Shapes of Complex Ions

10 Key excerpts on "Shapes of Complex Ions"

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  Chemistry

  Principles and Reactions

  • William Masterton, Cecile Hurley(Authors)
  • 2020(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 494 CHAPTER 19 Complex Ions ▼ 19-3 Geometry of Complex Ions The physical and chemical properties of complex ions and of the coordination compounds they form depend on the spatial orientation of ligands around the central metal atom. Here we consider the geometries associated with the coordi-nation numbers 2, 4, and 6. With that background, we then examine the phenom-enon of geometric isomerism , in which two or more complex ions have the same chemical formula but different properties because of their different geometries. Coordination Number 5 2 Complex ions in which the central metal forms only two bonds to ligands are linear; that is, the two bonds are directed at a 180 8 angle. The structures of CuCl 2 2 , Ag(NH 3 ) 2 1 , and Au(CN) 2 2 may be represented as H H H H H 9 N 9 Ag 9 N 9 H (Cl 9 Cu 9 Cl) 2 (N # C 9 Au 9 C # N) 2 1 Coordination Number 5 4 Four-coordinate metal complexes may have either of two different geometries (Figure 19.4). The four bonds from the central metal may be directed toward the corners of a regular tetrahedron . This is what we would expect from the VSEPR model (recall Chapter 7). Two common tetrahedral complexes are Zn(NH 3 ) 4 2 1 and CoCl 4 2 2 . Square planar complexes, in which the four bonds are directed toward the corners of a square, are more common. Certain complexes of copper(II) and nickel(II) show this geometry; it is characteristic of the complexes of Pd 2 1 and Pt 2 1 , including Pt(NH 3 ) 4 2 1 . Coordination Number 5 6 We saw in Chapter 7 that octahedral geometry is characteristic of many mole-cules (e.g., SF 6 ) in which a central atom is surrounded by six other atoms.

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  Chemistry

  With Inorganic Qualitative Analysis

  • Therald Moeller(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  28

  THE CHEMISTRY OF COMPLEXES

  Publisher Summary

  This chapter presents the nomenclature of complexes and describes their geometry, isomerism, and general properties. It discusses three different approaches to the explanation of bonding in complexes: valence bond theory, molecular orbital theory, and crystal field theory. It further discusses the equilibria among the components of complexes. The chapter illustrates a few practical applications of complexes. The term “complex” is usually reserved for metals combined with donors that also can exist independently either in the pure state or as ions in solution. The molecule or ion that contains the donor atom is called “ligand.” The neutral compound formed between a complex ion and other ions or molecules is called a coordination compound. The coordination number is the number of nonmetal atoms surrounding the central metal atom or ion in a complex. Most metal atoms or ions can accept more than one pair of electrons. The chapter explains the concept of chelation. All metal ions have the ability to form coordination compounds. Bonding in complexes, as in other compounds, is rarely strictly ionic or strictly covalent. The valence bond approach to complex bonding emphasizes covalent bonding, the crystal field theory emphasizes ionic bonding, and the molecular orbital theory brings about a compromise between the two.

  This chapter presents the nomenclature of complexes and describes their geometry, isomerism, and general properties. Three different approaches to the explanation of bonding in complexes are introduced. Two of them—valence bond theory and molecular orbital theory—we have used before. The third—crystal field theory—has not been mentioned before in this book. The equilibria among the components of complexes are discussed, and a few practical applications of complexes are mentioned.

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  Introduction to Coordination Chemistry

  • Geoffrey A. Lawrance(Author)
  • 2013(Publication Date)
  • Wiley

    (Publisher)

  The simple amended VSEPR point-charge model discussed in Chapter 3 is based in effect on the third of the above, but despite this provides a basis for predicting shape. However, it must, because of its limitations, be deficient in predicting shape in metal complexes. We shall see as we explore actual shapes below that it is, nevertheless, a good starting point.

  4.2 Forms of Complex Life – Coordination Number and Shape

  Molecules are certainly more varied than life forms. Even carbon-based compounds can be considered as unlimited in number, despite the fact that they almost exclusively involve four bonds around each carbon centre in a very limited number of shapes. When we move to coordination compounds, the range of coordination numbers and shapes is expanded considerably, so that coordination complexes live up to their name – they are inherently complex molecular forms. Fortunately, we can identify a number of basic shapes and even some system that governs outcomes – that is, there is some predictive aspect to shape in coordination complexes. We shall examine complexes from the perspective of coordination number below.

  4.2.1 One Coordination (ML)

  This unlikely coordination number suffers from the fact that a single donor bound to the metal would still leave the metal highly exposed, a situation that would most likely lead to additional ligands adding and thus increasing the coordination number. It is nevertheless prudent to describe it as extremely rare, because there is a small possibility that a suitably bulky and appropriately shaped ligand may achieve one-coordination. It may be more practicable in the gas phase under high dilution conditions, where metal–ligand encounters are limited.

  As a consequence, it is not surprising that there appears to be only one isolated structure claimed. This is of the indium(I) and thallium(I) complexes with a single M—C bond from a σ-bonded benzene anion that carries two bulky tri-substituted benzene substituents in ortho positions; these partially block approach of other potential donors to the metal cation (Figure 4.3

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  Chemistry

  The Molecular Nature of Matter

  • Neil D. Jespersen, Alison Hyslop(Authors)
  • 2021(Publication Date)
  • Wiley

    (Publisher)

  To understand the theory, therefore, it is essential that you know the shapes and orientation of the d orbitals. The d orbitals were described in Chapter 7, and they are illustrated again in Figure 21.14. First, notice that four of the d orbitals have the same shape but point in different directions. These are the d x 2 − y 2, d xy , d xz , and d yz orbitals. Each has four lobes of electron density. The fifth NOTE More complete theories consider the covalent nature of metal—ligand bonding, but crystal field theory nevertheless provides a useful model for explaining the colors and magnetic properties of complexes. NOTE The labels for the d orbitals come from the mathematics of quantum mechanics. FIGURE 21.13 Colors of complex ions depend on the nature of the ligands. Each of these brightly colored solutions contains a complex ion of Co 3+ . The variety of colors arises because of the different ligands (molecules or anions) that are bonded to the cobalt ion in the complexes. Michael Watson 1054 CHAPTER 21 Metal Complexes FIGURE 21.14 The shapes and directional properties of the five d orbitals of a d subshell. x y z d xz x y z d xy x y z d yz x y z d z 2 x y z d x 2 – y 2 d orbital, labeled d z 2, has two lobes that point in opposite directions along the z axis plus a small donut-shaped ring of electron density around the center that is concentrated in the xy plane. Of prime importance to us are the directions in which the lobes of the d orbitals point. Notice that three of them—d xy , d xz , and d yz —point between the x, y, and z axes. The other two— the d z 2 and d x 2 − y 2 orbitals—have their maximum electron densities along the x, y, and z axes. Now let’s consider constructing an octahedral complex within this coordinate system. We can do this by bringing ligands in along each of the axes as shown in Figure 21.15.

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  Chemistry

  The Molecular Nature of Matter

  • James E. Brady, Neil D. Jespersen, Alison Hyslop(Authors)
  • 2014(Publication Date)
  • Wiley

    (Publisher)

  Each of these brightly colored solutions contains a complex ion of Co 3+ . The variety of colors arises because of the different ligands (molecules or anions) that are bonded to the cobalt ion in the complexes. x y z d xz x y z d xy x y z d yz x y z d z 2 x y z d x 2 – y 2 Figure 21.13 | The shapes and directional properties of the five d orbitals of a d subshell. Michael Watson 21.5 | Bonding in Metal Complexes 1019 In an isolated atom or ion, all of the d orbitals of a given d subshell have the same energy. Therefore, an electron will have the same energy regardless of which d orbital it occupies. In an octahedral complex, however, this is no longer true. If the electron is in the d x 2 -y 2 or d z 2 orbital, it is forced to be nearer the negative charge of the ligands than if it is in a d xy , d xz , or d yz orbital. Since the electron itself is negatively charged and is repelled by the charges of the ligands, the electron’s potential energy will be higher in the d z 2 and d x 2 -y 2 orbitals than in a d xy , d xz , or d yz orbital. Therefore, as the complex is formed, the d subshell actually splits into two new energy levels, as shown in Figure 21.15. Here we see that regardless of which orbital the electron occupies, its energy increases because it is repelled by the negative charges of the approaching ligands. However, the electron is repelled more (and has a higher energy) if it is in an orbital that points directly at the ligands than if it occupies an orbital that points between them. In an octahedral complex, the energy difference between the two sets of d-orbital energy levels is called the crystal field splitting. It is usually given the symbol ∆ (delta), and its magnitude depends on the following factors: The nature of the ligand. Some ligands produce a larger splitting of the energies of the d orbitals than others. For a given metal ion, for example, cyanide always gives a large value of ∆ and F - always gives a small value.

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  Solvent Extraction Principles and Practice, Revised and Expanded

  • Jan Rydberg(Author)
  • 2004(Publication Date)
  • CRC Press

    (Publisher)

  Figure 3.7 shows common stereochemistries of the elements [5]. For electrostatic interaction, the radius ratio can be used as a guide to the possible geometry of the complex. However, when the metal-ligand bond has a significant covalent contribution, the geometry is fixed by the necessity to have the bonding orbitals of the metal and the donor atom of the ligand to overlap. Generally, this involves hybrid bond orbitals of the metal. The sp 3 and dsp 2 hybrid orbitals result in coordination number 4 with tetrahedral (sp 3 ) and square planar (dsp 2 ) geometry. For coordination number 6 the d 2 sp 3 hybrid Complexation of metal ions 91 Fig. 3.5 Correlation of log K n for formation of UO 2 L n (n=1 or 2) with ΣpK a of a series of salicylates. orbitals produce an octahedral configuration. Thus is tetrahedral while is square planar and is octahedral. Both and would have too much electrostatic repulsion to form the octahedral complexes. The reason why the cyanide forms a planar complex is because CN − tends to form bonds that are more covalent than those formed by Cl − , and that the dsp 2 Fig. 3.6 Structure of the 1:1 complex uranyl dibromosalicylate. Solvent extraction principles and practice 92 configuration (square planar) is associated with complexes with more covalent nature than the tetrahedral complexes of sp 3 configuration. However, neutral ligands such as H 2 O and NH 3 can form octahedral structures and Figure 3.8 shows examples of these geometric structures. In summary, the geometry of transition metal complexes is determined by the necessity (1) to group the ligands about the metal to minimize electrostatic repulsions and (2) to allow overlap of the metal and ligand orbitals. The first Fig. 3.7 Common stereochemistries of the elements.

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  Valency and Molecular Structure

  • E. Cartmell, G. W. A. Fowles(Authors)
  • 2013(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  Qualitative calculations and predictions about spectra and magnetic properties can be made by means of the molecular-orbital method and we shall discuss these problems in rather more detail later in this chapter and again in Chapter 12. 11.2.3 CRYSTAL-FIELD AND LIGAND-FIELD METHODS 3 » 6 » 7 The earliest electrostatic models attempted to account for the properties of complex compounds on the basis of interactions between the metal ions and ligands considered as point charges or dipoles, and while these successfully explained why tetrahedral and octahedral configurations were to be expected for four-co-ordinate and six-co-ordinate complexes, respect-ively, they could not account satisfactorily for the existence of square-planar complexes. Difficulties also arise with ligands such as carbon mono-xide, which formed stable complexes despite being non-polar. Moreover, the theory predicted that for a given set of ligands the smallest ions should form the strongest bonds, and yet second- and third-row transition elements are known to form more stable complexes than the smaller ele-ments of the first row. This simple electrostatic theory was subsequently extended to include the effects of the ligand point charges or dipoles on the d electrons of the metal ion. This extension, called the 'crystal-field' theory, was origin-ally applied to the behaviour of metal ions in a crystal lattice, but it may be applied equally well to complexes in which the d electrons of the metal 210 Complex Compounds Figure 11.6 d Orbitals: (a) d x 2_^ and ά χ 2; (b) dxy, dxz and d yz will be affected by the charge field produced by the six ligands (the ligand field).

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  Medicinal Chemistry

  An Introduction

  • Gareth Thomas(Author)
  • 2011(Publication Date)
  • Wiley

    (Publisher)

  coordination number of the metal atom, the most common being four and six. Ligands may be classified according to their coordination number of donor atoms as uni-, di-, tri-, tetra- dentate, etc. or using the Greek prefixes mono-, bi-, ter-, quadri- dentate, etc. ligands. Both systems are used in the literature.

  The geometric arrangements of the coordinated atoms of the ligands bonded to the metal in naturally occurring complexes usually approximate to those shown in Table 13.1 . For example, the atoms bonding to a four-coordinate zinc ion are usually arranged in a distorted rather than a regular tetrahedral shape about the zinc ion. It is possible for an element to exhibit more than one geometry in both different complexes and in the same complex (see section 13.2.4).

  Table 13.1

  The common geometrical arrangements of coordinated atoms (L) about a central metal, ion

  Coordination number	Common geometric arrangements	Examples of metals that can exhibit this arrangement
  2		Au(I), Ag(I), Hg(II), Cu(I)
  3		Cr(III), Fe(III)
  4		Tetrahedral: Co(II), Cu(II), Zn(II),Fe(III), Co(IV), Ti(IV), Ni(II)Square planar: Cu)II), Pt(II), Ni(II), Cr(II), Mn(III)
  5		Square pyramidal: Mo(IV), Cu(II), Fe(II).Trigonal bipyramidal: V(IV), Nb(IV), Ta(V)
  6		Co(IV), Fe(II), Mg(II), Cr(III)

  Multidentate ligands that form ring structures in which the ligand is bonded by more than one atom to a single metal atom are known as chelating agents . Chelating agents that bind strongly to metal cations to form stable water-soluble complexes are known as sequestering agents . The complexes produced by chelating agents are known as metal chelates or chelation compounds . The formation of ring systems may impose restrictions on the stereochemistry of the complex. For example, the flexible diethylenetriamine forms rings in which the three bonding atoms and the metal atom do not have to be in the same plane. However, rings will only be formed by the rigid, fully conjugated terpyridine if the bonding atoms and the metal atom are all in the same plane.

  The number of electrons donated to the metal atom

  Ligands may be classified as one, two, three… eight electron donors to the metal atom. For example, ligands such as methyl, fluorine and hydroxide ions that form a normal covalent bond with the metal are classified as one-electron donors whilst ligands that form dative bonds by the donation of two electrons from the ligand are said to be two-electron donors, and so on (Table 13.2 ). However, with some ligands, the number of electrons donated will depend on the nature of the complex in which they occur. For example, bromine normally acts as a one-electron donor, but in the form of a ligand bridge it acts as a three electron donor (see section 13.2.2). However, the symbol ηn before the name of a ligand implies that n atoms of the ligand are involved in the bonding but not necessarily n

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  Analytical Chemistry

  • Clyde Frank(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  Chapter Fifteen

  Complexes in Analytical Chemistry: Complexometric Titrations

  Publisher Summary

  A complex ion is one in which part or all of the coordination positions are occupied. Only in a gas phase at a high temperature, it is possible for a metal ion to exist in a simple uncoordinated state. The instant the metal ion is dissolved in a solvent, a solvent sheath forms around the metal ion and its anion by occupying the coordination positions. The extent of solvation and the number of coordinated solvent molecules are determined by the type of metal ion and solvent. The maximum number of ligands that can coordinate to the central ion is given by the coordination number for the central ion. Ligands are attached to the central ion at only one point and are called unidentate ligands. If each ligand has two or more coordinating sites, the ligands are called multidentate ligands. In this system, the coordination results in the formation of rings. This particular type of coordination compound is called a chelate and the ligands involved in the coordination are chelating agents.

  INTRODUCTION

  A complex ion is one in which part or all of the coordination positions are occupied. In general, only in a gas phase at a high temperature is it possible for a metal ion to exist in a simple uncoordinated state. The instant the metal ion is dissolved in a solvent, a solvent sheath (solvation) forms around the metal ion and its anion by occupying the coordination positions. The extent of solvation and number of coordinated solvent molecules will be determined by the type of metal ion and solvent.

  If the colors of dilute solutions of Cu(ClO4 )2 , CuSO4 , and CuCl2 are compared, it is observed that the shade of blue is different for the three Cu2+ solutions. It must be concluded that the Cu2+ in the three solutions is coordinated differently and coordination between Cu2+

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  Fundamentals of Inorganic Chemistry

  An Introductory Text for Degree Studies

  • J Barrett, M A Malati(Authors)
  • 1997(Publication Date)
  • Woodhead Publishing

    (Publisher)

  stability is enhanced by the interaction with its counter ion (in the case of a solid) or with the solvent (in solution). Highly charged (positive or negative) complexes are favoured by this factor (in spite of the electroneutrality principle). Metals in their zero oxidation states forming complexes with neutral ligands are exempt from this consideration. Interligand repulsion and/or steric hindrance Interligand repulsion may well influence the formula of a complex but is difficult to separate from steric hindrance (in which ligands would be closer together than indicated by their normal Van der Waals radii). There is some evidence for the ligands Cl and O 2 being too large in the Van der Waals sense to form ML 6 complexes where M is a first-row transition element. Iron(III) forms [FeCl 4 ] with the larger chloride ion, but forms [FeF 6 ] 3 with the smaller fluoride ion. In non-aqueous solution and in the crystalline state the [FeCl 4 ] ion is Sec. 10.3] The angular overlap approximation 219 tetrahedral, but in aqueous solution two water molecules are included in the axial positions. The [C0CI4] 2 complex ion is tetrahedral in the solid state and in aqueous solution. There are no cases of complexes with greater than four O 2 ligands. This may be due to steric hindrance or to the large repulsion forces in operation or to both effects. Generalizations Bearing in mind the above calculations and factors there are only three general cases which apply to the majority of existing complexes. Weak metal-ligand bonding If the metal-ligand interaction is weak the energy gaps between the d-type orbitals are relatively small. This allows all he anti-bonding orbitals to be occupied if necessary. It has the consequence that the complexes in this category are those with high spin—they have the maximum number ofunpaired electrons with parallel spins.

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