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How do we see colour?. Most transition metal compounds appear coloured. This is because they absorb energy corresponding to certain parts of the visible electromagnetic spectrum. The colour that is seen is made up of the parts of the visible spectrum that aren’t absorbed.
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How do we see colour? Most transition metal compounds appear coloured. This is because they absorb energy corresponding to certain parts of the visible electromagnetic spectrum. The colour that is seen is made up of the parts of the visible spectrum that aren’t absorbed. For example, a red compound will absorb all frequencies of the spectrum apart from red light, which is transmitted.
What happens when light is absorbed? In transition metal ions, the d sub-level is only partially filled. This means that electrons can move between d orbitals. In a transition metal complex, the relative energies of the d orbitals change. Electrons can be promoted to higher energy orbitals. For electrons to be promoted, they need to absorb light energy of a particular frequency. This frequency depends on the precise difference in energy between the d orbitals.
Factors affecting colours The colour of a transition metal compound is determined by the difference in energy between its d orbitals. This can be affected by several factors: • size and type of ligands • coordination number • strength of metal–ligand bonds • oxidation state. • complex shape [Cr(H2O)6]3+ [Ni(H2O)6]2+ [Fe(H2O)6]2+ [Fe(H2O)6]3+
Colours of complexes A transition metal will appear different colours in complexes with different ligands. For example: [Cu(H2O)6]2+ [CuCl4]2–
Determining concentration Ultraviolet–visible spectroscopy can be used to determine the concentration of a transition metal complex solution. A UV-Vis spectrometer passes different frequencies of light through a sample. Some of the light is absorbed while the rest passes through. A detector measures the absorbance of the sample. The amount of light absorbed is proportional to the concentration of the absorbing species.
Improving UV–Vis spectroscopy 1 Some transition metal complex ions have a very pale colour. Accurate quantitative determination of the concentration of these solutions is difficult, because the difference in absorption between the different concentrations is too small. Replacing the ligands in a complex changes its colour. This can increase the absorption of the transition metal species, allowing lower concentrations to be analysed using UV–Vis spectroscopy.
[Fe(H2O)6]2+ + 3bipy [Fe(bipy)3]2+ + 6H2O Improving UV–Vis spectroscopy 2 [Fe(H2O)6]2+ is a pale green colour. If the water ligands are replaced by other types of ligands, a much stronger coloured solution is produced. The absorption values increase by 103, allowing analysis of lower concentrations.
[M(H2O)6]3+ + H2O [M(H2O)5OH]2+ + H3O+ Hydrolysis reactions Transition metal salts dissolve in water to form aqua ions – complexes with water molecule ligands. Some of these aqueous complexes are acidic, for example [Fe(H2O)6]3+. This is because the Fe3+ ion is strongly polarizing and weakens the O–H bonds in the water ligands. The complex ion releases an H+ ion, producing an acidic solution. The reaction is called hydrolysis because the water molecule is being split. The general equation for this reaction is:
Why are some aqua ions less acidic? A solution of [Fe(H2O)6]3+ is highly acidic, whilst a solution of [Fe(H2O)6]2+ is only very weakly acidic. This is because the polarizing power of the metal ion depends on its size and charge. Smaller, more highly-charged metal ions exert a greater polarizing effect on the water ligands, so that more O–H bonds break, releasing H+ ions. As a general rule, M3+ ions are significantly more acidic than M2+ ions.
[M(H2O)6]2+(aq) + H2O(l) [M(H2O)5OH]+(s)+ H3O+(aq) [M(H2O)5OH]+(s) + H2O(l) [M(H2O)4(OH)2](s) + H3O+(aq) Hydrolysis of M2+ ions A series of hydrolysis reactions can occur until the overall charge on a complex is 0. For example, in transition metal 2+ ions, two H+ ions are removed.
[M(H2O)6]3+(aq) + H2O(l) [M(H2O)5OH]2+(aq)+ H3O+(aq) [M(H2O)5OH]2+(aq) + H2O(l) [M(H2O)4(OH)2]+(aq) + H3O+(aq) [M(H2O)4(OH)2]+(aq) + H2O(l) [M(H2O)3(OH)3](s) + H3O+(aq) Hydrolysis of M3+ ions Three H+ ions are removed from transition metal 3+ ions.
Reactions with bases If a base such as ammonia or hydroxide ions is added to a solution of transition metal aqua ions, H+ ions are removed from the water ligands until there is no overall charge on the complex. The final product is an uncharged, insoluble metal hydroxide that forms a precipitate. This reaction occurs in a series of steps, depending on whether the metal is a 2+ or 3+ ion.
2[Cr(H2O)3(OH)3](s)+ 3CO2(g) +3H2O(l) 2[Cr(H2O)6]3+(aq)+3CO32–(aq) [Cr(H2O)6]2+(aq)+CO32–(aq) CrCO3(s)+6H2O(l) Reactions with carbonate ions Sodium carbonate acts as a base with M3+ aqua ions to produce a hydrated metal hydroxide. However, adding sodium carbonate to M2+ ions produces an insoluble metal carbonate. The M2+ ions are less acidic, and carbonate ions are unable to remove protons from the water ligands, so they displace the ligands instead.
Cr(H2O)3(OH)3 + 3H3O+ [Cr(H2O)6]3+ + 3H2O Cr(H2O)3(OH)3 + 3OH– [Cr(OH)6]3– + 3H2O Amphoteric character Some metal hydroxides can react as both an acid and as a base and are known as amphoteric. For example, chromium hydroxide will react in the following ways: With strong acids: With strong, excess alkali:
[Cu(H2O)6]2+ + 4Cl–[CuCl4]2–+ 6H2O Ligand substitution A ligand substitution reaction occurs when a ligand in a complex ion is replaced by another type of ligand molecule. When concentrated hydrochloric acid is added to a solution of hexaaquacopper(II), chloride ions replace the water molecules as ligands. Chloride ions are much larger than water molecules so only four can fit around the copper ion. This means that the complex changes shape from octahedral to tetrahedral. The colour of the complex changes from blue to green.
[Cu(H2O)6]2+ + NH3 [Cu(NH3)(H2O)5]2+ + H2O [Cu(NH3)(H2O)5]2+ + NH3 [Cu(NH3)2(H2O)4]2+ + H2O [Cu(NH3)2(H2O)4]2+ + NH3 [Cu(NH3)3(H2O)3]2+ + H2O [Cu(NH3)3(H2O)3]2+ + NH3 [Cu(NH3)4(H2O)2]2+ + H2O Stability constants Hexaaquacopper(II) undergoes ligand substitution with ammonia in stages. At each stage, one water molecule is replaced by one ammonia molecule, until four molecules of water have been replaced by four molecules of ammonia. Each of these stages is an equilibrium reaction.
[Cu(H2O)6]2+(aq) + 4NH3(aq) [Cu(NH3)4(H2O)2]2+(aq) + 4H2O(l) [Cu(NH3)4(H2O)22+] Kstab = [Cu(H2O)62+] + [NH3]4 Stability constants For each stage of a reaction, an expression for the equilibrium constant can be written. This equilibrium constant for the overall reaction is called the stability constant: Kstab. Note that the square brackets now mean concentration in moldm-3. The concentration of water is left out because it is in great excess and its concentration is almost constant.
Comparing stability constants The stability constant, Kstab, is the equilibrium constant for the formation of a complex ion from its constituent ions in solution. Kstab values show how stable a complex ion is. Complex ions with large Kstab values are easily formed. Whether a ligand substitution will occur can be predicted from Kstab values, by comparing that of the current complex ion with the value for the substituted complex ion. The most stable ion is most likely to occur.
[M(H2O)6]2+ + EDTA4– [M(EDTA)]2– + 6H2O 2 species 7 species The chelate effect Complex ions containing multidentate ligands such as EDTA are called chelates. They have much larger Kstab values and are far more stable than complex ions containing unidentate ligands. This is because of the effect of entropy. When a multidentate ligand replaces unidentate ligands in a complex, it releases many molecules, increasing the entropy. The reverse reaction involves a large decrease in entropy, which is why it is so unfavourable.
Haemoglobin Haemoglobin is the molecule that causes blood to appear red. It carries oxygen from the lungs to cells in the body. Haemoglobin contains an Fe2+ ion which forms a haem complex with a tetradentate ligand called porphyrin. It also binds to a unidentate globin molecule. One coordination site is left that can bind loosely to an oxygen molecule. Oxygen is a poor ligand that is easily released to cells, where its concentration is low. Ligands that can form stronger bonds with the Fe2+ ion, such as carbon monoxide, bind irreversibly and destroy haemoglobin’s ability to carry oxygen. These substances are toxic.
Redox reactions Transition metals are able to exist in many different oxidation states, which is why they often undergo redoxreactions. Oxidation of transition metals occurs most easily in alkaline solution. This is because negative ions tend to form in alkaline solution and it is easier to lose electrons from a negatively-charged species. Reduction of transition metals occurs most easily in acidic solution.
[Co(H2O)6]2+ + 2OH– [Co(H2O)4(OH)2] + 2H2O [Co(H2O)4(OH)2] + 6NH3 [Co(NH3)6]2+ + 2OH– + 4H2O 4[Co(NH3)6]2+ + O2 + 2H2O 4[Co(NH3)6]3+ + 4OH– 2[Co(OH)6]4– + H2O2 2[Co(OH)6]3– + 2OH– Oxidation of cobalt(II) In ammoniacal solution, Co2+ is oxidized to Co3+ by oxygen in the air. Several reactions occur because ammonia acts as both a base and a ligand: Co2+ can also be oxidized by hydrogen peroxide (H2O2) after adding an alkali such as sodium hydroxide.
2CrO42– + 2H+ Cr2O72– + H2O Cr2O72– + 14H+ + 3Zn 2Cr3+ + 7H2O + 3Zn2+ Zn + 2Cr3+ Zn2+ + 2Cr2+ Reduction of chromium(VI) In aqueous solution, chromium has an oxidation state of 6+. It exists in alkaline solution as CrO42– and as Cr2O72– in acidic solution. Chromium(VI) can be reduced to Cr3+ and Cr2+ by zinc in acid solution . Cr2+ is easily oxidized to Cr3+ in the presence of oxygen, but hydrogen is produced during the reduction, which excludes air.
