bonding is suggested (but not guaranteed).
    10. Trends in lattice enthalpies: The born - mayer equation can be used to retionalize trends in lattice enthalpies in terms of charge numbers and ionic radius: large charge number and small radii result in high lattice enthalpies.
    11. Thermal stability - size correlations: Large cations stabilize large polyatomic anions (and vice versa); in particular, the decomposition temperatures of thermally unstable compoounds (such as carbonates) increase with cation radius.
    12. Oxidation number - size correlations: Species with  high oxidation numbers are stabilized by small anions. In particular, fluorine has a greater ability compared with the other halogens to stabilize the high oxidation states of metals. Oxygen helps to stabilize species in high oxidation states.
    13. Solubility - size correlations: Compounds that contain ions with widely different radii are generally soluble in water; conversely, the least water - soluble salts are those of ions with simalar radii.

5 Acids and Bases.
    1. Bro/nsted acids and bases: In the Bro/nsted definition, acids are proton donors and bases are proton acceptors.
    2. Proton transfer equilibria: A bro/nsted equilibrium exists in solution between conjugate acids and bases, and has the form Acid₁ + Base₂ <=> Acid₂ + Base₁.
    3. Acidity and autoprotolysis constants: The strength of a bro/nsted acid is expressed in terms of its acidity constant, K_a, and the extent of self - protonation of water is expressed in terms of autoprotolysis constant of water, K_w.
    4. Solent leveling: Water has a leveling effect that brings the strengths of all stronger acids down to the acid strength of H₃O⁺ and the strengths of all the strong bases down to the strength of OH^-. Similar leveling effects are found in other solvents, such as liquid ammonia and methanol.
    5. Classes of oxoacid: Bro/nsted acids in which the acidic hydrogen atom is attached to an O atom ae classified as aqua acids, hydroxoacids, and oxoacids. The observed strengths of mononuclear oxoacids can be sysematized in terms of Pauling's rules.
    6. Classes of oxides: Oxides are classified as acidic, amphoteric, and basic. The charecter of an oxide varies systematically through the periodic table and (particularly in the d block) with the oxidation state of the element.
    7. Effect of pH on amphoteric oxides: As the pH of a solution is increased, the aqua ions of metals that have amphoteric oxides generally precipitate and then redisolve.
    8. Lewis acids and bases: In the Lewis definition, acids are e- pair acceptors and bases are e- pair donors.
    9. Varieties of Lewis acids: Lewis acids include metal cations in complexes, molecules with an incomplete octet, molecules with a complete octet that can rearrange their e-s to accept another pair, molecules or ions that can expand their octet, and closed shell molecules that can use their antibonding orbitals to accommodate e-s.
    10. Reactions of lewis acids and bases: The three important types of acid - base reaction  are complex formation, displacement (substitution), and double displacement.
    11. Strengths of Lewis acids and bases: Electronic and steric effects determine the strengths of Lewis acids and bases. Electronic effects are summarized by the distinction between hard and soft acids and bases.
    12. E and C parameters: Quantitative, empirical correlation of the thermochemical aspects of complex formation is expressed by the E and C parameters.
    13. Solvent properties: Solvents are usually Lewis acids or bases, and their abilities to act in this manner are summerized quantitatively by the donor and acceptor numbers of the solvent.

6 d-metal complexes.
    1. Complexes: Assemblies of units (ligands) about a central atom are called complexes; they are characterised by their coordination numbers and characteristic geometries. the ligands act as lewis bases and the central metal atom acts as a lewis acid.
    2. Coordination geometries: Important coordination numbers range from 2 to 12 with 4, 5, and 6 being most common for d block ions. The most common geometries include tetrahedral, square planner, trigonal bipyramidal, square pyramidal and octahedral.
    3. Isomerism: The possibility of different spatial arrangments of ligands about a metal atom gives rise to geomeetrical and optical isomerism.
    4. Ligand field splitting: the bonding in complexes can be modeled by crystal field theory in which the ligands are represented by partial negative charges. The theory leads to the ligand field splitting parameter (D_O for octahedral complexes and D_T for tetrahedral complexes).
    5. The spectrochemical series: The size of the lignad field splitting parameter leads to the ordering of ligands in the spectrochemical series.
    6. E- configurations of octahedral complexes: The e- config. of a complex is predicted by applying the building up principle to the d orbitals. If the ligand field splitting is large, then the lowest energy is obtained by filling the lower set of d orbitals (t_2g in octahedral complexes). If the ligand field splitting is small, then the lowest energy is obtained by occupying the upper set (e_g in octahedral symmetry) before pairing e-s in the lower set.
    7. High- and low-spin complexes: When the ligand field splitting parameter is larger than the e- pairing energy, a low- spin complexe results; when the opposite is true, a high-spin complex results. The number of unpaired spins can sometimes be determined by measuring the magnetic susceptibilty of the complex.
    8. Ligand - field splitting parameter: The energy of an e- config. relative to the average energy of the d e-s in a complex is called the ligand field stabilization energy. It accounts for trends in the observed enthalpies of hydration of complexes and provides a means of correlating thermodynamic and spectroscopic data.
    9. E- config. of tetrahedral complexes: The e- config of tetrahedral complexes can be explained in terms of a smaller ligand field splitting parameter and a splitting of the d orbitals into two sets with the e orbitals below the t₂ orbitals. Almost all tetrahedral complexes are high - spin complexes.
    10. Ligand field theory: In ligand field theory, the bonding is modeled in terms of molecular orbitals that are formed by overlap between the d orbitals of the central metal ion and SALCs of the ligand orbitals. The ligand field splitting parameter is identified with the energy separation of the frontier orbitals of the complex.
    11. p Bonding: Ligand field theory accommodates the effects of p bonding between the ligand and the metal ion. When the ligand acts as a p donor, the ligand field splitting parameter is decreased, and when the ligand acts as a p acceptor, the parameter is increased.
    12. The spectrochemical series: The effects of p bonding correlate with the position of a ligand in the spectrochemical series, with p-donor ligands lying low in the series and p- acceptor ligands lying high in the series. Very strong s donors, such as CH^-₃ and H^-, are also high in the series.
    13. Spectroscopic, magnetic and thermochemical correlations: The ligand field splitting parameter can be used to correlate the trends in electronic spectra, magnetic properties, and some thermochemical properties of complexes of d-block metals.
    14. Formation constants: the thermodynamic stabilities of complexes are expressed in terms of stepwise and overall formation constant. Stepwise formation constants K_n commonly decrease as n increases.
    15. Chelate effects: Chelate complexes are generally more stable than their nonchelated analogs.
    16. Irving - williams series: The irving - williams series summerarizes the variation in thermodynamic stability of complexes with change in the central metal ion.
    17. Electrophiles and nucleophiles: Incoming groups are classified as electrophiles or nucleophiles; their relative strengths are assessed by comparing rate constants for substitution reactions.
    18. Associative substitution: In an associative substitution reaction the initial complex passes through an activated complex with an increased coordination number.
    19. Dissociative substitution: In a dissociative substitution reaction the activated complex has a lower coordination number than the initial complex. Most octahedral substitutions are dissociative rather than associative.
    20. Relative labilities: Among aqua metal ions, the ones that form relatively weak bonds (on account of their low charge and large ionic radius) are more labile than those with high charge and small radius.

7 Oxidation and Reduction.
    1. Oxidation and Reduction: Oxidation is the loss of e-; reduction is the gain of e-. These definition are generalizations of the original definitions based on reaction with oxygen to produce an oxide.
    2. Thermodynamic aspects of metal production: Many elements exists in nature as oxides or as sulfides that are easily converted to oxides. An Ellingham diagram summerizes the temperature variation of the thermodynamic aspects of the reaction of metal oxides to the element.
    3. Varieties of industrial element production: Industrial extraction of elements from ores uses low temperature carbon reduction for readily reduced speicies (Zinc), higher temperature carbon reduction for less readily reduced species (iron), and high temperature (electric furnace) carbon reduction for some (magnesium) or electrolysis (aluminum) for least readily reduced species. Several elements (eg haalogens) exist in a reduced state and are extracted by oxidation, which is commonly electrolytic.
    4. Reduction half-reactions and standard potentials: The thermodynamics of redox reactions in solution are conveniently tabulated in terms of reduction half - reactions and their standard potentials E^q. The overall standard potential of an electrochemichal cell is proportional to DG^q for the cell reaction. The Nernst equation summarizes the dependence of potentials on concentrations.
    5. Redox stability in solution: The stability of a species in solution refers to its thermodynamic tendancy toward oxidation or rduction by the medium or to disproportionation.
    6. Overpotential: The kinetics of redox reactions often determine the observed processes. An overpotential is often required to achieve a significant reaction rate.
    7. Reaction mechanisms: The mechanisms of redox reactions include outersphere e- transfer, inner - sphere e- trnsfer, and atom or group tansfer.
    8. Latimer, Frost, and Pourbaix diagrams: Latimer diagrams provide compact portrayal of E^q values for an element. Frost diagrams provide a striking representation of trends in the stabilities of oxidation states of elements. Pourbaix diagrams provide a summary of the influence of pH and potentials on the identity of the predominant species.
    9. The effect of pH and complexation: Where proton transfer or complexation occurs together with redox processes, potentials depend on pH and complex formation.
    10. Natural waters: The behavior of natural water systems is conveniently organized in terms of dependence on potential and pH (and expressed in terms of Pourbaix diagrams).

8 The Metals.
    1. Group 1 and 2 metal properties: the s-block elements ar low - melting and highly electropositive metals with the group oxidation state dominant. The principal exceptions are beryllium and magnesium, which are physically harder and kinetically less reactive.
    2. Nonaqueous solutions of group 1 and 2 metals: Alkali metals dissolve in low acidity solvents (such as liquid ammonia) to yield highly reducing solutions that contain the metal cationand e^- (solv) or sometimes the M^-(solv) anion. Metal solutions also can be prepared for the hevier group 2 metals.
    3. Oxidation state stability in the d block: In aqueous redox chemistry, the M²⁺ ion is favored toward the right of the 3d metals, and the highest oxidation state becomes more stable on descending each of the Groups 4 through 8.
    4. Coordination number in the d block: In contrast to the common occurrence of 4 and 6 - coordinate 3d cations, those in the 4d and 5d series are larger and often display higher coordination numbers.
    5. Polyoxometallates: he Polyoxometallates observed for group 5 and 6 metals, especially vanadium, molybdenum, and tungsten, usually consist of edge- and vertex- shared [MO₆]^n- octahedra.
    6. Metal - metal bonding in the d block: Metal - metal bonded compounds are encountered for low - oxidation state metals on the left of the d block in conjunction with p-donar ligands such as halide and for metals on the right of the d block combineed with p-acceptor ligands.
    7. Noble character: The noble metals and the coinage metals are not oxidized by H ions in acids. Square - planar structures are the rule for their d⁸ complexes and octahedral structures are common for their d⁶ complexes.
    8. Oxidation states of p block metals: The heavy p-block metals thallium, lead, and bismuth are most stable in an oxidation state that is two less than the group oxidation state.
    9. Oxidation states in the f block: The +3 oxidation state predominates for the lanthanides and the heavy actinides. For uranium, oxidation states +4 and +6 are most important.
    10. f-Block metal complexes: Because of their large size, high coorination umbers are common for 4f and 5f cations. Linear dioxo complexes are characteristic of the actinides in +5 and +6 oxidation states.
 
9 Hydrogen and its compounds.
    1. Isotope effects: The effects of isotopic substitution on spectra and reactions are more pronounced for H compounds than for others on account of the large percentage mass change when H is replaced by D or T.
    2. Production of H₂: The commercial production of H is currently based mainly on steam reforming and dehydrogenation of hydrocarbons.
    3. Classification: The H compounds are classified as saline, metallic, and molecular. Saline hydrides are formed by s-block elements, metallic hydrides by many d- and f-block elements, and molecular hydrides by many p-block elements.
    4. Stabilities: the stabilities of the p-block hydrides relative to their elements decreases markedly down a group. The endoergic character of many p-block hydrides necessitates indirect synthetic routes for their preparation.
    5. Reactoin patterns: The reactions of H include homolytic cleavage on metal surfaces and some e- rich metal complexes, heterolytic cleavage on surfaces such as ZnO, and radical chain reactios with many nonmetals in the gas phase.
    6. Synthesis of H compounds:The synthesis of compounds of H is generally based on one of three strategies: direct combination of the elements, protonation of a basic-ide salt, or double replacement between a metal hydride and halide of a more electronegative element.
    7. Hydrogen bonding: H bonding is responsible for the low volatility and high electric permittivities of NH₃, H₂O, and HF. It also results in the organisation of water molecules in ice and clathrate hydrates.

10 Main-Group organometallic compounds.
    1. Classidication: The organometalic compounds of the s-block elements range from molecular or polymeric lithium, beryllium and magnesium compounds, which have highly polar bonds M^d+C^d-, to the even more ionic compounds of the heavier metals. Themore covalent compounds of the p block are classified as e- poor through group 13/III, e- precise for Group 14/IV, and e- rich for many compounds in Group 15/V.
    2. Structures: The highly useful lithium organometallic compounds form cluster compounds in non - coordinating or weakly coordinating solvents.
    3. Stability: The thermodynamic stability of organometallic compounds (as revealed by Gibbs free energies of formation) decreases down a group.
    4. Synthesis: The primary routes to organometallic compounds are: (1) Active metal plus an organohalide: M +RX (X = halogen); (2) Transmetallation: M+M'R (where M is more electropositive than M'); (3) Double replacement (metathesis): MR + EX (where M is more electropositive than E); (4) Hydrometallation: MH + RHC = CH₂ (where MH is a molecular hydride, often borane or silane).
    5. s-block reactivity: the s-block organometallic compounds are highly reactive with proton sources and are good carbanion sources in metathesis reactions with the halides of more electronegative elements. Organolithium and grignard reagents are the most frequently used carbanion reagents in synthesis.
    6. Group13/III structures: BR₃ and GaR₃ compounds are unassociated trigonal and planar. triorganoaluminum compounds are usally associated through alkyl or aryl bridges.
    7. Group 13/III Reactivity: cabanion character follows the order AlR₃ > GaR₃ >> BR₃; accordingly, trialkylaluminum compounds are frequently used in double replacement reactions with the halides of more eletronegative elements.
    8. Group 14/IV structures: Silicon, germanium,, and lead form tetrahedral MR₄ compounds. With bulky groups, R₂E = ER₂ (E = Si, Ge, or Sn) is formed.
    9.Group 14/IV reactivity: Associative substittion reactions are characteristic of organometallic compounds such as SiR₃Br and the greater access to the larger Si atom accounts for greater reactivity of silicon compounds than their carbon analogs.
    10. Group 15/V structures: The compounds AsR₃, SbR₃, and BiR₃ of Group 15/V are trigonal pyramidal, with a stereo - chemically active lone pair. The less common ER₅ organometallic compounds of the group are either trigonal bipyramidal or square pyramidal.
    11. Group 15/V Reactivity: the compounds AsR₃, and to a lesser extent SbR₃, form complexes with many soft d-block metals. These formal +3 oxidation state compounds can be converted to formal +5 species such as (C₆H₅)₃AsO and [(C₆H₅)₄As]⁺.

11 The Boron and Carbon groups.
    1. The Elements: In groups 13/III and 14/IV the light elements are distinctly nonmetals and the heavy elements are metals. The borderline occurs between boron and aluminium in group 13/III and between germanium and tin in group 14/IV.
    2. Extraction: Elemental boron (group 13/III) is not important commercially. In group 14/IV, carbon from graphite mining or pyrolysis of hydrocarbons or coal, finds many uses. Reduction by carbon is used in the recovery of the other commercially important elements. Silica requires very high temperatures for its reduction, whereas the heavy metal oxides of tin and lead are reduced at moderate temperatures.
    3. Boron halides: the simple trihalides are notable for their lewis acidity and by the facile halide nucleophilic displacement reactions they undergo, to produce BR₃ and [BR₄]^-, where R = alkyl, aryl, alkoxy, amido, and the like. A series of boron subhalides is know, the most important of which have the formula B₂X₄.
    4. Boron oxides: Boron has a high affinity for oxygen and is mined in the form of metal borates. Boron oxide, B₂O₃, in combination with SiO₂ and alkali metal oxides is widely used to make glass which is tolerant to thermal shock. The coordination number of B in oxides may be either 3 or 4.
    5. Boron nitrogen compounds: Boron nitrogen compounds are often structural analogs of carbon compounds because BN is isoelectronic with CC. Thus, one form of BN has a structure similar to graphite and another resembles diamond. Other BN analogs of carbon compounds include Lewis acid - base complexes (such as H₃NBF₃) and analogs of alkenes and amino acids.
    6. Boranes and carboraes: The boranes constitute a large group of compounds which are conveniently represented by delocalized bonding, such as the 3 - center, 2 - e- BHB bond. Most of them fall into one of three classes: closo, [B_nH_n]^-; nido, B_nH_n+4; andarachno, B_nH_n+6. An isoelectronic series of carboranes is known, which formally represents the replacement of BH^- by the isolobal CH group.
    7. Allotropes of carbon: Elemental carbon exists in many forms, ranging from the extended crystalline solids, diamond and graphite, to semicrystalline analogs of graphite and clusters such as C₆₀.
    8. Carbides: Carbon compounds include the familiar hydrocarbons and halogenated hydrocarbons, and simple oxygen and nitrogen compounds. In addition, carbon forms a series of compounds with metals and metalloids: saline carbides, which are formed by highly electropositive metals display ionic character in many of their reactions; metallic carbides, formed with many d-block metals are often hard materials with high electronic conductivity; metalloid carbides are often hard covalent insulating or semiconducting solids, such as SiC.
    9. Silicon and germanium halides: The halides of silicon and germanium are notable for their mild Lewis acidity,  which is achieved by hypervalence, and their utility as starting materials by the nucleophilic displacement of halide to form compounds such as Si(CH₃)₄ and SiH₄.
    10. Silicates: the silicates are metal - silicon - oxygen compounds that contain the tetrahedral SiO₄ building block that may share one or two of its O atoms with an adjacent SiO₄. These arrangments correspond to shared tetrahedral vertices or edges respectivly. In SiO₂, the four O atoms around a Si atom have four surrounding SiO₄ units. The progressive introduction of alkali metal and alkaline earth metal oxide to form metal silicates leads to decreased sharing of O atoms between the SiO₄ tetrahedra.
    11. Aluminosilicates: Aluminosilicates, together with silicates, constitute the bulk of the minerals in the Earth's crust. A vertex- or edge-shared MO₄ tetrahedron is the building block, where M is either Si or Al. Aluminosilicate minerals include the three - dimensional feldspars; many two - d minerals, such as clays, talc, and mica; and a few with one - d topology.
    12. Molecular sieves: Most molecular sieves are aluminosilicate minerals or synthetic compounds that have open structures through which small molecules can diffuse. These compounds find use as desiccants through their water absorbing properties, ion exchange materials, and solid acid catalysts.

12 The Nitrogen and Oxygen groups.
    1. Physical state: In contrast to the gaseous diatomic species N₂ and O₂, the heavier elements in groups 15/V and 16/ VI are solids under normal conditions. Atmospheric nitrogen and oxygen are separated on a massive scale by low temperature distillation for use as a refigerent and inert gas (N₂) and in steel making (O₂).
    2. Elemental nitrogen: Nitrogen has low reactivity but bacteria manage to reduce it at room temperature. The commercial Haber process requires high temperatures and pressures to yeild ammonia, which is a major ingredient in fertilizers and an important chemical intermediate.
    3. Elemental Phosphorus: Elemental phosphorus is recovered from the mineral apatite by carbon arc reduction. The resulting white phophorus is a molecular solid (P₄). Treatment of apatite with sulphuric acid yeilds phosphoric acid which is converted to fertilizers and other chemicals.
    4. Halides: The halides of nitrogen and oxygen have limited stability, but their heavier congeners form an extensive series of compounds. Typical formulas are EX₃ and EX₅ for group 15/V and EX₂, EX₄, and EX₆ for group 16/VI. Nitrogen, oxygen, and bismuth have somewhat limited halide chemistry because of their resistance to oxidation.
    5. Positive oxidation states of Nitrogen: Such species are produced by the oxidation of ammonia rather than elemental nitrogen on account of the kinetic inertness of the later.
    6. Nitrogen oxoanions: Nitrate, NO₃^-, and nitrite, NO₂^-, are the most important oxoanions of nitrogen. Their reactions can be sluggish, especially those of nitrate. The mechanisms often involve atom transfer and rates are facilitated by low pH.
    7. Phophorus oxides and oxoanions: The oxides of phosphorus include P₄O₆ and P₄O₁₀, both of which are cage compounds with T_d symmetry. Important oxoanions are the P(I) species hypophosphite, H₂PO₂^-, the P(III) species phosphite HPO₃^2-, and the P(V) species phosphate, PO₄^3-. The existence of P - H bonds and highly reducing character of the two lower oxidation states is notable. Phosphorus(V) also forms an extensive series of O-bridged polyphosphates. In contrast to N(V), P(V) species are not strongly oxidizing.
    8. Arsenic: Arsenic(V) is more easily reduced than P(V), which may account for the toxicity of AsO₄^3-, which mimics PO₄^3- sufficiently to enter biological cells.
    9. Phosphorus - nitrogen compounds: The range of PN compounds is extensive, and includes cyclic and polymeric phosphazenes, (-PX₂=N-)_n.
    10. Oxygen allotopes: Oxygen has two allotropes, dioxygen (O₂) and ozone (O₃). Dioxygen has a triplet ground state which oxidizes hydrocarbons by a radical chain mechanism. Reaction with an excited state molecule can produce a fairly long lived singlet state species, which is found in photochemical smog and often reacts as an electrophile. Ozone, is an unstable and highly aggressive oxidizing agent.
    11. Allotropes and polymorphs of sulphur: Sulfur has many allotropic forms including rings and metastable polymer; its solid form is polymorphic.
    12. Sulfur oxoanions: The oxoanions of sulfur include the good reducing agent sulfite ion, SO₃^2-, the rather unreactive sulfate ion, SO₄^2-, and the strongly oxidizing peroxosulfate ion, O₃S-OO-SO₃^2-.
    13. Metal oxides: The oxides formed by metals include the basic oxides with high oxygen coordination number that are formed with most M⁺ and M²⁺ ions. Oxides of metals in intermediate oxidation states often have more complex structures and are amphoteric. Ruthenium and osmium tetroxides, MO₄, are molecular and not basic.
    14. Cyclic and cluster compounds: Cyclic and cluster compounds are known for many of the heavier p-block elements. The valence electron count and structures of group 14/IV anionic clusters can often be correlated by Wade's rules.

13 The halogens and the noble gases.
    1. Structures of the elements: There is greater structural uniformity for the elements within the halogen group and those within the noble gas group then for any other p-block groups. The former are diatomic molecules and uniformly nonmetals, whereas the latter are monatomic gases with low reactivity.
    2. Recovery and use of halogens: Chlorine, bromine, and a high proportion of iodine is produced by oxidation of the halides. Chlorine is by far the most commercially important halogen, although its uses in the preparation of chlorocarbon compounds is coming under close scrutiny for environmental reasons.
    3. Recovery and use of noble gases: Exept for helium, which is extracted from gas wells, and radon, which is radioactive, the noble gases are extracted from the atmosphere by distillation. Helium and argon are used as inert gases and helium is used as a refrigerant.
    4. special properties of fluorine: The distinctive properties of fluorine include the stabilization of metal ions in their high oxidation states, the high volatility of many molecular fluorine compounds, and strong e- withdrawing effects on covalent compounds, which give rise to strong proton and Lewis acidity of fluorine containing acids.
    5. Polyiodides and interhalogens: Polyiodides most of which have the formular I_n^-, with n odd, may be veiwed as loosely associated aggregates of I₃^-, I₂, and I^-; the most important species is I₃^-. Many interhalogen compounds exist, most of which are composed of a central Cl, Br or I atom bound to periphal F atoms. The structures are generally predictable from the VSEPR model. Some interhalogen compounds, such as ClF₃ and BrF₃, are aggressive fluorinating agents.
    6. Halogen oxoanions: the halogen oxoanions are the most important oxygen - containing halogen compounds.
    7. Redox reactions of halogen compounds: Halogen oxides and oxoanions are often thermodynamically potent oxidizing agents. Many of the intermediate oxidation states undergo redox disproportionation.
    8. Reaction mechanisms of halogen oxoanions: Atom transfer mechanisms are common, such as Cl transfer from the facile oxidant OCl^- and O transfer from the sluggish but potent oxidant ClO₄^-.
    9. Compounds of noble gases: Far more compounds are known for Xenon than for any of the other noble gases. The fluorides XeF_n, with n= 2, 4, and 6, are the most important, but Xe-O, Xe-C, and Xe-N bonds are known.

14 Electronic spectra of complexes.
    1. Atomic energy levels: Atomic orbital configurations do not completely specify the state and energy levels of an atom: interelectronic repulsions must be taken into account. Angular momentum quantum numbers aid in grouping microstates that have similar energies.
    2. Racah parameters: Interelectronic repulsion effects can be summarized with a small number of parameters called Racah parameters, which are known from atomic spectra.
    3. Ligand field transitions: Electronic transitions between the HOMO and LUMO of compexes can often be assigned as ligand field transitions between the d orbitals of the metal atom. A consideration of ligand field stabilization energies and Racah parameters is used to identify the transitions. In complexes of low symmetry, the HOMO and LUMO are not the only orbitals involved in transitions.
    4. Charge transfer transitions: The other major source of low wnergy electronic transitions is charge transfer between the metal atom and the ligands (CT bands).
    5. Selection rules: Spectral intensities depend upon selection rules. These rules are based on the conservation of spin of the complex and on considerations of symmetry by using group theory.
    6. Specroscopic techniques: The Specroscopic techniques for investigation of d-block complexes include electronic absorption spectra, luminescence, circular dichroism, Raman scattering, and EPR.
    7. Metal - metal bonding: The electronic structure of metal - metal bonded systems may be expressed conceptually by the consideration of the frontier orbitals of fragments that contain a single metal atom.
    8. Mixed - valence complexes: Intense intervalence charge transfer absorption bands are found in multinuclear system with metal atoms in more than one oxidation state.

15 Reaction mechanisms of d-block complexes.
    1. Ligand substitution reactions: Ligand substitution reactions involve the replacement of a coordinated ligand by an incoming ligand from solution.
    2. Types of mechanism: Stoichiometric mechanisms are designated D (dissociative) if the first step is dissociation (loss) of a coordinated ligand, A (associative) if the first step is coordination of an incoming ligand, and I (interchange) if the reaction involves addition of the incoming ligand with an expansion of the coordination shell in the activated complex (an energy maximum). Inerchanges are further differentiated as Ia if the activated complexis primarily favoured by bond making with the incoming ligand, or Id if the energy maximum is determined primarily by breaking the bond with the outgoing ligand.
    3. Identification of mechanism: Mechanisms are inferred from the concentration, temperature, and pressure dependence of rates, from trends in rates as a function of incoming and outgoing ligand, and from the stereochemistry of reactants and properties.
    4. The trans effect: The structure of the products may be controlled by the trans effect in substitutions on square planar complexes.
    5. Isomerization reactions: Isomerization reactions may occur by substitution pathway or by intermolecular rearrangements. Two of the latter for octahedral complexes are the Bailar twist and the Ray - Dutt twist. The Berry pseudorotation is common in five - coordinate complexes.
    6. Redox reaction mechanisms: redox reactions involve changes in oxidation state of the reacting complexes. Inner - sphere redox reactions involve a transient bridging ligand, which carries e- density from one metal atom to the other. Outer - sphere redox reactions involve e- transfer from one reactant species to the other.
    7. Photochemical processes: Photochemical reactions involve the absorption of light to produce a highly reactive electronic state.
    8. Prompt reactions: Prompt photochemical reactions often involve the rapid dissociation of a ligand upon photoexitation. These photosubstitution reactions often result from d - d transitions.
    9. Delayed reactions: Delayed photochemical reactions often involve e- transfer, either ligand - to - metal or metal - to - ligand. Long - lived excited states associated with electronic transitions in metal - metal bonded compounds may also lead to delayed photochemical reactions.

16 d- and f-block organometalic compounds.
    1. Organometalic compounds: Organometallic compounds contain at least one metal - carbon bond.
    2. Hapticity: the Hapticity h of a compound designates the number of points of attachment of an organic ligand to a metal atom.
    3. the 16/18 electron rule: d-block organometallic compounds of metals on groups 6 to 8 generally have 18 valence electrons around the central metal atom. For groups 9 and 10, the norm is 16 or 18 electrons.
    4. CO - metal interaction: the Ligand CO is counted as a two - e- s-donor. Additionally, it is a net acceptor of e-s via p interaction from the coordinated metal atom. It usually bonds through C to one, two or three metal atoms.
    5. Metal carboyl syntheses: Metal carbonyls are usually synthesized by direct reaction of CO with the metal or reaction between CO and a metal salt in the presence of a reducing agent.
    6. Metal carbonyl IR spectra: The number of observed IR bands correlates with the structures of substituted mononuclear carbonyls: CO stretching frequencies
                                                       \         \
are lowered by metal bridging (-CO>CO>-CO) and by electron density on the metal (neutral > Mononegative > Dinegative complexes).
                                                       /         /
    7. Substitution of CO: Substitution of CO by other ligands generally occurs by initial CO dissociation, and is promoted by heat or photolysis.
    8. b-hydrogen elimination: b-hydrogen elimination from an ethyl or larger hydrocarbon ligand produces  an M-H bond and freee alkene.
    9. Alkkylidene ligands: Alkylidene ligands have a formal M=C double bond. Alkylidenes designated Fischer carbenes are formed with mid to late d-block metals and are electrophilic at the metal - bound C atom. Schrock carbenes are formed with early d-block metals and the C atom is nucleophilic.
    10. Alkene and polyene ligands: For electron counting purposes, alkene and polyene igands donate two - e-s for each metal bound C-C group.
    11. Early d- and f-block organometallic compounds: The organometallic compounds of early d- and f- block elements are highly oxophilic, and are aggressive C-H bond cleavers.
    12. Metal clusters: Metal clusters involve direct M-M bonds and generally obey the 18 - e- rule for M₅ and smaller clusters. The structures of clusters are correlated with e- count by the Wade - Mingos - Lauher rules.
    13. Metal cluster syntheses: Metal clusters may be synthesized by thermal ligand dissociation followed by M-M bond formation between the coordinatively unsaturated fragmetns, redox condensation, or unsaturated MM or MC bonds.
    14. Multi - metal interactions: Such interactions may promote reactivity of alkyl or CO ligands.

17 Catalysis.
    1. Definition of catalysis: Catalysis may be viewed as a cyclic process in which a species, the catalyst, increases the rate of conversion of reactants to products and is constantly regenerated.
    2. Action of catalysts: The catalyzed reaction folows a different path from the uncatalyzed reaction, and has a lower activation free energy. A catalyst does not infuence the postition of equilm. of a reaction.
    3. Homogeneous and heterogeneous catalysis: In homogeneous catalysis the catalyst and reactants are in the same phase, usually liquid but sometimes gas. Heterogenous catalysts are present in a different phase from from the reactants. Most hetrogeneous catalysts are solids, which are easily separated from liquid or gaseous products.
    4. Metal complexes as homogeneous catalysts: Some common steps in the catalysis of hydrocarbon reactions by metal complexes are coordination of a reactant to the metal, insertion of a group into a metal - carbon bond (or its reverse, reductive elimination), and dissociation of a product from the metal center.
    5. Since the catalyst is not consumed, it is convenient to depict catalytic processes as a cycle in which the metal complex interacts with reactants and eventually produces products with regeneration of the catalyst.
    6. Important homogeneous catalytic reactions: Examples include alkene hydrogenation, hydroformylation (combination of an alkene with CO and hydrogen), acetic acid synthesis from methanol and carbon monoxide, the formation of acetaldehyde from ethene and carbon monoxide, and alkene polymerization.
    7. Commercial advantages of heterogeneous catalysis: Solid catalysts (often consisting of metal particles on a solid or acid sites on a solid) are often preferred in commercial processes because the separation of catalyst from products is simple and therefore economical.
    8. Uniform and mulitiphasic heterogenous catalysts:  Uniform solid catalysts consist of a single phase material such as a reactive zeolite. Multiphasic catalysts often consist in large proportoin of a more or less inert high - surface - area - support ( often silica or g - alumina) which holds fine particles of a catalytic species, such as platinum or rhodium metal.
    9. Mechanism of heterogenous catalysis: Chemisorption brings a reactive molecule on to a surface and modifies its character. Products are released from the surface by desorption. Surface migration and combination are other steps characteistic if many heterogeneous catalytic reactions.
    10. Important heterogeneous catalytic reactions: Some of the many important heterogeneous catalytic processes are oxidation of SO₂ to SO₃ in sulfuric acid production, hdrogenation of alkenes, conversion of hydrogen and N to ammonia, polymerisation of alkenes, catalytic total oxidation of hydrocarbon pollutants (automobile catalyic converterrs), surface acid catalyzed hydrocarbon isomerasation, and electrochemical catalysis as in chlorine production.

18 Structures and properties of solids.
    1. Defects: Point defects in structure are found in solids and may influence properties such as ionic migration and colour. Two common point defects are the Schottky defect, which is a vacancy created by the displacement of an atom from its normal site to the surface of the crystal and the Frenkel defect, which is the displacement of an atom from its normal site to an interstitial site in the crystal.
    2. Nonstoichiometric compounds: Nonstoichiometric compounds deviate from simple whole - number ratios for the constituent atoms in the compound. they are common for the H compounds of early d-block metals and for oxides of metals that can readily adopt more than one oxidation state (such as iron).
    3. Solid electrolytes: in solid electrolytes one or more types of iron diffuses at an appreciable rate. Examples are Ag₂HgI₄ above 50 ^oC and sodium b-alumina where Ag⁺ and Na⁺, respectively, are mobile.
    4. Cooperative magnetism: Ferromagnetism and antiferromagnetism involve the long range order of unpaired  e-s on the individual ions.
    5 Superconductiviry:Superconductivity is observed for many meals and some compounds of nonmetals. High temperature superconductors are metal oxides that have structures related to that of perovskite.
    6. Glasses: the vitreous phase of materials usually contain local order (eg four O atoms around a Si atom) but no long range order.
    7. Metal sulfides: The sulfides of some metals exist in layered structures and some of them (TaS2, for instance) can accommodate ions or molecules between the layers by the process of intercalation or insertion. Intercalation can be performed chemically or electrochemically.

19 Bioinorganic chemistry.
    1. Inorganic components of organisms: