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Ternary and quasiternary compounds consist of either one cationic and two anionic, or one
anionic and two cationic components. The size of the ions and the bonding forces between
them determine the structure of the compounds. A compound is regarded as a complex when
its lattice contains structural groups in which the interaction between the constituent atoms
is substantially greater than or dtferent from that between these structural groups and
the other components of the compound. Metallo-complexes are built up of one cationic
and two anionic components; in contrast to the complexes with cationic central atoms
which have been thoroughly investigated, they contain an anionic central atom surrounded
by cationic ligands. A list is given of many compounds which can be formulated as metallocomplexes after discussion of methods for their preparation and their chemicalproperties;
the reason that only few metals can funciion as ligands of an anionic central atom i r
deduced from the pertinent crystal structure.>.
A. Introduction
In modern chemistry, a salt is defined as a substance
whose components are arranged in an i o n i c lattice [2].
If we consider a salt K,Ap in which the lattice sites are
occupied by cations K and anions A, the type of lattice
assumed is determined by the sizes of the ions involved,
the bonding forces between them, and the principle of
electrical neutrality. In ternary compounds
or K,A(l)pA(Z)q,
either the K or the A sites are occupied by two different
elements. The magnitudes of the bonding forces K(1)-A
and K(2)-A or K-A(1) and K-A(2) are normally different, and hence play a significant part in determining the
crystal structure. Typical mixed crystals are formed
when the bonding forces are approximately equal in
magnitude, and typical complex salts result when the
bonding forces are vastly different.
Although complex salts of the form K(1),[K(2),Ap]
have been investigated in great detail ever since the
studies of Werner, little attention has been given to
ternary compounds of the type [K,A( 1),]A(2),.
The names of such complexes can be formed in various ways
according to the I.U.P.A.C. Rules for Nomenclature of
Inorganic Chemistry [3]:
1. The addition of protons to mono-atomic anions yields
cations of the type [XH,]+, which are named by adding
-onium to the Latin name of the anionic element (Rule 3.14).
I f the protons are displaced by other cations (a possibility
not considered in the I.U.P.A.C. Rules), the names remain
unchanged (Rule 7.321) and terms such as trisilversulfoniurn
nitrate for [SAg3] [NO31 are obtained.
2. If substitutive nomenclature is to be avoided in inorganic
chemistry, the name of the complex [SAg3]+ can be formed
by listing the partners. In this case neither the cationic ligand
[I] Excerpts from the Habilitation Thesis of G. Bergerhoff, Universitat Bonn, 1962 (D 5).
[2] See, for example, G. Schwnrzenbach: Allgerneine und anorganische Chemie. S. Hjrzel, Stuttgart 1948, p. 76.
[3] I. U. P. A. C . Nomenclature of Inorganic Chemistry 1957.
Butterworths, London 1959.
silver (Rule 7.321) nor the central atom sulfur (Rule 7.24)
are given a distinguishing termination. The name of [SAg3]
[NO31 would then be trisilversulfur nitrate.
3 . Rule 7.321, which requires that the name of a cationic
ligand be used without change, does not list any pertinent
example. This is to be regretted, since otherwise it would
have become clear that it seems advisable to denote the
position of a positive ligand within a complex also by using
the ending -0.
[SAg3] [NO31 may thus be referred to as trisilversulfonium
nitrate, trisilversulfur nitrate, triargentosulfonium nitrate,
or triargentosulfur nitrate. In this article, the last system of
nomenclature will be used, in accordance with Lieser [4].
Complexes with an anionic central atom and cationic ligands
may be conveniently designated general as metallo-complexes.
B. A Systematic Classification of
Multicomponent Compounds
1. Ternary Compounds
In order to arrive at an unambiguous definition of
metallo-complexes, let us arrange the structures of the
ternary compounds K( 1),K(2),AP and K,A( 1)pA(2)q
into a series, with the difference in bond strengths of
K(l)-A and K(2)-A or of K-A(l) and K-A(2) as parameter. Four groups are formed.
Group I
Solid Solutions and Double Salts with
Binary-Salt Type Structures
If the sizes and polarizability of the elements K(l) and
K(2) at the K-sites or A(1) and A(2) at the A-sites are
very similar, then K(1) and K(2) or A(1) and A(2) are
mutually i n t e r c h a n g e a b l e . They are randomly distributed, and the relative proportions of the elements
need not even be stoichiometric, so that mixed crystals
can occur as well as double salts (Table 1).
[4] K. H. Lieser, Z . anorg. allg. Chem. 295, 106 (1958).
Angew. Chem. internat. Edit. 1 Vol. 3 (1964)
No. I0
Group 2 - Double Salts with Structures Derived from
Binary-Salt Structures by Ordered Distribution of the
Two Types of Cations or Anions
A greater difference in bond strengths forces the elements to assume an ordered distribution, resulting in a
new type of structure. However, the essential structural
relationships of the binary salts are retained, and consequently the coordination numbers of K(l) and K(2) or
A(l) and A(2) are practically equal (Table 1).
their components, they are called complex compounds”.
This definition is unsatisfactory nowadays, for we know
that on principle complexes dissociate into their components in solution, although there are great differences
among the complex formation constants and the rates
of dissociation. Moreover, dissociation of complexes is
regarded as a chemical reaction, the course of which
depends on the nature of the reaction partners. A true
complex compound may even be present when its
components can be detected in solution. It seems
useful to abandon this “dynamic” definition, and as
Table 1. Examples of the classification of ternary compounds [a].
a : K(l)mK(2)nAp
I Structure [bl
1 1
NaCl type,
mixed crystals
Fluorite type,
double salt
Related to
zinc blende
Perovskite type
C N of K(1):IZ
C N of K(2) :6
Spinel type
C N of K(1):4
C N of K(2): 6
b: KmA(l)pA(z)q
Related to burcite
C N of A(I):6
C N of A(2): 12
Layer lattice
C N of A(1):4
C N of A(2): 5
I Two-shell complex
11I 1
Chain complex
[a] Although Cud0H)aBrZ and Hg(NH2)Br are quaternary compounds,
they are included here since the proton practically disappears in the
electron cloud of its bonding partner.
[b] C N = Coordination number.
Group 3 - Lattice Compounds in which the Two Types
of Cations or Anions have Different Properties
If the difference in bond strengths is even greater, then
K(l) and K(2) or A(l) and A(2) assume different
coordination numbers, although the nature of the bonds
between the cations and the anion or between the
anions and the cation remains similar. The resulting
structures can no longer be traced back to the binarysalt type (Table 1).
Group 4 -- Complex Compounds
The series is terminated by compounds which were
defined by Werner [I21 in 1913 as follows: “If compounds of higher order are characterized by stability in
aqueous solution, in that they do not dissociate into
[ 5 ] W . H. Zachariasen, Acta crystallogr. 2, 388 (1949).
[6] W. Finkelnburg and A . Srein, J. chern. Physics 18, 1296 (1950).
[7] L. Pauling and L. 0. Brockway, Z. Kristallogr., Mineral., Petrogr. 82, 188 (1932).
[S] F. Aebi, Helv. chirn. Acta 31, 369 (1948).
[9] A . F. Weffs: Structural Inorganic Chemistry. University
Press, Oxford 1962, pp. 360, 487.
[ l o ] B. Reuter and K. Hardel, Angew. Chem. 72, 138 (1960).
11 11 W. Nieuwenkamp and J. M . Bijvoef, Z. Kristallogr., Mineral., Petrogr. 81, 469 (1931).
[I21 A . Werner: Neuere Anschauungen auf dem Gebiete der anorganischen Chemie. Friedr.Vieweg, Braunschweig 1913, p. 28.
Anyew. Clwm. internal. Edit.
NaCl type,
mixed crystals
Fluorite type,
double salt
Island complex
Chain complex
Layer complex
I Structure [bl
1 Vol. 3 (1964) 1 No. 10
with the concept of a “salt”, to give a “static” definition
based on the structure of the solid state:
Complex compounds are compounds whose lattices contain structural groups in which the interaction between
the constituent atoms is substantially greater than or
different from that between these structural groups and
the other components 0.f the compound.
Of course, the interaction between the atoms of a
crystal lattice cannot be determined readily by experiment; however, assignments can be made in
most cases from considerations based on bonding
theory. For example, it may be plausible, from
the electron configuration of the atoms or from comparison of the interatomic distances, to assume that the
bonding for a given compound is homopolar. Sometimes it can be determined from the size and charge
of the ions, and from their polarizability and polarizing
action, whether the compound in question is a complex.
Structural groups consisting of discrete units made up
of a central atom surrounded by ligands are usually
known as “island complexes”. If these ligands in turn
act as central atoms and become surrounded by a
second outer ligand sphere, which is also to be regarded
as part of the structural group, the compounds are
referred to as t w o - s h e l l or s u p e r c o m p l e x e s [13,14].
[13] F. ffein: Chemische Koordinationslehre. S . Hirzel, Leipzig
1950, p. 217.
[ 141 H . Renzy: Lehrbuch der anorgdnlschen Chemic. Akademische
Verlagsgesellschaft Geest u. Portig, Leipzig 1959, Vol. 11, p. 364.
However, the ligands of a complex can also link up
several central atoms to form polynuclear complexes,
and these logically include complexes which extend in
one or more directions to infinity (polymeric complexes)
[15]. It appears to be convenient to refer to the atom
with the g r e a t e r coordination number as the central
atom; it is then possible to differentiate between
complexes with cationic and anionic central atoms.
Some examples are shown in Table 1.
2. Quasiternary Compounds
The above division into four groups provides a basis
for a systematic classification of inorganic complex
structures; which largely reflects the relationships between the various lattice types. Groups 1 and 2 have
been characterized already by the fact that their structures can be traced back to those of binary compounds.
If lattice sites are occupied by complexes (in the sense
of discrete structural groups), ternary compounds, such
as K2[S04], can be regarded as quasibinary compounds,
and quaternary or higher-order compounds, such as
[K(&0)6] [Al(HzO)6] [so4]2, as quasiternary.
The series can now be extended to include quasiternary
compounds. Some examples are shown in Table 2.
distinguished between “the great majority of coordination compounds, in which the centers of complex
formation are elementary atoms which are more electropositive than the elementary atoms which accept the
linkages of the added molecular components”, and a
“small group of compounds in which the opposite is the
case”. By way of example, Werner [28] quoted the compound [IAg3][NO&, in which Hellwig [29] had observed migration of a complex ion [lAgj]Zi- during
electrolysis. This salt and similar silver salts are still the
classical examples of this type of complex, although
the number of compounds which can be formulated as
metallo-complexes has steadily grown. EvidentJy the
lack of a good example of a typical island complex was
the reason for alternative formulations. Thus. Hawk
[30] regarded whole “salt molecules” as ligands, e.g.
in [Hg(HgI2)21[C104]2, because he required cationic
central atoms as in conventional complexes. Yatsimiw
kii [31] listed numerous compounds which can be
described as complexes with anionic central atoms, but
did not prove his assumptions. It is only recently that
modern physical methods have been applied to silver
compounds and that mercury compounds in particular
have been investigated by X-ray methods. As prerequisites for the formation of metallo-complexes, Lieser
[32] named high polarizability of the negative central
Table 2. Examples of the application of the systematic classification t o quasiternary compounds.
Related t o brucite
Related to brucite
Apatite type
C . Metallo-Complexes
1 , General Discussion and Examples
The examples of Group 4b in Tables I and 2 will be
referred to as metallo-complexes. Werner [28], ii1
the systematic application of his coordination theory,
[I51 See [13], pp. 180, 200, 290, and [2], p. 101.
[I61 See [12], p. 305.
[I71 See references in Table 3.
1181 C.Brink and C. H. MacGillavry, Actacrystallogr. 2, 158 (1949).
[IS] D. Balz and K . Plieth, 2. Elektrochem., Ber. Bunsenges.
physik. Chem. 59, 545 (1955).
[20] H . G. v. Schnering, R . Hoppe, and J. Zemann, Z. anorg. allg.
Chem. 305, 241 (1960).
1211 L. Gerhardr, Naturwissenschaften 33, 56 (1946).
[22] W. Feitknech/, Fortschr. chem. Forsch. 2, 702 (1953).
[23] Landolt-Bornstein: Zahlenwerte und Funktionen aus Physik, Chemie, Astxonomie, Geophysik und Technik. Springer-Verlag, Berlin-Gottingen-Heidelberg 1950, Vol. I, Part 4, p. 58.
[24] W. Nuwacki and R. Scheidegger, Helv. chim. Acta 35, 375
[25] A . Zemann and J . Zeinann, Acta crystallogr. 14, 835 (1961).
[261 H . Wondrafschek, Neues Jb. Mineralog., Abh. 99, 113(1963).
[271 A . Ferrari, L . Cavalca, and M . Nardelli, Gazz. chim. ital. 81,
945 (1951).
[28] See [12], p. 300.
atom, and Yatsimirskii its low electron affinity. Obviously not all metals are capable of forming complexes
of this type; according to Yatsimirskii, a high ionization
potential is necessary. Frequently, the difficulties encountered in preparing compounds of this type indicate
that their free enthalpy of formation is not very great,
so that even small effects can reverse its sign. The extracomplex ion A(2) is also of importance, since it can
change the lattice energy, and can even act as a complexing ligand for a cationic central atom.
Table 3 shows some compounds which can be formulated as metallo-complexes. Their structures are
known only in a few cases. The Table, which is not
meant to be complete, shows how niany structures
might be ordered using the concept of “metallo-complexes”. However, only an elucidation of these structures will finally prove the value of this list.
[29] K. HeUwig, Z . anorg. allg. Chem. 25, 157 (1900).
[30] E. Hayek, 2. anorg. allg. Chem. 223, 382 (1935).
[31] K. B. Yutsimirskii, Dokl. Akad. Nauk S.S.S.R.77,819(1951).
[32] K. H. Lieser, 2. anorg. allg. Chem. 292, 114 (1957).
[33] A . Weissand A . Weiss, 2. anorg. allg. Chern. 282, 324 (1955).
[34] W. Riidorj’and K.Brodersen,Z.anorg. allg.Chem. ?87,24(1956).
[35] R . Airddi, Ann. Chimica 48, 491 (1958).
[36] W . Riidorj’and K . Brodersen, 2. Naturforsch. 9b, 164 (1954).
1371 W. Riidorfand K. Brodrrsen, Z . anorg. allg. Chem. 270, 145
A n g e w . Chem. internat. Edit.
Vul. 3 (1964)
1 No. I0
Table 3. Compounds xhich can be formulated as metallo-complexe\
Central atom
Stoichiometry 01
the complex ion
Known counter-ion
Structure of the
complex i o i i
C1k. OH-. SO:-, MnO,
and others
Two-shell complex
Two-shell complex
T‘wo-shell complex
Two-shell complex
Frame work
[ * ] This group includes a few “basic salts”; most of these form structure of Groups I , 2, and 3
Angew. Chem. internat. Edit. 1 Vol. 3 (1964) 1 No. 10
2. Preparation of Metallo-Complexes
Metallo-coniplcxes can be convenicntly prcparcd by three
incthoda :
P KA(1)
q KA(2)
(P .I. m) KA(1)
+ A(I),,Xq
+ m) KA(2) + P A(I)X lA(lP)Kq+ml-W)q + A(2),XP
Metallo-complexes are formed according to Equation (a)
either in melts 1711, e . g . [IAg,] [N03]2, or on annealing by
reaction in the solid state, e . g . [S2Hg3]C12. This stage may be
preceded by another reaction, such as a redox reaction in the
preparation of [SAg,] [SO,] from silver nitrate and sulfur.
The salts KA(1) are often only sparingly soluble in water, but
dissolve to a n appreciable extent in a sufficient excess of
KA(2) solution. Thus, the salts AgI, AgCNS, AgZTe, etc. are
soluble in concentrated aqueous solutions of silver nitrate,
complexes being precipitated on careful dilution. Mercury(J1)
[38] A. Weiss and E. Michel, Z. Naturforsch. ISb, 679 (1960).
[39] Th. Poleck and K. Thiimmel, Ber. dtsch. chem. Ges. 16, 2435
[40] H . Puff; Angew. Chem. 74, 659, 507 (1962); Angew. Chem.
internat. Edit. I , 41 I (1962); H. PuR, Naturwissenschaften 49,
299, 464 (1962).
[41] A . Weiss and G . Hofmann, Z . Naturforsch. 15b, 679 (1960):
S . Scavnicar, Z. Kristallogr., Mineral., Petrogr. 118, 248 (1963).
[42] A. Weiss, G. Nagorsen, and A . Weiss, Z . anorg. allg. Chem.
274, 151 (1953); D . Grdenic and S . Scavincar, Acta crystallogr. 8,
275 (1955): K . Aurivittius, ibid. 16, A 25 (1963).
[43] D . Grdenic and F. Zado, J. chem. SOC.(London) 1962, 521.
1441 A . Weiss, S . Lyng, and A. Weiss,2. Naturforsch. 156,678 (1960).
[45] A. Weiss, S. Nagorsen, S . Lyng, and A . Weiss, Angew. Chem.
74, 119 (1962); Angew. Chem. internat. Edit. 1, 115 (1962).
[46] E. Hayek, Z. anorg. allg. Chem. 223, 382 (1935).
[47] K. Aurivillius, Acta chem. scand. 8, 523 (1954).
[48] A . Tulinsky, C. R . Worthington, and E. Pignataro, Acta crystallogr. 12, 623 (1959).
[49] C. C . Addison and A.Watker, Proc. chem. SOC.(London)
1961, 242.
[50] G . M . Kurdyumov and K . N . Semenenko, Chem. Abstr. 52,
6995e (1958).
[51] If. D . Hardt, Z. anorg. allg. Chem. 293, 47 (1957).
[52] A . Gupta, Sci. and Cult. (Calcutta) 25, 426 (1960); Chem.
Zbl. 132, 13467 (1961).
[53] H . D . Hardt and F. Stavenow, Z. anorg. allg. Chem. 301,267
[54] A . G. Ostroffand R . T. Sanderson, J. inorg. nuclear Chem. 9,
45 (1959).
[55] G. Denk, Chem. Ber. 92, 2236 (1959).
[56] W. P. Binne, Acta crystal1ogr.t 4, 471 (1951).
I571 H. Bode and E. Voss, Electrochimica Acta I , 318 (1959).
[ 5 8 ] L . E. Orgel, Nature (London) 187, 504 (1960); H. D. Hardt
and W. Moller, Z. anorg. allg. Chem. 313, 57 (1961); K. Starke,
J. inorg. nuclear Chem. 13, 254 (1960).
[59] G. Bergerhoff, Z. anorg. allg. Chem. 299, 328 (1959).
[60] H . Stamm and H . Wintzer, Ber. dtsch. chem. Ges. 71, 2212
[61] J. K . Bouknight and G. McP. Smith, J. Amer. chem. SOC.61,
28 (1939).
[62] G. McP. Smith and W . L. Sexmon, J. Amer. chem. SOC.46,
1325 (1924).
[63] A . Bernard; and G. Rossi, Gazz. chim. ital. 52, 140 (1922),
[64] K T a k a i , Bull. chem. SOC.Japan 28, 403 (1955).
[65] A. Baroni, Atti Accad. naz. Lincei, Rend. [6] 22, 76 (1939).
[66] A . Naumann, Ber. dtsch. chem. Ges. 43, 313 (1910).
[67] E. Blasius and K . Riebnrfsch, Z. analyt. Chem. 176, 336
1681 B. Reuter and C. G . Beuthe, Angew. Chem. 72, 594 (1960):
B. Reuter and H. W . Levi, Z. anorg. allg. Chem. 291, 254 (1957).
[69] G . Bergerhoff and E. Schultze-Rhonhof, unpublished work.
[70] K. H . Lieser, Z. anorg. allg. Chem. 305, 255 (1960).
1711 See [23], Vol. 2, Part 3, p. 99.
iodide is soluble in hot aqueous mcrcury(I1) perchlorate, and
14Hg31 [C10,]2 crystallizes out on cooling [46].
Addition of A(2) according to Equation (b) leads to the foriiiaiion of {[OBe4] [CH3C00Ib} from beryllium hydroxideand acetic anhydride since excess hydroxide ions arc taken
tip by the acetic anhydride to form acetic acid.
Very many metallo-complexes are formed according to
Equation (c). Thus, at the beginning of the precipitation
with hydrogen sulfide of many heavy metals, such as thallium,
lead, or mercury, brightly colored “sulfobasic” salts appear,
e.g. [ST13]CI. “Basic salts”, which may also be metallocomplexes, are generally formed when the p H is increased,
i . e . on addition of OH- or 0 2 - ions. The formation of
complexes with a central nitrogen atom is so far known only
for mercury, with ammonia as the source of A(1); in this
reaction the driving force is probably the insolubility of the
“salts of Millon’s base”. Phosphine can give phosphorus
complexes with silver or mercury compounds.
However, these methods of preparation usually yield the
compounds at best only in microcrystalline form. Larger
crystals can be obtained if A(l) is slowly formed by chemical
reaction. Of the series methyl iodide, isopropyl iodide, and
2,4-dinitroiodobenzene, only the middle member forms large
crystals of triargentoiodine nitrate with concentrated silver
nitrate solution:
+ H20 + 3 AgNO3
IIAg31 “Oilz
+ HNO3 +
Carbon disulfide can be used as the source of sulfur for the
preparation of large crystals of triargentosulfur nitrate :
+ 2 HzO + 6 AgNO3 -+ 2 [SAgxl “0x1 + 4 HN03 + COz.
The crystal size can be further increased if the compounds
are formed in silica gel, for in this colloidal capillary system,
the crystal seed is held fast and can grow by influx of material
from a11 sides by diffusion. For example, crystals of [SAg3]
[NO31 up to 1 mm long have been obtained as well-defined
The capillary structure of silica gel containing silver nitrate
can be very elegantly demonstrated by covering the gel with
a layer of reducing agent; this procedure results in growth
of silver whiskers up to 1 cm long in which, according to
rotating crystal photographs, the filament direction is always [110].
3. Structure of Metallo-Complexes
The structures of only a few metallo-complexes are
known at present. All of them represent well-known
complex types.
Thus,trischloromercury(I)-oxygen chloride[O(HgCl)~]Cl
is a two-shell complex or supercomplex (Fig. 1).
In the first sphere, the central oxygen atom occupies the
apex of a trigonal pyramid [77], the base of which is formed by three mercury atoms at distances of 2.06 A (sum
of the covalent radii = 2.10 A apart); these in turn are
linked to chlorine at a distance of 2.29 A (Hg-C1 distance
in HgC12 = 2.34 A), forming a second sphere of covalently
bonded ligands around the central oxygen atom. In
contrast, the chloride ions at a distance of 2.99 and
3.09A from the mercury atoms should be regarded as
_ _
[72] K . H. Lieser, Z . anorg. allg. Chem. 305, 133 (1960).
[73] E. Hayek, Mh. Chem. 68, 29 (1936).
1741 H . Morse, Z. physik. Chem. 41, 709 (1902).
[75] I. Lindqvist, Acta crystallogr. 7, 636 (1954).
[76] A . M. Golub and W . W . Skopenko, Dokl. Akad. Nauk
U.S.S.R. 138, 601 (1961); Chem. Zbl. 134, 6151 (1963).
[77] K . Aurivillius, Acta crystallogr. 16, A 25 (1963).
Angew. Chem. internat. Edit.
Vol. 3 (1964) 1 No. 10
lying outside the complex. Another example is “basic
beryllium acetate” {[OBe4][CH3C00]6}. In this complex,
the central oxygen atom is tetrahedrally surrounded by
four beryllium atoms; the carhoxyl groups link the
beryllium atoms in pairs, in the manner of chelates,
thus forming an inner complex salt. The beryllium
atoms, like the oxygen atom, therefore have a coordination number of four, so that the compound might
also be regarded as a tetranuclear complex of beryllium; however, oxygen certainly is the central atom.
Fig. 3. Section of three layers of (02Hg3) (S04). After [45]. The upper
layer has been continued backwards. Hatched circles = mercury;
open circles = oxygen; terahedra = sulfate.
Fig. 1 . Projection of the [O(HgCl)] [CII structure in the direction of
one of its threefold axes (anions omitted) [77].
“Basic mercury chlorate” or hydridomercury(1)-oxygen
chlorate [OHHg] [C103] forms a chain complex (Fig. 2),
in which the oxygen and the mercury atoms alternate.
argentosulfur nitrate [SAg3][N03]. The mercury(1)-complexes have a tridymite structure in which the tetrahedrally coordinated silicon atom is replaced by nitrogen
or phosphorus. and in which mercury takes the place
of the 2-coordinate oxygen atom. The oppositely charged
ions are situated in the spaces between the tetrahedra.
Similarly, [SAg3][N03], which is cubic, contains approximately trigonal prismatic SAg, structural groups [78],
linked through common corners to form a three-dimensional framework (Fig. 4). The cavities are sufficiently
large to accommodate the nitrate groups which are
ionical ly bonded.
Fig. 2. Chain structure of tbe [HOHp] (ClOal complex (anions) omitted.
However, the oxygen atom has a coordination number of
three, owing to the additional bond to the hydrogen,
while that of the mercury atom is only two, since the
chlorate ions lie at the ionic bond distance. “Mercuryamido bromide” [(NH2)Hg]Br has a similar structure.
The lattice is traversed by chains of alternating mercury
and nitrogen atoms. The coordination number of the
nitrogen atom is four, so that it must be regarded as the
central atom.
The class of layer complexes is represented by the
mercury(1)-oxygen compound [02Hg3][SO,], and also
by “mercuryimido bromide” [(NH)zHgs] LHgBr31Br.
According to preliminary structural determinations,
mercury atoms and oxygen or nitrogen atoms form
networks, the meshes of which receive the oppositely
charged ions (Fig. 3).
The oxygen or nitrogen atom has a coordination number of three and is the central atom of the complex. The
sulfate groups in trimercury(1)-dioxygen sulfate are
bonded ionically (Hg-0 distance in the complex =
2.03 A; Hg-OSO3 distance = 2.47 A).
Examples of framework complexes are dimercury(1)nitrogen salts such as [NHgzlBr, dimercury(1)-phosphorus chloromercurate(I1) [PHgz] [HgCL], and triAngew. Chem. internat. Edit.
Val. 3 (1964) f No. 10
Fig. 4. Three-dimensional structure of the SAg NO, complex showing
the SAg6 polyhedra and the nitrate ions ( 0 = N; o = 0).
[78] The structure published earlier 1591 has meanwhile been
more accurately calculated by the method of least squares, using
an IBM 704 computer. The program made it possible to presuppose the planar arrangement of the nitrate groups and the
N-0 distance of 1.218 A 1791, and used the difference Fobs-Fcdo
for the weighting, with w = l/(Fobs-Fcalc)Z
0.5 [go]. A
reliability factor of R = 14% (only with F
0) was obtained.
The following parameters were finally obtained:
Ag: 12b;x=O,308+ 2 y = 0.9942%3 Z = 0.3582f 2 B = 2.821. 5
S: 4a; x = 0.1367 k 8
B=2.3 i t 1
N : 4a; x = 0.4145 1
0: 12b ; x = 0.532
y = 0.318
z = 0.391 ( B = 1‘2
69 I
4. Requirements for the Formation of MetalloComplexes
finition, be lower than that of the central atom, and that
for steric reasons, ligands belong to more than two
coordination polyhedra only in special cases.
Examination of all the nietallo-complex structures
described shows that island and p3lynuclear complexes
are not encountered. Both types have a comm,m factor,
namely that their structural units must contain terminal
metal atoms; a hypothetical island complex [SAg3]and a dinuclear complex [Ag2S.4g2SAg# -. would contain metal atoms with a coordination number of I .
The tendency towards formation or structures o f this
type is obviously small.
These reyuircinents are satisfied for example by elements
of the B groups of the periodic table. i n the case of
ions with a dlo-electron configuration this behavior can
be attributed to the development of a ds-hybrid, for
which the excitation energy is relatively small, particularly for the Cu+, Ag+, Au+, and Hgz+ ions[81]. The
crystal field theory shows that an octahedral environment is then so strongly distorted that two strong
linear bonds remain.
It is known that the orbitals of cationic central atoms
undergo hybridization when ligands are added, so that
the resulting coordination polyhedron is more or less
symmetrical. Anionic ligands, on the other hand, can be
bonded by a singly occupied atomic-p-orbital, i. e. onesidedly. In the sulfate ion, for example, if double-band
contributions are neglected, the electron structure msy
be pictured as follows: sulfur, with the ground state
s2p4, gives up two electrons, and the remaining sIp-3electrons occupy four new tetrahedrally arranged equivalent orbitals. These hybridized orbitals overlap with
the singly occupied atomic p-orbitals of four oxygen
ions in the configuration s2p5, which are formed by the
uptake of four electrons - two from the sulfur, and twc?
through formation of the counter-ions.
A second group of elements with apparently lower
coordination numbers forms ions which possess s-electrons in addition to ten d-electrons. A hybrid of s- and
p-orbitals is now partially occupied by the electrons of
the atom,and this leads to development of extremely
unsymmetrical and low coordination states. Numerous
examples of this are provided by the structures of conipounds of TI+, Sn2+, Pb2+, Bi3 I, and Sb3+.
The structure of metallo-complexes becomes understandable if it is assumed that the more electropositive
atom of a complex, whether it be the central atom or a
ligand, always appears to have a stronger tendency
towards hybridization than the more electronegative
atom. In general, the negative atom (central atom) is
already hybridized. The tendency of the positive ligands
to form hybrids will therefore be even greater. However,
since hybridization allows the existence of at least two
bonds in two directions, the ligand cannot be terminal
in the structural group, i. e. it can never be monodentate.
Instead, it will either attach itself to a second central
atom, giving rise to polynuclear complexes, or will become saturated by other anionic ligands to form a
two-shell complex. Conversely, the formation of twoshell complexes is not to be expected with cationic
central atoms, since a cationic element cannot terminate a structural group. Thus, Werner’s [16] example
[Co(NH3)5NCSAg] [NO313 probably has some other
structure. The only exception to this rule is the hydrogen
atom. Owing to its small size, it is largely embedded in
the electron cloud of its bonding partner, so that
hydroxo-complexes are not regarded as two-shell complexes, nor hydroxides as complexes.
I t follows also from the structure of the metallo-complexes that not every element can participate in complexes of this type. As may be seen from Table 3, all
elements which occur as a n i o n s are suitable, but of the
cationic elements, only those which have low specific
coordination numbers can take part. The reason is that
the coordination number of the ligands must, by de-
5. Properties of Metallo-Complexes
Lieser determined the stability constants of argentohalogen complexes [82] in aqueous solution; these
increase with the size of the central atom. However, it
is not known to what extent the tendency of metal
ligands towards coordination in solution is satisfied by
the associated water molecules. The roles of the central
atom and the ligand can be interchanged in competitive
reactions, so that complexes with cationic central atoms
are formed. Thus, for example, all argento-complexes
decompose upon contact with molecules containing
nitrogen, e. g. ammonia or acetonitrile, the corresponding silver complex being formed immediately.
The metallo-complexes are predominantly polynuclear
complexes. The most notable of their reactions are
those in which the coniplex ion remains stable and the
compounds assume ion-exchanger characteristics.Thus,
Lieser was able to exchange nitrate ions in [IAg3] “ 0 3 1 2
for hydroxide or chloride ions. Weiss et al. displaced
alkali ions from K[NHg(SO3)] with tetraalkylanvnonium ions. Similarly, [SzHg3]CI2 reacts with dilute
nitric acid to yield [S2Hg3] [NO312 1401, and almost any
desired group can be built into the framework of the
[CHg3O]+ ion.
These chemical properties show clearly how reasonable
it is to single out structural groups from lattices,
even when the former occur as “infinitely” extended
components of a complex compound.
I should like lo thank Prof. 0. Schmitz-Du Mont and the
Deutsche Forschirngsgemrinschaft for their. generous
.y~pportCV-tllt.7 tvorli.
Received, December 2nd, 1963 [A 383/179 I€]
German version: Angew. Chem. 76, 697 (1964)
Translated by EKpress Translation Service. London
[79] R . L . Suss, R. Vidale, and J. Donohue, Acta crystallogr. 10,
567 (1957).
[80] C.Scheringer, Acta crystallogr. 16, 546 (1963).
[XI] L. E. Orgel, J. chem. SOC.(London) 1958,4186; J. D . Dunitz
and L. E . Orgel, Advances inorg. Chem. Radiochem. 2 , 3 4 (1960).
[82] K . H . Lieser, Z . anorg. allg. Chem. 304, 296 (1960).
Angew. Chem. internat. Edit. / Vol. 3 (19641 1 No. 10
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