Properties of Transition Metals - Essay Example

TRANSITION METALS They possess the features of metals, hence the reason why they are regarded as transition metals. These elements possess an incomplete d shell. This definition excludes cadmium, Zinc and mercury from the list of transition metals because of their d 10 configuration. The elements are hard and have high boiling and melting points. As one moves across the periodic table from the left to the right, the d orbitals are filled more and more. According to this loose definition, they are 40 elements in number and range from 21-30, 39-48, 71 to 80 and finally 103 -112. The term transition emanates from their location in the periodic elements. They represent the transition from the group two elements to group 13 elements. They are able to up more than eight electrons in their shells. An example is Scandium that has 21 electrons and forms an electronic configuration of 2-8-9-2 (Henry, 2002) The maximum number of electrons for an element in each shell is 2-8-18-32-32-18-2. This means the 3rd and 4th shell of Scandium is unstable and would give up electrons in reaction without a lot of energy. Their electrons are bound loosely, which is responsible for their malleability and conductivity. The exhibit low ionization energies compared to other elements and have a varied oxidation states which is why they form different partial and full ionic compounds. Examples of Transition Metals and the reasons why they form different coloured solutions Transition metals forms coloured solutions when they are dissolved in liquid solutions. There are many reasons for this phenomenon but one known reason why they form different colours is because of the different oxidation states that they exhibit. Below are some of the examples with their reactions under different oxidation states. COPPER Copper (II) ions Reaction in Solution The known simple copper ion is the hexaaquacopper (II) ion [Cu (H20)6]2+ which is typically blue in colour. When the hexaaquacopper (II) ions are mixed with hydroxide ions, the hydrogen ions on the compound are removed from the water ligands. The resulting complex has no charge and forms a precipitate that is insoluble in water. The complex is light blue in colour [Cu (H20)6]2+ + 2OH- [Cu (H20)4(OH) 2] + 2H20 (Henry, 2002) But when excess ammonia is added to the solution, then a complex royal blue solution is formed as follows. [Cu (H20)6]2+ + 2NH3 [Cu (NH3)4(H2O) 2]2+ + 4H2O When combined with carbonate ions, it forms light blue precipitate as follows Cu2+ (aq) + CO32- (aq) CuCO3 (s) (Henry, 2002) IRON Iron ions Reaction in solution Ferrous iron has two notable oxidation states of +2 and +3, i.e. Fe +2 and Fe 3+. It has also the +6 oxidation state in its FeO42- [ferrate (VI) ion. The known simple ions that are in solution form are the hexaaquairon (III) ion- [Fe (H20)6]3+ and the hexaaquairon (II) ion- [Fe (H20)6]2+. They both possess acidic properties, but the Iron (III) is stronger than the Iron (II). (Henry, 2002) Reactions of ions of iron with hydroxide ions When ferrous iron is combined hydroxide ions, say from sodium hydroxide, the hydrogen ions are removed from the iron ions. When enough of hydrogen ions have been displaced, a neutral complex precipitate with no charge is formed and is insoluble in water as follows Iron (II) ions [Fe (H20)6]2+ 2OH- [Fe (H20)4(OH) 2] + 2H2O The resulting solution is green in colour. On standing, the precipitate is easily oxidized in alkaline conditions. The oxygen in the air causes the iron (II) hydroxide to be oxidized to iron (III) hydroxide on the top of the solution. This makes the precipitate to darken and turn orange near the top of the test tube (Soloveichik, 1993) Iron (III) ions [Fe (H20)6]3+ 3OH [Fe (H20)4(OH) 2] + 2H2O The resulting solution is orange/brown in colour. (Soloveichik, 1993) Reaction with carbonate ions Iron (II) ions with carbonate ions Fe2+ (aq) + CO32+ (aq) FeCO3 (s) The resulting complex is a green precipitate of carbonate. Iron (III) irons combination with carbonate If hexaaquairon (III) ions are combined with carbonate ions, a neutral complex results together with carbon dioxide gas. 2 [Fe (H2O) 6] 3+ + 3CO32+ 2[Fe (H2O) 3(OH) 3] + 3CO2 + 3H2O CHRONIUM The simplest ion formed by Chromium is hecaaquachronium (III) ion, [Cr (H20)6]3+. This ion complex is a difficult to describe colour that is violet-blue-grey. When produced in a test tube reaction it is usually green. Reaction of chromium (III) ions with hydroxide ions The hydroxide ions from sodium hydroxide remove the hydrogen ions from chromium ion solution leaving a neutral complex with no charge. Below is the reaction equation: [Cr (H20)6] + 3OH [Cr (H20)3(OH) 3] + 3H20 However, this reaction process doesn’t end here. More ions of hydrogen are removed to relinquish ions such as [Cr (OH) 6]3+ and [Cr (H20)2(OH) 4]-. For instance; [Cr (H20)3(OH) 3] + 3OH- [Cr (OH) 6]3-+ H20 As the precipitate redissolves, the colour changes from greyish to dark green colour under excess sodium hydroxide solution. Hexaaquachronium (III) ions reaction with Ammonia In this case, ammonia acts as a ligand and as a base. When ammonia solution is added to the complex, hydrogen ions are removed to form a neutral complex that is greyish as follows [Cr (H20)6]3+ + 3NH3 [Cr (H20)3(OH) 3 +3NH4+ When excess ammonia is added to the precipitate, the water ligand is replaced by ammonia to give the hexaamminechromium (III) ions. The following reaction equation serves to illustrate this. The resulting precipitate is purple in colour. [Cr (H20)6]3+ + 6NH3 [Cr (NH3)6]3+ + 6H20 When hydrogen peroxide is added to the hexaaquachromium (III) ions solution with a little, the resulting solution is yellow in colour MANGANESE Manganese (II) ions reaction in solutions Hydroxide ions from sodium hydroxide react to remove the hydrogen ions from the manganese ion complex resulting into a neutral complex with no charge. The precipitate is insoluble in water. [Mn (H2O) 6]2+ + 2OH- [Mn (H2O) 4(OH) 2] + 2H2O (Soloveichik, 1993) The precipitate is pale pink but on standing, it’s oxidized to dark brown at the top of the test tube to form manganese (III) oxide. When the hexaaquamanganese (II) ions is combined with ammonia solution, a pale brown precipitate is formed which is later oxidized to a darker manganese (III) oxide when exposed to air. [Mn (H2O)6]2+ +2NH3 [Mn(H2O)4(OH)2] +2NH4+ (Soloveichik, 1993) COBALT Cobalt forms hexaaquacobalt (II) ion in solution of [Co (H2O)6]2+. When this simple ion is combined with hydroxide ions from sodium hydroxide, the hydrogen ions are removed and appended to the cobalt ion. The resulting neutral complex ion precipitate has no charge and is insoluble in water. Below is the reaction path [Co (H2O) 6]2+ + 2OH- [Co (H2O) 4(OH) 2] + 2H2O VANADIUM Vanadium has different oxidation states of +5, +4, +3, and +2. Perhaps to observe the different characteristics through which vanadium changes, the reduction of Vanadium (V) to Vanadium (II) is the best to illustrate the process. The source of Vanadium +5 oxidation state will be NH4 VO3. This is insoluble in water but is soluble in sodium hydroxide. The reaction is carried out under acidic conditions at a time when dioxovanadium ion is present. The reaction is warmed whereby the + 5 oxidation stage is yellow, the second stage is green ( a combination of the former yellow and the resulting blue) and lastly the oxidation state +4 is light blue in colour. As the stage progresses, oxidation state +3 {V (H20)6 3+} is blue, whereas oxidation state +2 {V (H20)6 2+is purple. When Vanadium (II) is oxidized (when it comes into contact with air), it turns into a green Vanadium (III) ion. (Dolgoplosk, 1983) Reasons why different Oxidation States would give different colours Transition metals complex ions exhibit varied colours while the non-transition metals don’t. The partially filled d orbitals are responsible for this scenario (Dolgoplosk, 1983) Octahedral Complexes When the transition metals bond with the ligands, repulsion occurs between the d orbitals electrons and the ligands which consequently steps up the energy of the d orbitals. Because of the characteristic arrangement of the d orbitals in space, the energy is not stepped in a similar manner, but splits into two groups. Below is the diagram showing how Cu2+ d electrons are arranged before and after bonding with six water molecules (Churchill, 2000). The moment 6 ligands are put around an ion of a transition metal, the d orbital split into two groups, whereby two have higher energy than the remaining three. The difference between their energy is represented by the blue line to the left of the diagram above. This gap/differences depends on the transition metal ion, ligands nature and the oxidation state (for instance 2+ or 3+) When the ion solution is exposed to white light, some energy is utilized to promote the electron from the lower orbitals to the upper set space. Every wavelength has a specific energy that accompanies it. Red light depicts the lowest amount of energy within the visible region. On the other hand violet has got the highest energy. Now suppose the complex ion d orbitals energy gap corresponds to the yellow light energy. In this case the yellow light would be absorbed just because it will utilize in promoting the electron. The other colours are left intact. The human eye would behold a dark blue colour because of its complimentarity with yellow. (Dolgoplosk, 1983) Tetrahedral Complexes The common tetrahedral complexes do possess 4 ligands arranged around the complex ion. They have an effect on the complex ion d electrons. The arrangements of the ligands and d orbitals shapes are very different. The overall effect being that when the d orbitals split, three have greater energy than the other two (opposite arrangement in the octahedral complex.) Apart from the difference in details, the explanation of colour origin in octahedral complexes with respect to light wavelengths is same as other complexes (Churchill, 2000) The Ligand Nature Some ligands may possess strong electrical fields and cause high energy gaps with respect to the d orbitals when they are split into groups. On the other hand, some have weaker electrical fields producing smaller gaps. Depending on the size of the gap, a particular light wavelength will be absorbed. Below is a demonstration of their strength: (Dolgoplosk, 1983). Conventionally, the higher the splitting, the more energy is required to promote lower groups orbital electrons to the higher levels. In this respect the greater the energy, the shorter the light wavelengths. For instance, is excess ammonia is added to the hexaaquacopper (II) ions, the pale blue colour is displaced by the dark blue colour because some water molecules are replaced by ammonia in the complex ion CI- F- OH- H20 NH3 CN- The first complex absorbs the red light to give the complimentary cyan colour, while the second one absorbs the yellow region to give the complementary dark blue colour. In this case the yellow light has got more energy than the red light. Ammonia increases splitting of d orbitals as compared to water. Chromium (III) ions- [Cr (OH) 6]3+ will form the dark green and purple colours because of the changing colors. (Churchill, 2000) Differing Oxidation States When the oxidation state of the transition metal increases, the volume of splitting of the d orbitals also enlarges. Changes in an elements’ oxidation state changes the type of colour absorbed which is made visible to the eye. For instance as chromium (II) ions are oxidized, they form bluish violet colour complex of chromium (III) ions. In this case the splitting is greater if the complex is octahedral than tetrahedral. An example is the Cobalt (II) ion, [Co (H2O) 6]2+ and the [CoCI4]2-. In this case the change in the colours will be as a result of the change of ligands number and the change in itself. (Dolgoplosk, 1983) Manganese is another element that exhibit different oxidation states. It has + 2 in Mn 2+, + 3 in Mn2O3, +4 in MnO2 , +6 in Mn O42- , +7 in MnO4-. References Churchill, F (2000). Why Transition Metal exhibit varied Colours. Amsterdam: Longhorn Publishers. Dolgoplosk, B.A ( 1983). The Mechanism of the Decomposition of Orgarnometallic Compounds of Transition Metals and the Role of Intermediate Species in Catalysis. London: Prentice Hall Henry, T. (2002). Transition Metals: A Chemistry Analysis. Singapore: Springer Publishers. Robert, Y ( 2004). Oxidation States of Transition Metals. London: Prentice Hall Soloveichik, G. L (1993). Bimetallic Transition Metals Hydride Complexes. New York: Macmillan. Read More