We are familiar that the periodic table positioned some elements between the group II and group III are known as transition elements. Transition or d-block elements are so called as their position in the periodic table is between the s and p block elements, and their properties are transitional among the highly reactive metallic elements of the s-block that usually form ionic compounds, and the elements of the p-block that are largely non-metallic forming covalent compounds.
In the d-block the penultimate shell, that is, the orbital subsequent to the outermost orbital, of electrons is being expanded from eight to eighteen via the addition of d-electrons. The word transition element is only used for an element that consists of partially or incompletely filled d-orbitals. They occupy three rows or periods; (Periods 4, 5 and 6 of the periodic table) having ten elements in each and every row. All the transition elements are metals of great economic significance. The table illustrates the d-block transition metals.
Table: The transition elements:
Period 4: 1st transition series Sc Ti V Cr Mn Fe Co Ni Cu Zn
Period 5: 2nd transition series Y Zr Nb Mo Tc Ru Rh Pd Ag Cd
Period 6: 3rd transition series La Hf Ta W Re Os Ir Pt Aw Hg
Number of electrons 1 2 3 4 5 6 7 8 9 10
The first transition series metals from scandium (Sc) to zinc (Zn) in period 4 are of main focus and they are principally significant in industry. Metals in this series comprise titanium (Ti), vanadium (v), chromium (Cr), Manganese (Mn), Iron (Fe), cobalt (Co), nickel (Ni) and copper (Cu). The transition metals encompass very identical properties and are quite dissimilar from the reactive s-block metals of Groups I and II.
Electronic Configuration of 1st transition series metals:
We are familiar that the energy level diagram of atomic orbitals is ls < 2s < 2p < 3s < 3p < 4s < 3d < 4p. Therefore for elements having atomic numbers 1 to 18, the electrons are filled usually to the orbitals. This ends by Argon whose electronic configuration is 1s2 2s2 2p6 3s2 3p6. The subsequent element, potassium (that is, atomic number 19) and 19th electron goes to the 4s orbital which is of a lower energy than the 3d. For calcium (that is, atomic number 20), the 20th electron goes to complete the 4s orbital. Next is scandium (that is, atomic number 21), the first transition element, the 21st electron goes to the 3d orbital, that is the next available orbital in terms of the energy level diagram. The filling of the 3d orbital continue from scandium to zinc giving the first transition series. Therefore, the outer 4p orbitals of the elements from scandium to zinc are not filled due to their higher energy levels, whereas the inner or penultimate 3d orbitals are preferentially filled as they are at lower energy levels. The table shown below represents the electronic configuration of s-block and d-block metals of Period 4.
Table: Electronic configuration of s- and d- block metals of Period 4
Chromium and copper atoms comprise only one electron in their 4s orbitals. This is as an outcome of the special stability related by half or fully filled sub levels. The 3d5 and 3d10 electronic configurations of chromium and copper correspondingly illustrate that the five 3d orbitals are singly or doubly filled.
Physical Properties of 1st Transition Series Metals:
1) Each and every transition element is typical metals having high boiling and melting points. They are hard, dense and lustrous example: Iron, copper and zinc. They encompass good mechanical properties - malleable and ductile.
2) They are usually good conductors of heat and electricity example: copper and iron.
These properties of transition metals can be illustrated in terms of strong metallic bonding. In addition to the 4s electrons, the 3d electrons in the atoms of such metals are available for metallic bonding. The additional electrons make the metallic bonds in the transition elements extremely strong. In comparison, the corresponding s-block metals, potassium and calcium, in the similar period are soft and of lower melting point. The metallic bonds in s-block metals, which encompass merely one or two available electrons for bonding, are not as strong as those in d-block metals, with extra 3d electrons.
The atomic size of d-block metals as well influences their properties. In comparison, both the s-block and d-block metals in period 4 encompass atoms having the same number of principal energy levels (n = 1, 2, 3, 4) however the transition metals encompass larger atomic number. As an outcome, the atoms of the transition metals are smaller than those of the s-block metals. For illustration, Iron (having atomic number 26) atom by a +26 nuclear charge will encompass a stronger attraction on its electrons and pull them closer (decrease atomic size) than potassium (atomic number 19) having a +19 nuclear charge. This difference is reflected in their atomic sizes - iron (0.12 nm) and potassium (0.20 nm).
The combined effects of strong metallic bonds and small atomic sizes accounts for the high melting and boiling points, densities and tensile strengths of the transition metals. The melting points of manganese and zinc are comparatively lower as their atoms have half-filled and fully-filled d-orbitals correspondingly. This electronic arrangement give stability to the atoms, therefore the d-orbital electrons are not available for metallic bonding. The table given below summarizes several properties of the first transition metal series.
Fig: properties of the first transition series metals
The chemical reactivity of elements is mainly based on their outermost electrons. In the s-block and p-block the chemical properties of the elements in a given period differ (example: from metal to non-metals) from left to right. For the transition metals this doesn't occur simply as electrons are added gradually to the inner d-orbitals and not the outermost orbitals, as in the s-block and p-block elements. The nuclei of transition metals encompass a stronger pull on their electrons than the nuclei of s-block metals. The s-block metals, thus, encompass lower ionization energies and are more reactive than the transition metals.
i) They don't react with the cold water such as the alkali metals. They though react with steam at high temperatures.
3Fe(s) + 4H2O (g) → Fe3O4 + H2
Zn + H2O → ZnO + H2
ii) They are oxidized whenever exposed to air or heated to red hot. In the damp air, iron rusts to form a reddish brown layer of hydrated iron (III) oxide - this can be noticed on iron roofing sheets of our houses.
4Fe + 3O2 → 2Fe2O3 + 3H2O (rust)
iii) Iron, cobalt and nickel react by dilute mineral acids example: HCl, H2SO4, to give hydrogen gas as they are higher than hydrogen in the electrochemical series.
Fe + 2HCl → FeCl2 + H2
Each and every transition metals react with concentrated acids to form the corresponding salt. Copper reacts by either dilute or concentrated trioxonitrate (V) acid to form blue copper (II) trioxonitrate (IV) solution
Cu + 2HNO3 → Cu (NO3)2 + H2
Variable Oxidation States:
One of the most striking characteristics of the transition metals is that they show variable valency. Most of the metals exhibit a broad range of valences example: Fe3+ and Fe2+; Cu2+ and Cu+; V3+ and V4+. This is in contrast to the s-block metals, where the single valency for all time equivalents the group number. The variable oxidation states illustrated by transition metals are due to the 3d electrons are available for bond formation. The word oxidation state is generally preferred to valency. All the metals illustrate an oxidation state of +2 that can be related by the loss of the two outermost s-electrons from the neutral atom.
Fig: oxidation states shown by the transition metals
Therefore Sc could encompass an oxidation number of (+2) if both 's' electrons are employed for bonding and (+3) whenever two 's' and one 'd' electrons are comprised. Ti consists of an oxidation state (+2) whenever both 's' electrons are employed for bonding, (+3) whenever two 's' and one 'd' electrons are utilized and (+4) whenever two 's' and two 'd' electrons are employed. Likewise, V illustrates oxidation numbers (2), (3), (4) and (5). In case of Cr, by employing the single 's' electron for bonding we obtain an oxidation number of (+1); therefore by employing varying numbers of 'd' electrons oxidation states of +(2), (3), (4), (5) and (6) are possible. Mn consists of oxidation states (2), (3), (4), (5), (6) and (7).
We are familiar that the oxidation state of any element can be worked out from the molecular formula of its compounds. For illustration, the oxidation state of Mn in KMnO4 is:
K+ = +1
4O2- = 4 x (-2) = -8
As it is neutral, positive and negative charges should balance; therefore remaining positive charge is +7
That is, + 1 + x - 8 = 0
x = 8 - 1 = +7
This is the oxidation state of Mn.
The lower oxidation states of the transition metals take place in ionic compounds and tend to form reducing agents example: Cr2+ salts, and basic oxides example: Mn2+O2-. The higher oxidation states are found in the covalent compounds (through sharing electrons) and tend to form oxidizing agents example: K2Cr2O7 → Cr (VI) and acidic oxides example: Mn2O7 → Mn (VII). Compounds of transition metals having intermediate oxidation states form amphoteric oxides.
Complex ion formation:
Transition metals form numerous coordination compounds, complex ions, in contrary to the s- and p- block elements. A complex ion consists of a central positive ion linked through a cluster of some other ions or molecules termed as ligands. The number of ligands is termed to as the coordination number. The bonding between the central metal ion and the ligands might be either predominantly electrovalent or predominantly coordinate. The ligands act as electron donors. Illustrations of ligands are: NH3, H2O: CN.
For instance, the blue complex ion, tetraammine copper (II) ion [Cu (NH3)4]2+ in which the central ion, Cu2+, is coordinately bonded to four ammonia a molecules as ligands.
Fig: Tetraammine copper (II) ion
Color of Transition Metal Ions:
Ionic and covalent compounds of transition metals are generally markedly colored, in contrary to compounds of the s-and p- block elements that are almost for all time colourless. Color formation is related with the capability to promote an electron from one energy level to the other on absorption of energy. The color of transition metal ions generally serve up as a helpful guide in recognizing a compound.
For the first transition series metals, the colors are related by partially filled 3d orbitals (that is, 3d1 to 3d9). Scandium and zinc ions are colourless as they don't encompass partially filled 3d orbitals. The table represents the characteristic colors of transition metal ions.
Table: Color of the 1st transition series metal ions in water.
Catalytic capability is one of the pronounced features of the transition metals. The ease with which ions of transition metals change their oxidation states allow them to act as the catalysts. The most general catalyst and their uses are illustrated below:
TiCl4 (Ziegle-Natta catalyst) Polymerization of ethene to polyethene
Ni Hydrogenation of vegetable oil to margarine
MnO2 Decomposition of H2O2 to water and oxygen
V2O5 Contact method for H2SO4 manufacture
Fe Haber method for NH3 manufacture
Industrial Uses of Transition Metals:
1) Iron is employed in the manufacture of steel that are used in the making of tin-plates, corrugated sheets, car bodies, screws, nails, pipes and so on.
2) Copper is broadly employed for making electric wires, metal-work because of its attractive appearance, and in alloys like bronze and brass.
3) Zinc is employed broadly for galvanizing iron and steel, in alloys like bronze, brass and in making dry cells and package foils.
4) Some of them are employed as catalyst for significant industrial process and to impact color to objects example: ceramics, paints and so on.
The features of metals can be enhanced by mixing them with other elements to form alloys. The alloy is a substance made by adding one or more elements to a base or parent metal to get desired properties. The added elements are generally metals or carbon.
Alloys are economically significant as most of the metallic things we use are manufactured of alloys and not the pure metals. The table illustrated below gives some common alloys, their compositions and uses.
Table: Properties of alloys
Hazards of Radioactivity:
Gamma radiation is utilized to destroy the cancerous cells. This is one utilize gamma radiation is put into F-radiation as well destroys healthy cells and also too much exposure to it can do more harm than good. The degree of damage based on the energy and kind of radiation. The effect of radiation is as well cumulative and small doses over a long period of time will as well cause serious damage to biological systems. Radioactive waste is extremely dangerous and should be disposed properly to avoid unnecessary exposure to its risks.
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