Chapter 15 Complex Ions

Outline 1. Composition of complex ions 2. Geometry of complex ions 3. Electronic structure of complex ions 4. Formation constants of complex ions

Compounds of Transition Metals Components Transition metal Associated ions and molecules Counter-ion

Components of Complex Ions and Compounds of Them Composition of complex ions Geometry of complex ions Electronic structure of central metal atom or ion Equilibrium constant for the formation of the complex ion

Composition of Complex Ions Cu2+ (aq) + 4NH3 (aq) ⇌ Cu(NH3)42+ (aq) When ammonia is added to a solution of aqueous copper(II) ion, the color changes from pale blue to deep blue The species that accounts for the color is a complex made up of copper(II) ion and four associated ammonia molecules

The Nature of the Complex Ion, Cu(NH3)42+ Each ammonia molecule has a lone pair on the nitrogen A bond forms between the ammonia and the copper(II) ion Both electrons come from ammonia Result is a coordinate covalent bond

Terminology The species is referred to as a complex ion Central metal ion Small molecules or ions surround it These are called ligands The number of atoms bonded to the metal is the coordination number This may or may not be equal to the number of ligands, as we will see

Complex Ion or Compound? The species, Cu(NH3)42+, is clearly a complex ion; because of the charge, it cannot be a compound Complex ions can associate with other ions to form compounds: [Cu(NH3)4]F2 Two fluorides are needed to balance the positive charge on the complex The square brackets indicate the complex ion, i.e., everything inside is part of the complex ion

Cu(NH3)42+ The central metal ion is Cu2+ The ligands are the NH3 molecules The coordination number is 4

Figure 15.1

The Nature of the Metal Ion Complex ions form from Transition metals Main-group metals: Al, Sn, Pb Complex ions exist in aqueous solution Consider Zn(NO3)2 In water, the ion exists as Zn(H2O)42+

Lewis Acid-Base Principles Lewis bases donate lone pairs Lewis acids accept lone pairs

The Nature of the Complex Ion Consider that hydrated metal cations are acidic Brønsted Zn(H2O)42+ ⇌ Zn(H2O)3(OH)+ (aq) + H+ (aq) Lewis The metal can accept a pair of electrons and is therefore a Lewis acid Consider that the ligand possesses at least one lone pair This makes the ligand a Lewis base The complex ion can be described as the product of a Lewis acid-base reaction

Lewis and Brønsted Acids and Bases Lewis bases are also Brønsted bases Can accept an electron pair (or pairs) Lewis acids need not be Brønsted acids The Lewis model broadens the definition of an acid

Charges of Complexes Charge of complex = oxidation number of central metal + charges of ligands Platinum(II)

Table 15.1 – Complexes of Pt2+

Example 15.1

Chelating Agents Ligands Any molecule or anion with at least one unshared pair of electrons Some ligands have more than one pair of unshared electrons These ligands can coordinate using multiple pairs of electrons Ligands that can form more than one bond to a central metal ion are called chelates

Two Common Chelating Ligands Oxalate, C2O42- Ethylenediamine, H2N-CH2CH2-NH2 Both of these ligands coordinate in two places per ligand species Bidentate chelating ligand

Figure 15.2

Geometry of Complex Ions Review of Chapter 7 Recall that the geometry of a molecule can be determined by the way in which the central atom coordinates with terminal atoms For a complex ion, the geometry is determined by the shape taken by the complex, as determined by the coordination number and nature of the metal ion and ligands

Coordination Number The most common coordination number is 6 Geometry is octahedral Two other coordination numbers are common 2; Geometry is linear 4; Geometry may be square planar or tetrahedral

Table 15.2

Figure 15.2

Example 15.2

Coordination Number 2 2-coordinate complexes are linear CuCl2- Ag(NH3)2+ Au(CN)2-

Coordination Number 4 Two geometries exist Tetrahedral Square planar Zn(NH3)42+ CoCl42- Square planar Characteristic of Cu2+ and metal ions with d8 configurations (Pt2+, Ni2+)

Figure 15.3

Coordination Number 6 6-coordinate complexes are octahedral The six ligands are equidistant from the central metal The octahedron can be considered a derivative of a square plane, with two ligands added, one above and one below the plane

Example 15.3

Isomerism Two or more species with the same formula, but different chemical and physical properties are called isomers Complex ions can show several different types of isomerism Only type to be considered here is geometric isomerism Only the spatial orientation of ligands differs between geometric isomers

Isomerism in Square Planar Complexes Pt(NH3)2Cl2 Can have two different isomers: cis and trans Complexes of the form Ma2b2 will show cis-trans isomerism a and b are different ligands

Meaning of cis and trans Cis positions are 90° apart Trans positions are 180° apart

Figure 15.4

Isomerism in Octahedral Complexes Co(NH3)4Cl2 or Ma4b2

Chelates and Isomers In general, chelating ligands can bridge only cis positions The bridge is not long enough to stretch across a trans position Chelates, due to their binding in two locations, generally produce more stable complexes Partially a result of the geometry (ring size) Partially as a result of the nature of the bonds

Figure 15.6

Figure 15.7

Physical Properties of Isomers Note in the last figure that the color of the two complexes differ The cis isomer is reddish-purple The trans isomer is green Geometric isomerism can lead to great differences in the physical and chemical properties of compounds containing complex ions

Electronic Structure of Complex Ions Crystal Field Model explains Color of transition metal complexes Magnetic properties of transition metal complexes Considers the nature of the Metal Ligands

Transition Metal Cations In a simple transition metal cation There are no outer s electrons Electrons are distributed among the five d orbitals by Hund’s Rule Recall that Hund’s rule results in the maximum number of unpaired spins Magnetic properties depend on distribution of electrons Diamagnetism: no unpaired electrons Paramagnetism: unpaired electrons

Iron(II) ion Iron(II) ion has 26-2 or 24 electrons Shorthand notation: [Ar]3d6 Orbital diagram: [Ar] (↑↓) (↑ ) (↑ ) (↑ ) (↑ ) Fe2+ is paramagnetic; there are four unpaired electrons

Figure 15.8 – Colors of Transition Metal Compounds

Example 15.4

d-orbitals Recall that there are five d orbitals In uncomplexed metal cations, these orbitals have the same energy

Figure 15.9

Octahedral Complexes We can collect the d orbitals into two groups: A high energy pair, the dx2-y2 and dz2 A low energy triplet, the dxy, dxz and dyz

Why the Split into Two Groups? In the absence of any ligands, the d orbitals have the same energy Note that the two orbitals whose energy is considered high lie on the xyz axes Electron density in metal interacts with the electron density of ligands brought toward the metal on the axes Note that the two orbitals whose energy is considered lower lie between the xyz axes There is less interaction between the metal electron density and that of the ligand

The Splitting Energy, Δo The crystal field splitting energy is given the symbol Δo The magnitude of Δo will determine the way in which the electrons fill the d orbitals in the metal ion of the complex If Δo is large, electrons will pair in the lower energy orbitals before occupying the higher energy ones If Δo is small, electrons will distribute themselves in all five orbitals, pairing only in cases where there are more than five electrons

Figure 15.10

Figure 15.11

Electronic Arrangements in Complexes When Δo is large, a low spin complex results Electrons fill the lower three orbitals before occupying the upper two When Δo is small, a high spin complex results Electrons distribute themselves to all five orbitals by Hund’s rule

Notes on High and Low Spin For a given cation, a high spin complex will always have a larger number of unpaired electrons than a low spin complex The value of Δo is determined by the nature of the ligand(s) Strong field-ligands produce low spin complexes Example: CN-, NH3 Weak field-ligands produce high spin complexes Example: H2O, Cl-

Example 15.5

Color Most transition metal complexes are brightly colored Exception: those with empty d sublevels (e.g., Sc3+; those with full d sublevels (e.g., Zn2+) The splitting of the d sublevel results in an energy difference that corresponds to the visible region of the electromagnetic spectrum Visible light is absorbed in the transition of an electron in the ion Some of the wavelengths of white light are removed Complex appears colored

Figure 15.12

Titanium(III) Consider Ti3+ in a complex There is only one d electron Because there is only one possible electronic transition, it is possible to calculate Δo for this ion The ion absorbs at 510 nm (green) The complex appears as the color complement of green (i.e., red-violet or purple)

Calculating Δo for Titanium(III) By knowing the energy of the light absorbed, it is possible to calculate the value of Δo For ions with more than one d electron, calculating the energy Δo is more difficult

Δo and Wavelength The smaller Δo is, the longer the wavelength of absorption Small Δo stem from weak field ligands Large Δo stem from strong field ligands

Light Absorption and Color

The Spectrochemical Series of Ligands CN- > NO2- > en > NH3 > NCS- > H2O > F- > Cl- Strong field weak field

Equilibrium and Complex Ions We can consider the formation of a complex ion in aqueous solution as an equilibrium between The reactants Bare metal ion and ligands The product The complex ion As with acid-base and gas phase equilibria, an expression for K can be written

Formation Constants Consider Cu2+ (aq) + 4NH3 (aq) ⇌ Cu(NH3)42+ (aq) We can write the expression for Kf (the formation constant for the ion: In many cases, Kf is a very large number, indicating that the complex which forms is very stable

Table 15.4

Example 15.6

Example 15.6 (Cont’d)

Practical Example Most cleaning compounds containing ammonia state that the product is not to be used on objects made of silver Consider the Kf for Ag(NH3)2+: 1.7 X 107 The complex of silver with ammonia is a very stable one Using an ammonia-based cleaner on silver will partially dissolve the silver

Chelates in Nature Chelating agents are abundant in nature Iron is complexed with the chelate heme EDTA is used to complex metal ions Treat heavy metal poisoning Mask metal ions that promote food spoilage

Heme

EDTA

Key Concepts 1. Relate the composition of a complex ion to its charge, coordination number, and the oxidation number of the central metal 2. Sketch the geometry of a complex ion and identify geometric isomers 3. Give the electron configuration and the orbital diagram for the transition metal at the center of the complex 4. Derive the orbital diagram (high or low spin) for the complex 5. Relate the formation equilibrium constant to the concentrations of metal ion, ligands and complex ion