Metal Complexes – HSC Chemistry

 
This is part of the HSC Chemistry course under the section Analysis of Inorganic Substances.

HSC Chemistry Syllabus

  • Conduct qualitative investigations – using flame tests, precipitation and complexation reactions as appropriate – to test for the presence in aqueous solution of the following ions:

– cations: barium (Ba2+), calcium (Ca2+), magnesium (Mg2+), lead(II) (Pb2+), silver ion (Ag+), copper(II) (Cu2+), iron(II) (Fe2+), iron(III) (Fe3+)
– anions: chloride (Cl–), bromide (Br–), iodide (I–), hydroxide (OH–), acetate (CH3COO–), carbonate (CO32–), sulfate (SO42–), phosphate (PO43–) 

    Metal Complexes for HSC Chemistry

    What are Metal Complexes?

    Metal complexes are often coloured substances which consist of a central metal atom, or ion, bonded to surrounding molecules or ions known as ligands. These ligands can be neutral molecules like water `(H_2O)` and ammonia `(NH_3)`, or ions like chloride `(Cl^–)` or hydroxide `(OH^–)`.

     

     
    Figure 1. Copper hexahydrate complex drawn as a coordination sphere (left) and in physical form (right) 

    Metal Complex Formation

     

    The reaction between metal ions and ligands to form metal complexes can reach an equilibrium state. The equilibrium constant (`K_{eq}`) for this reaction can vary significantly depending on the nature of the metal ion, the ligands, the temperature, and other conditions.

    The overall charge of a metal complex is determined by adding together the oxidation state of the metal ion and the charges of all the ligands attached to it. For instance, if a metal ion with a +2 charge binds with four ligands, each with a -1 charge, the overall charge of the complex would be -2 since `(2 + (4 \times (-1)) = -2.`

    Coordination Number

    The coordination number of a metal ion in a complex refers to the number of ligand attachment sites around the metal. It is determined by the size, charge, and electronic configuration of the metal ion, as well as the size and electronic properties of the ligands. Common coordination numbers are 2, 4, and 6, leading to various geometric arrangements such as linear, square planar, and octahedral. 

    The metal complex demonstrated in figure 1 has a coordination number of 6.  

    Coordinate Covalent Bonding

    Coordinate covalent bonds are a specific type of covalent bond where both electrons involved in the bond come from the same atom, typically the ligand. The ligand donates a pair of electrons to an empty orbital of the metal ion, creating a bond that stabilises the metal complex.  

     

    Figure 2. Sequenced formation of coordinate covalent bonds between silver ion and ammonia molecules. Both electrons are donated from the lone pair on the nitrogen of the ammonia to form a singular covalent bond with the metal ion. 

    Types of Ligands

    Ligands can be classified as either unidentate or bidentate. Unidentate ligands are ligands which have only one electron donor atom. This leads to the formation of only a singular coordinate covalent bond with the metal,

    Examples of unidentate ligands include chloride, ammonia and water. 

     

    Figure 3. Both ligands have a coordination number of 6 since the central atom is attached to 6 unidentate ligands. 

     

    Bidentate ligands are ligands which have two electron donor atoms. These electron donors are able to form coordinate covalent bonds simultaneous, leading to the ability to create rings which are known as chelates.

    Oxalate is an example of a bidentate ligand: the two negatively charged oxygen atoms have electron lone pairs that can be donated to form coordinate covalent bonds with metal ions.


    Figure 4. The zirconium complex (right) is attached to 4 oxalate ions (left). Since oxalate is a bidentate ion, which form coordinate covalent bonds at two sites each, the coordination number of the complex is 8. 

    Example: Dissolution of AgCl in Ammonia Solution

     The dissolution of AgCl in ammonia solution is a classic example of metal complex formation. The ammonia molecules act as ligands, bonding with the silver ions to form a [Ag(NH3)2]+ complex, which is soluble in water. This process is represented by the following equations

     

     

    Salt Dissociation:

    $$AgCl(s) \rightleftharpoons Ag^+(aq) + Cl^–(aq)$$

      

    Complex Formation:

    $$Ag^+(aq) + 2NH_3(aq) \rightleftharpoons [Ag(NH_3)_2]^+(aq)$$

     

    Net Equation:

    $$AgCl(s) + 2NH_3(aq) \rightleftharpoons [Ag(NH_3)_2]Cl(aq)$$ 

    Example: Hydration of Aluminium(III)

    Formation equations illustrate the step-by-step assembly of metal complexes. For example, the formation of an Aluminium(III) complex with water ligands can be represented as:

     

    Salt Dissociation:

    $$Al(NO_3)_3(s) \rightarrow Al^{3+}(aq) + 3NO_3^–(aq)$$

     

    Complex Formation:

    $$Al^{3+}(aq) + 6H_2O(l) + 3NO_3^–(aq) \rightarrow [Al(H_2O)_6^{3+}](aq) + 3NO_3^–(aq)$$

     

    Net Ionic Equation:

    $$Al^{3+}(aq) + 6H_2O(l) \rightarrow [Al(H_2O)_6]^{3+}(aq)$$ 

     

     

      RETURN TO MODULE 8: APPLYING CHEMICAL IDEAS