Complex Ions Worksheet PDF

## SECTION 2 - Complex ions ### 1) Some definitions | Term | Definition | |---|---| | Ligand | particle with lone pair of electrons that bonds to metals by a co-ordinate bond | | Complex | metal ion with co-ordinately bonded ligands | | Co-ordination number | number of co-ordinate bonds from ligan...

## SECTION 2 - Complex ions ### 1) Some definitions | Term | Definition | |---|---| | Ligand | particle with lone pair of electrons that bonds to metals by a co-ordinate bond | | Complex | metal ion with co-ordinately bonded ligands | | Co-ordination number | number of co-ordinate bonds from ligands to metal ion | | Lewis base | lone pair donor | | Lewis acid | lone pair acceptor | ### 2) Formation of complexes Common, simple ligands include H₂O, NH₃ and Cl^- ions, all of which have a long pair of electrons. Complexes are formed when ligands such as these form co-ordinate bonds using their lone pairs to a metal ion. In this example, 4 chloride ions (Cl^-) act as ligands, each forming a single co-ordinate bond to the Cu²⁺ ion using a lone pair, forming the [CuCl₄]^2- ion. Ligands are acting as Lewis bases when they bond to transition metals as they donate a lone pair to form a co-ordinate bond. The metal ion acts as a Lewis acid as it accepts lone pairs. ### 3) Shapes of complexes The table shows the most common shapes of complexes. | Co-ordination number | Shape | Occurrence | Example | |---|---|---|---| | 2 | Linear | Ag⁺ complexes | [Ag(NH₃)₂]⁺ (in Tollen's reagent) | | 4 | Square planar | Pt²⁺ and Ni²⁺ complexes | [PtCl₄]^2- | | 4 | Tetrahedral | With larger ligands (e.g. Cl^-) (when ligands are too big for six to fit) | [CuCl₄]^2- | | 6 | Octahedral | Most complexes with small ligands (e.g. H₂O, NH₃) | [Cu(H₂O)₆]²⁺ | ### 4) Types of ligands All ligands must have one (or more) lone pair(s) of electrons to form the co-ordinate bond(s) to the metal ion. #### Unidentate ligands - ligands which form one co-ordinate bond to a metal ion | Ligand | Example complex | |---|---| | :Cl^- | [CuCl₄]^2- | | :OH^- | [Cr(OH)₆]^3- | | :CN^- | [Ag(CN)₂]^- | | H₂O: | [Fe(H₂O)₆]²⁺ | | :NH₃ | [Cr(NH₃)₆]³⁺ | #### Bidentate ligands - ligands which form two co-ordinate bonds to a metal ion | Ligand | Example complex | |---|---| | 1,2-diaminoethane | [Cr(NH₂CH₂CH₂NH₂)₃]³⁺ | | Ethanedioate ion (C₂O₄^2-) | [Cr(C₂O₄)₃]^3- | #### Multidentate ligands - ligands which form more than two co-ordinate bonds to a metal ion | Ligand | How it bonds | Example | |---|---|---| | EDTA^4- forms 6 bonds | | [Cu(EDTA)]^2- | | Porphyrin forms 4 bonds | | | ## Haemoglobin Haemoglobin is a globular protein that contains four Fe²⁺ centres, each with a porphyrin ligand taking up four of the six co-ordination sites. One of the other two sites is bonded to the rest of the haemoglobin structure leaving one site to which oxygen can bond as a ligand. The oxygen is carried around the body in blood bonded to iron in haemoglobin in this way. It is not a particularly good ligand and so easily comes off the iron when needed to transfer to cells for respiration. Cyanide ions (CN^-) and carbon monoxide (CO) are better ligands than oxygen (O₂) and so will bond to haemoglobin in preference to oxygen. This is why carbon monoxide and cyanide compounds are toxic as they prevent the transfer of oxygen around the body. ## TASK 1 - Drawing complexes | Formula | Sketch | Shape | Bond angles | Metal oxidation state | Ligands | Co-ordination number | |---|---|---|---|---|---|---| | [Ag(CN)₂]^- | | | | | | | | [Cr(NH₃)₆]³⁺ | | | | | | | | [Ni(NH₂CH₂CH₂NH₂)₃]²⁺ | | | | | | | | [Co(NH₂CH₂CH₂NH₂)₂Cl₂]⁺ | | | | | | | | [Pt(NH₃)Cl₃] | | | | | | | | [Fe(C₂O₄)₃]^3- | | | | | | |