P.A. Cox - Inorganic chemistry (793955), страница 10
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For simple ionic compounds the oxidation state is equal to thecharge on the ions. Roman numbers should be used to distinguish oxidation states from ionic charges, e.g. NaI and Cl−Iin NaCl. In polar covalent bonds the electrons are assigned to the more electronegative atoms as if the bonding wereionic. Oxidation states are calculated by the following rules.(i) Bonds between the same element are not counted. Elements have oxidation state zero. In an ion such as peroxidethe electrons in the O—O bond are distributed equally, making O−I.(ii) Except in cases such asthe most electronegative and electropositive elements in a compound have anoxidation state equal to their normal ionic charge: KI, CaII, F−I, O−II.(iii) The sum of the oxidation states must equal the charge on the species, and is therefore zero in a neutral compound.Using this rule and (ii) above, we have HI in H2O, H−I in CaH2 and MnVII in.(iv) Complex formation, and donor-acceptor interaction in general (see Topic C8) do not alter the oxidation state.Both [Ni(NH3)6]2+ and [Ni(CN)4]2− have NiII, complexed by NH3 and CN− respectively.A redox reaction is any reaction involving changes of oxidation state.
In Equation 1 the changes are from Zn0 to ZnIIand from HI to H0. The reactionis not a redox reaction as no change of oxidation state takes place.Specifying the oxidation state of an element can be a useful way of naming compounds, especially when variablestoichiometries are possible (see Topic B5). Thus we have iron(II) chloride (FeCl2) and iron(III) chloride (FeCl3). Theolder names ‘ferrous’ and ‘ferric’ respectively are still encountered for such compounds but are potentially confusing.In current terminology the -ous and -ic suffixes (referring to a lower and a higher oxidation state, respectively) are onlyused for some oxoacids (e.g.
H2SIVO3, sulfurous acid, and H2SVIO4, sulfuric acid; see Topic F7).Balancing redox reactionsIn any complete redox reaction the changes in oxidation state must balance so that the totals on the two sides arethe same. Difficulties can arise with ions in solution, as the ionic charges may not be the same as the oxidation states.Consider the unbalanced redox reaction in acidified aqueous solution:B4—OXIDATION AND REDUCTION39It is easiest to balance the redox changes by first splitting this into two half reactions, one involving oxidation, and theother reduction.
The oxidation step is(3)with electrons (e−) being removed. The conversion ofto Mn2+ involves a change of oxidation state from MnVIIto MnII and so is a reduction requiring five electrons. To balance the half reactionfour oxygen atoms are required on the right-hand side, which (in aqueous solution) will be in the form of H2O. Thereaction(4)is then completed by balancing hydrogen with 8H+ on the left-hand side, as this reaction takes place in acid. The overallredox reaction is now written by combining the two half reactions in such a way that the free electrons are eliminated.This requires 5 moles of Equation 3 to every 1 mole of Equation 4, givingIn alkaline solution it is more appropriate to use OH− rather than H+ (see Topic E2).
The other species present may alsobe different from those in acid, as many metal cations form insoluble hydroxides or even oxoanions (see Topic E4). Asan example, consider the reaction of aluminum metal with water to form [AlIII(OH)4]− and H2. The balanced halfreactions areandwhich may be combined in the appropriate proportions (two to three) to giveA particular advantage of the half-reaction approach is that it leads naturally to the discussion of the thermodynamics ofredox reactions in terms of electrode potentials (see Topic E5).Extraction of the elementsVery few elements occur naturally in uncombined form (see Topic J2).
Most are found in compounds where they are inpositive or (less often) negative oxidation states (e.g. TiIV, ZnII and Cl−I in TiO2, ZnS and NaCl, respectively).Extraction of these elements therefore requires redox chemistry, using appropriate reducing or oxidizing agents.Thermodynamic considerations are very important (see Topic B3).Iron is produced in greater quantities than any other metal, by reduction of Fe2O3 with carbon (coke). The overallreaction approximates to40SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESAt 25°C, ΔGΘ for this reaction is +151 kJ mol−1 so that it is not thermodynamically feasible at room temperature.However, it is strongly endothermic (ΔHΘ=+234 kJ mol−1) and so by Le Chatelier’s principle the equilibrium is shiftedin favor of the products at higher temperatures.
In a blast furnace it takes place above 1000°C, heat being provided fromthe combustion of carbon in air, which is blown through the reaction mixture.Carbon is a convenient and cheap reducing agent for metal oxides, but for many elements it cannot be used. Withsome highly electropositive metals (e.g.
Al) the oxide is too stable (i.e. itsis too negative), and the temperaturerequired for reduction by carbon is too high to be technically or economically viable. Some elements (e.g. Ti) react withcarbon to form a carbide. In these cases other redox processes are necessary. Table 1 summarizes the common methods.Hydrogen can be used to reduce oxides or halides, or a very strongly reducing metal such as sodium or calcium toreduce halides.In electrolysis a redox process with positive ΔG is induced by providing electrical energy.
Reduction takes place atthe cathode (the negative electrode, which provides electrons), and oxidation at the anode (the positive electrode).For example, electrolysis of molten NaCl gives elemental Na at the cathode and Cl2 at the anode. Many veryelectropositive elements (e.g. Na, Ca, Al) and a few very electronegative ones (F, Cl) are obtained by this method.Table 1. Extraction of elements from their compoundsSection B—Introduction to inorganic substancesB5DESCRIBING INORGANIC COMPOUNDSKey NotesFormulaeNamesStructure and bondingRelated topicsStoichiometric (empirical) formulae describe only the relative numbersof atoms present.
Molecular formulae and/or representations givingstructural information should be used when they are appropriate. Thephysical state of a substance is often specified.Systematic nomenclature can be based on three systems, binary,substitutive (similar to that in organic chemistry) or coordination.Many nonsystematic or trivial names are used.The coordination number and geometry of an atom describe thenumber of bonded atoms and their arrangement in space. Oxidationstates rather than valencies are generally used for describing differentpossible stoichiometries.Methods of characterizationOxygen (F7)(B7)Complexes: structure andHydrogen (F2)isomerism (H6)FormulaeIt is important to distinguish the stoichiometric or empirical formula of a molecular substance from itsmolecular formula.
The former expresses only the relative numbers of atoms present, in the simplest possible ratio.For example, the compound of stoichiometry P2O5 contains P4O10 molecules. Molecular formulae should be used whenthey are known. Methods for determining empirical and molecular formulae are described in Topic B7. On the otherhand, in a solid where clear molecular or other units do not exist the empirical formula is generally used. For example,NaCl is an ionic substance and the formula does not imply that molecules are present.When solids contain identifiable groups such as molecules or complex ions the formula is written to indicate this:for example, NH4NO3 is much more informative for ammonium nitrate than the empirical formula N2H4O3.
This isoften used in molecular formulae, for example, in NH2OH (1) and Ni(CO)4, which are intended to show the groupingsof atoms present. For coordination compounds formed by transition metals formulae are written with squarebrackets as in [Ni(NH3)6]Br2, which indicates that six NH3 are attached directly to Ni, but not the two Br. Complexions formed by main-group elements can be written in a similar way, for example, [PCl4]+ and [BF4]−, although usage isnot very systematic.42SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESWhen a metallic and a nonmetallic element are present, the metallic one is always written first, as in NaCl and PbO2.For compounds between two or more nonmetals they are listed conventionally in the following order, based roughly ona sequence of increasing electronegativity:For example, we have OF2 and ClO2, which are therefore called oxygen difluoride and chlorine dioxide, respectively(see below).When the physical state of a substance is important it is specified as in NaCl(s), H2O(l) and HCl(g) for solids,liquids and gases, respectively.
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