P.A. Cox - Inorganic chemistry (793955), страница 48
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Later elements in the 4d and 5d series arerelatively more inert.Neutral atoms have both s and d valence electrons, but in chemicallyimportant states are often regarded as having purely dn configurations.Many-electron atoms (A3)IntroductiontononThe periodic table (A4)transition metals (G1)Trends in atomic properties(A5)ScopeTransition metals are elements of the d block that form compounds where electrons from d orbitals are ionized orotherwise involved in bonding.
Typical transition metal characteristics include: the possibility of variable oxidationstates; compounds with spectroscopic, magnetic or structural features resulting from partially occupied d orbitals; anextensive range of complexes and organometallic compounds including ones with very low oxidation state (zero or evennegative); and useful catalytic properties shown by metals and by solid or molecular compounds.
Different transitionmetals display these features to different degrees, but together the properties form a sufficiently coherent pattern thatthe elements are best dealt with by themed Topics rather than individually or as groups.Although formally part of the d block, the elements of group 12 do not show typical transition metal characteristics,as the d orbitals are too tightly bound to be involved in chemical bonding. These elements are better regarded as posttransition metals, and are dealt with in Section G (Topic G4).210SECTION H—CHEMISTRY OF TRANSITION METALSVertical trendsThe smaller size of the 3d orbitals compared with 4d and 5d has some important consequences.• Electron repulsion is large between electrons in 3d orbitals. Exchange energy effects are more significant (seeTopic A3); also, successive ionization energies (IEs) rise more sharply compared with later series (see Topic A5).• 3d orbitals are not much larger than the 3p orbitals of the argon core (3p)6.
Good overlap with other atoms is hard toachieve, and covalent or metallic bonding involving 3d orbitals is weak compared with 4d and 5d.One consequence of the IE trend is that higher oxidation states are less stable (more strongly oxidizing) compared withthe 4d and 5d series. For example, in group 7is much more strongly oxidizing thanand in group 8 FeO4is unknown although RuO4 and OsO4 are stable compounds.The bond-strength trend 3d<<4d<5d is the reverse of that normally found in main groups (see Topic C8). Itsinfluence can be seen in the atomization enthalpies of the elements, reflecting the strength of bonding in the metallicstate, and shown in Fig.
1 for elements of the three series. The very high atomization energies of elements such astungsten (5d group 6) are reflected in their extremely high melting and boiling points, a property important inapplications such as electric light bulb filaments. Sublimation energies in the middle of the 3d series are much less,partly because the relatively poor overlap of 3d orbitals gives weaker bonding, and partly because of the exchangeenergy stabilization of the free atoms, which have several unpaired electrons (e.g.
six with Cr). Compounds withunpaired electrons in d orbitals are also much commoner in the 3d series, those of the 4d series more often forming lowspin configurations or having d electrons involved in metal-metal bonds (see Topics H2 and H5).Between the 4d and 5d series the expected decrease of IEs and increase of radius is counteracted by the increase ofnuclear charge involved in filling the 4f shell before 5d (see Topic A4). 5d elements in early groups are very similar tothe corresponding 4d ones, although this feature is less marked in later groups.Horizontal trendsThe chemical trends along each series are dominated by the increase in nuclear charge and in the number of valenceelectrons. Earlier elements can achieve the group oxidation state corresponding formally to ions with a noble gasconfiguration (up to MnVII in the 3d series and RuVIII and OsVIII in 4d and 5d).
Increasing effective nuclear charge bringsan increase in IEs as shown for the 3d elements in Fig. 2. Not only does the group oxidation state become very stronglyoxidizing for later elements, but redox potentials for any given states (e.g. M3+/M2+) also increase along the series, asthe extra lattice or solvation energies of the higher state become less able to compensate for the higher IE values (seeTopics D6 and G1).With increasing IEs comes also a general decline in electropositive character.
Early elements in each series arethermodynamically extremely reactive towards oxygen and other electronegative elements (although the formation ofan inert oxide film may kinetically prevent the solid elements from further oxidation). Later elements are less reactive,a trend that culminates in the ‘noble’ or ‘coinage’ metals Cu, Ag and Au of group 11. The trend is exacerbated in thelater 4d and 5d elements by the high atomization energies, and the elements Ru, Rh, Pd, Os, Ir and Pt form a groupknown as the platinum metals, often occurring together in nature, sometimes as metallic alloys.
The change inelectronegativity is also shown by different patterns of chemical stability: whereas earlier elements of both seriesgenerally form more stable compounds with ‘harder’ anions such as oxide and fluoride (and are found in nature in oxideminerals), the later ones are ‘softer’ in character and are more often found as sulfides. The trend along the series thusprovides a link between the chemical characteristics of the pre-transition and post-transition metals (see Topic G1).H1—INTRODUCTION TO TRANSITION METALS211Fig. 1. Standard enthalpies of atomization for elements of the three series.A general decline in atomic size is another consequence of increasing effective nuclear charge.
Figure 2 also shows theionic radii of M2+ ions of the 3d series. The expected decrease across the series is modulated by ligand field effects(see Topic H2).Electron configurationsElectron configurations of the neutral atoms are complex and have both d and s electrons in outer shells. For example,in the 3d series most atoms have the configuration (3d)n(4s)2, where n increases from one to 10; chromium and copperare, however, exceptions with (3d)5(4s)1 and (3d)10(4s)1, respectively.
The configurations depend on a balance of twofactors:(i) 3d orbitals are progressively stabilized relative to 4s across the series;(ii) repulsion between electrons is large in the small 3d orbitals, and so minimum energy in the neutral atom isachieved in spite of (i) by putting one or two electrons in the 4s orbitals.An important consequence of this balance is that in forming positive ions, 4s electrons are always removed first. Thusfor M2+ ions and ones of higher charge, outer-shell electrons are left only in the 3d orbitals.
Figure 2 lists the value of nin the configuration (3d)n for M2+; values for higher charges may be found from these by subtracting the appropriatenumber of electrons (e.g. Ti4+ (3d)0 and Fe3+ (3d)5). These numbers can be used to interpret the IE trends shown:whereas I1 and I2 rise fairly steadily (with small irregularities resulting from the exceptional configurations of Cr andCu), the I3 plot shows a pronounced break after manganese. With six or more electrons in the d shell some must pairup, thus giving greater electron repulsion and a lower IE than expected from the previous trend (see Topic A5).Ligand field theory deals with the important consequences of the progressive filling of the d shell.
It is normal tospecify the d electron number associated with the appropriate transition metal ion, even though the bonding is notassumed to be completely ionic. For example, any FeIII compound is assigned the configuration (3d)5, a PtII compound(5d)8 (corresponding to Ni2+ in the same group). In compounds with very low oxidation states, or with ligands such as212SECTION H—CHEMISTRY OF TRANSITION METALSFig. 2. Data for ions of the elements Ca-Zn showing: radii of M2+ ions, third IE and the sum of first and second IEs, and the (3d)n configurations ofM2+.organic groups where bonding is largely covalent, a different electron counting scheme is often used (see Topics H9 andH10).
In applying the 18-electron rule one needs to count the total number of valence electrons in a neutral atom,irrespective of whether they are d or s. This is simply the group number, thus eight for Fe and 10 for Pt. If ligand fieldarguments are used for very low oxidation states the electrons in the appropriate ion are assigned entirely to d orbitals.For example, a CoI compound would be regarded as (3d)8 even though the free Co+ ion has the configuration (3d)7(4s)1. The justification for this procedure is that the energy balance between d and s orbitals changes on compoundformation; what were s orbitals in the free ion become strongly antibonding molecular orbitals in a complex and are nolonger occupied in the ground state.Section H—Chemistry of transition metalsH2LIGAND FIELD THEORYKey NotesOctahedral splittingsHigh and low spinLigand fieldstabilization energyOther geometriesRelated topicsLigand field splitting of the d orbitals arises from a combination of σand π bonding interactions with ligands.
In octahedral (Oh) geometrytwo orbitals (eg) are at higher energy than the other three (t2g). Thespectrochemical series puts ligands in order of field strength. High-fieldligands are strong σ donors and π acceptors.High-spin complexes have as many d electrons unpaired as possible,and are common with 3d series elements. Low-spin complexes haveas many electrons as possible in the lower set of orbitals, and arecommon in the 4d and 5d series.Ligand field stabilization energy (LFSE) is calculated relative to theaverage of all d orbital energies.
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