P.A. Cox - Inorganic chemistry (793955), страница 34
Текст из файла (страница 34)
In Equation 2 n=2, which gives K around 1024 at 298 K for this reaction. As the potential of asingle half-cell is not measurable, so an equilibrium constant based on a single potential has no meaning.Nonstandard conditionsUnder nonstandard conditions the electrode potential of a couple can be calculated from the Nernst equation:146E5—ELECTRODE POTENTIALSwhere [Ox] and [Red] are the activities of the species involved; it is a common approximation, especially in dilutesolutions, to assume that these are the same as the molar concentrations.
With n=1 at 298 K, a factor of 10 difference inthe activity changes E by 0.059 V.When a reaction involves H+ or OH− ions, these must be included in the Nernst equation to predict the pHdependence of the couple. Thus for thehalf-cell reaction shown in Table 1. a factor of [H+]8 should beincluded in the [Ox] term, leading to a reduction in potential of (8/5)×0.059=0.094 V per unit increase in pH.
pHchanges may also have a more subtle influence by altering the species involved. For example, in alkaline solution the ionMn2+ precipitates as Mn(OH)2. Standard potentials at pH=14 refer to reactions written with OH− rather than H+ (seeTopic B4).Potentials may also be strongly influenced by complex formation. In general, any ligand that complexes morestrongly with the higher oxidation state will reduce the potential. For example, cyanide (CN−) complexes much morestrongly with Mn3+ than with Mn2+, and at unit activity reduces the Mn3+/Mn2+ potential from its standard value of+1.5 V to +0.22 V.
Conversely, the potential increases if the lower oxidation state is more strongly complexed.Diagrammatic representationsA Latimer diagram shows the standard electrode potentials associated with the different oxidation states of an element,as illustrated in Fig. 1 for manganese. Potentials not given explicitly can be calculated using Equation 1 and taking carefulaccount of the number of electrons involved. Thus the free energy change for the Mn3+/Mn reduction is the sum ofthose for Mn3+/Mn2+ and Mn2+/Mn.
From Equation 1 thereforeFig. 1. Latimer diagram for Mn at pH=0.In a Frost or oxidation state diagram (see Fig. 2) each oxidation state (n) is assigned a volt equivalent equal to ntimes itsvalue with respect to the element. The potentialin volts between any two oxidation states is equal tothe slope of the line between the points in this diagram. Steep positive slopes show strong oxidizing agents, steepnegative slopes strong reducing agents. Frost diagrams are convenient for displaying the comparative redox properties ofelements (see Topics F9 and H3).Frost diagrams also provide a visual guide to when disproportionation of a species is expected. For example, inFig. 2 the Mn3+ state at pH=0 is found above the line formed by joining Mn2+ with MnO2.
It follows that the Mn3+/Mn2+ potential is more positive than MnO2/Mn3+, and disproportionation is predicted:SECTION E—CHEMISTRY IN SOLUTION147Fig. 2. Frost diagram for Mn at pH=0 (solid line) and pH=14 (dashed line).The equilibrium constant of this reaction can be calculated by noting that it is made up from the half reactions forMnO2/Mn3+ and Mn3+/Mn2+ each with n=1, and hasfrom Fig. 1.
giving K=2×109. TheVVIstates Mn and Mn are similarly unstable to disproportionation at pH=0, whereas at pH=14, also shown in Fig. 2.only MnV will disproportionate.Latimer and Frost diagrams display the same information but in a different way. When interpreting electrodepotential data, either in numerical or graphical form, it is important to remember that a single potential in isolation hasno meaning,Kinetic limitationsElectrode potentials are thermodynamic quantities and show nothing about how fast a redox reaction can takeplace (see Topic B3). Simple electron transfer reactions (as in Mn3+/Mn2+) are expected to be rapid, but redoxreactions where covalent bonds are made or broken may be much slower (see Topics F9 and H7).
For example, thepotential is well above that for the oxidation of water (see O2/H2O in Table 1), but the predictedreaction happens very slowly and aqueous permanganate is commonly used as an oxidizing agent (although it shouldalways be standardized before use in volumetric analysis).Kinetic problems can also affect redox reactions at electrodes when covalent substances are involved.
For example, apractical hydrogen electrode uses specially prepared platinum with a high surface area to act as a catalyst for thedissociation of dihydrogen into atoms (see Topic J5). On other metals a high overpotential may be experienced, as acell potential considerably larger than the equilibrium value is necessary for a reaction to occur at an appreciable rate.Section F—Chemistry of nonmetalsF1INTRODUCTION TO NONMETALSKey NotesCovalent chemistryIonic chemistryAcid-base chemistryRedox chemistryRelated topicsHydrogen and boron stand out in their chemistry. In the other elements,valence states depend on the electron configuration and on the possibilityof octet expansion which occurs in period 3 onwards. Multiple bonds arecommon in period 2, but are often replaced by polymerized structureswith heavier elements.Simple anionic chemistry is limited to oxygen and the halogens, althoughpolyanions and polycations can be formed by many elements.Many halides and oxides are Lewis acids; compounds with lone-pairs areLewis bases.
Brønsted acidity is possible in hydrides and oxoacids. Halidecomplexes can also be formed by ion transfer.The oxidizing power of elements and their oxides increases with groupnumber. Vertical trends show an alternation in the stability of the highestoxidation state.Electronegativity and bondChemical periodicity (B2)type (B1)Electron pair bonds (C1)Covalent chemistryNonmetallic elements include hydrogen and the upper right-hand portion of the p block (see Topic B2, Fig. 1). Covalentbonding is characteristic of the elements, and of the compounds they form with other nonmetals. The bondingpossibilities depend on the electron configurations of the atoms (see Topics A4 and C1).
Hydrogen (Topic F2) isunique and normally can form only one covalent bond. Boron (Topic F3) is also unusual as compounds such as BF3have an incomplete octet. Electron deficiency leads to the formation of many unusual compounds, especiallyhydrides (see also Topic C7).The increasing number of valence electrons between groups 14 and 18 has two possible consequences. In simplemolecules obeying the octet rule the valency falls with group number (e.g. in CH4, NH3, H2O and HF, and in relatedcompounds where H is replaced by a halogen or an organic radical).
On the other hand, if the number of valenceelectrons involved in bonding is not limited, then a wider range of valencies becomes possible from group 15 onwards.This is most easily achieved in combination with the highly electronegative elements O and F, and the resultingcompounds are best classified by the oxidation state of the atom concerned (see Topic B4). Thus the maximumpossible oxidation state increases from +5 in group 15 to +8 in group 18.
The +5 state is found in all periods (e.g.PF5) but higher oxidation states in later groups require octet expansion and occur only from period 3 onwards (e.g.SF6 andin group 18 only xenon can do this, e.g. XeO4).150SECTION F—CHEMISTRY OF NONMETALSOctet expansion or hypervalence is often attributed to the involvement of d orbitals in the same principal quantumshell (e.g. 3d in period 3; see Topics A3 and A4).
Thus six octahedrally directed bonds as in SF6 could be formed withsp3d2 hybrid orbitals (see Topic C6). In a similar way the multiple bonding normally drawn in species such as(1)is often described as dπ-pπ bonding. These models certainly overestimate the contribution of d orbitals. It is alwayspossible to draw valence structures with no octet expansion provided that nonzero formal charges are allowed. Forexample, the orthonitrate ionis drawn without double bonds (2), andcould be similarly represented.One of many equivalent valence structures for SF6 where sulfur has only eight valence-shell electrons is shown in 3.Three-center four-electron bonding models express similar ideas (see Topic C6).
Such models are also oversimplified.It is generally believed that d orbitals do play some role in octet expansion, but that two other factors are at least asimportant: the larger size of elements in lower periods, which allows higher coordination numbers, and their lowerelectronegativity, which accommodates positive formal charge more easily.Another very important distinction between period 2 elements and others is the ready formation of multiple bondsby C, N and O (see Topic C8).
Many of the compounds of these elements have stoichiometries and structures not repeatedin lower periods (e.g. oxides of nitrogen; see Topic F5).Some of these trends are exemplified by the selection of molecules and complex ions in Table 1. They have beenclassified by (i) the total number of valence electrons (VE), and (ii) the steric number of the central atom (SN), whichis calculated by adding the number of lone-pairs to the number of bonded atoms and used for interpreting moleculargeometries in the VSEPR model (see Topic C2).
![](https://s.studizba.com/z.php?f=/templates/si/i/noavatar.png&w=100&h=100&t=1"){kind=avatar}
