The d-block elements fall in the Periodic table between the s-block and the p-block. They are often called the Transition Elements. Only the first row of transition elements, scandium to zinc will be considered. Scandium and zinc are not classed as transition elements as they show some fairly obvious differences from the elements titanium to copper. They have only one oxidation state, their compounds are white, and they do not behave as catalysts.
The electronic configurations for the first row of transition elements is:
|Element||Atomic Number||Electronic Configuration|
We can define a transition element as one which forms at least one ion with a partially filled d sub-shell. There are five separate d-orbitals each of which can accommodate two electrons of opposite spin. These five orbitals have the same energy. In practice each of these orbitals must be singly occupied by an electron before pairing takes place. When all five 3d orbitals are either singly or doubly filled a degree of stability is conferred on the atom or ion. This explains the electronic configurations of the atoms chromium and copper which are respectively 3d54s1 and 3d104s1. It also explains why Fe2+ (3d6) is easily oxidised to Fe3+ (3d5) but Mn2+ (3d5) is not readily oxidised to Mn3+ (3d4).
The transition elements show a horizontal similarity in their physical and chemical properties as well as the usual vertical relationship. The horizontal similarity contrasts sharply with the trend in traversing a row of the s and p block elements, e.g. Li to F, where it is the difference between them that is the most striking characteristic. This is because in the transition elements, although the nuclear charge is increasing by one, an electron is being added to an inner d-orbital. In Li to F a valency electron is being added. As an electron is being added to an inner orbital, the difference from element to element is only small. Ionisation energies and electronegativities tend therefore to increase only slightly along the series.
The transition elements are silvery metals (apart from copper). They are dense metals with high melting points and boiling points. They tend to be hard with high tensile strength and good mechanical properties. They are good conductors of heat and electricity and most have a close-packed structure of atoms. They are much less reactive than the s-block metals.
Strong forces exist between the separate atoms in a metal and these are known as metallic bonds. This can be explained by the movement of the outer shell electrons of the metal, as they move randomly throughout a lattice of regularly spaced positive ions. The moving electrons are referred to as a 'sea of electrons'. Each positively charged ion is attracted to the sea of electrons and vice versa. These electrostatic attractions bind the structure together.
When a force is applied to a metal, the layers of atoms can slide over each other in a process known as slip. After slipping the atoms settle once again into a close-packed structure. This is why metals can be hammered into different shapes (malleable) or drawn into wire (ductile). In an alloy, differently sized atoms interrupt the orderly arrangement of atoms in the lattice and make it more difficult for the layers to slide over each other.
The transition elements have certain characteristic properties:
Transition elements have electrons of similar energy in both the 3d and 4s levels. Successive ionisation energies show a gradual increase as first the 4s and then the 3d electrons are removed. This means that one particular element can form ions of roughly the same stability by losing different numbers of electrons. Therefore, all transition metals from titanium to copper show two or more oxidation states in their compounds. When a first row d-block element loses electrons to form positive ions, the 4s electrons are lost first, followed by the 3d electrons. In no case are electrons lost from the 3p orbitals.
The common oxidation states for each element include +2, +3 or both. +3 states are relatively more common at the beginning of the series whereas +2 states are more common towards the end. Up to manganese the highest oxidation states involve all 3d and 4s electrons. After manganese, there is a decrease in the number of oxidation states shown by each element. The lower oxidation states are found in simple ionic compounds, e.g. Cr3+, Mn2+, Fe3+. The higher oxidation states give rise to covalent compounds, e.g. MnO4-, Cr2O72-, or complex compounds.
An effective demonstration of the range of oxidation states of a d-block element can be shown by reacting a solution of ammonium metavanadate (vanadium V) with zinc and hydrochloric acid.
Examination of electrode potential values shows that it is possible to reduce vanadium (V) to vanadium (IV) to vanadium (III) and finally to vanadium (II).
The solution of ammonium metavanadate is yellow due to the presence of dioxovanadium (V) (VO2+) ions in acid solution. When this solution is shaken with zinc, it changes gradually through green to blue oxovanadium (IV) (VO2+), then to green vanadium (II) ions and eventually to violet vanadium (II) ions.
A catalyst is a substance which alters the rate of a chemical reaction without being consumed by the process. A small amount of catalyst is able to catalyse the reaction of a large amount of reactant. A catalysed reaction has a lower activation energy than an uncatalysed reaction. It is thought that a catalyst provides a new reaction pathway (mechanism) of lower activation energy. There are two types of catalysts:
Heterogeneous Catalysts. The catalyst is in a different phase from the reactants. This usually means a solid metal catalyst with the reactants in the gas or liquid phase.
When gas molecules approach the surface of a solid catalyst there is a tendency for them to interact with the surface atoms. This interaction is called adsorption. If new bonds are formed between the reactant molecules and the catalyst, it is called chemisorption. If only a weak interaction occurs it is called physical adsorption.
When ammonia is synthesised from hydrogen and nitrogen in the Haber Process, the hydrogen and nitrogen first diffuse towards the surface of the iron catalyst and are then adsorbed on the surface. The nitrogen and hydrogen then react on the surface of the catalyst. The ammonia then desorbs from the catalyst surface and then diffuses away.
Homogeneous Catalysts. The catalyst is in the same phase as the reactants. This usually means that they are in the aqueous phase. Often the transition metal ion forms an intermediate compound with one or more reactants which then breaks down to form the products. Catalytic activity is usually associated with variable oxidation states of the catalyst. This is illustrated with the reaction between 2,3-dihydroxybutanoate ions and hydrogen peroxide.:
It is thought that the catalyst Co2+(pink) ions reduce the hydrogen peroxide to form an intermediate containing Co3+ (green). The Co3+ oxidises the 2,3-dihydroxybutanoate before reducing itself back to Co2+.
A complex compound is formed when a number of molecules or negatively charged ions surround a central d-block atom or ion. These surrounding molecules or ions possess a lone pair of electrons and are called ligands. They form dative covalent bonds with the metal. The number of bonds from ligands to the central ion is known as the coordination number. The complexes formed may have an overall positive charge, a negative charge or no charge, e.g. [Fe(H2O)6]3+, [NiCl4]2-, Ni(CO)4.
The shape of a complex depends on its coordination number.
Complexes with coordination number 6 usually have an octahedral shape, e.g. [Fe(CN)6]3-, hexacyanoferrate (III) ion.
Complexes with coordination number 4 usually have a tetrahedral shape, e.g. NiCl42-. Some four-coordinate complexes have a square planar shape, e.g. Ni(CN)42-.
Complexes with coordination number 2 usually have a linear shape, e.g. Ag(NH3)2+.
A systematic procedure is used.
|[Cr(H2O)6]3+||hexaaquachromium (III) ion|
|[Fe(CN)62-||hexacyanoferrate (II) ion|
|[CuCl4]2-||tetrachlorocuprate (II) ion|
Different ligands form complexes with different stabilities. Ammonia displaces water ligands when added to blue aqueous copper (II) ions. The water is displaced until a deep blue complex is formed.
The stability of a complex is expressed in terms of an equilibrium constant, called the stability constant, Kstab. Usually logarithmic values are used, log10 Kstab. The values quoted give the stability of the complex relative to the aqueous ion. The higher the value of log10 Kstab, the more stable the complex.
Ligands that occupy one coordination position are called monodentate, e.g. Cl-, H2O, NH3, CN-. Those that can occupy more than one position are called polydentate ligands. 1,2-diaminohexane, NH2CH2CH2NH2, occupies two coordination positions and is said to be bidentate. EDTA occupies six positions and is said to be hexadentate. Where the metal ion is held in a ring, this is called a chelate ring.
Compounds of d-block metals are frequently coloured in the solid state and in solution. The colour of the ion depends upon the ligands bonded to the metal ion. The pale blue [Cu(H2O)6]2+ changes to dark blue in the presence of excess ammonia and to green if sufficient chloride ions are added. Hydrated cobalt(II) ions are pink but in the presence of excess chloride ions the blue complex CoCl42- is formed.
The colour of a transition metal complex depends on:
When electrons move from one energy level to a higher one, they absorb a quantum of electromagnetic energy equal in energy to the gap between the levels. The frequency, , is given by the equation E = h . If is in the visible region of the spectrum, this will result in the substance being coloured. If a substance absorbs green light, it will let through red and blue and will appear purple. Unpaired electrons absorb light energy by becoming promoted from their ground-state energy levels to their excited-state energy levels. The wavelength of the light absorbed depends on the energy difference E between these states. This energy difference also depends on the nature of the ligands in the complex.
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