UNIT 09 THE d and f-BLOCK ELEMENTS
Syllabus
· Position in the Periodic Table
· Electronic Configuration of Transition Elements
· Properties of Transition Elements
· Comparison of the First Row Transition Metals through the d-Electron Configuration
· General Group Trends in the Chemistry of the d-block Metals
· Occurrence of Principles of Extraction of Some d-Block Metals
· Some important Compounds of Transition Elements
· The Lanthanides
· The Actinoids
In the periodic classification of elements, the elements have been divided into four blocks. This division is based on the type of atomic orbital which receives the last electron in the atom. The d-block elements are called transition elements. This block consists of elements lying between s and p-blocks i.e., between Groups 2 and 13 , starting from fourth period on wards. In these elements, the outermost shell contains one or two electrons in their s-orbital (ns) but the last electron enters the last but one d-subshell i.e., (n - 1) d. The elements of this block have general characteristic properties intermediate between the elements of s-block and p-block. In other words d-block elements represent change (or transition) in properties from most electropositive s-block elements to least electropositive p-block elements. Therefore, these are called Transition Elements.
Definition and Electronic Configurations
The transition elements are those elements which have partially filled d-subshells in their elementary form or in their commonly occurring oxidation states.
The definition covers coinage metals, Cu, Ag and Au as transition metals because in their commonly occurring oxidation states , they have partially filled d-subshells. For example, Cu2+ has 3d9 configuration, Au3+ has 5d8 configuration, although all these atoms have completely filled d-subshell in their elementary states(Cu ; 3d10 : Ag ; 4d10 : Au ; 5d10 ). However, it may be noted that, the above definition does not include elements of Group-12 i.e., Zn, Cd and Hg. These elements do not have partially filled d-subshells in their elementary state or in their commonly occurring oxidation states. However, these may be treated along with transition elements because they are quite similar to other transition elements in some of their chemical properties.
Electronic configurations
In the transition elements , the d-orbitals are successively filled. The general electronic configuration for atoms of d-block is : (n-1) d1-10 ns1-2
Transition Series
The transition elements consist of three complete rows of the elements and one incomplete row. These rows are called first, second and third transition series which involve the filling of 3d, 4d and 5d-orbitals respectively. These series are also called Transition series.
First transition Series
Scandium (Z = 21) to Zinc ( Z = 30). 3d-orbitals are gradually filled.
Second transition Series
Yttrium(Z=39) to cadmium ( Z = 48). 4d-orbitals are gradually filled.
Third transition Series
Lanthanum (Z = 57) , Hafnium( Z = 72) to mercury (Z = 80) . 5d-orbitals are gradually filled.
Fourth transition Series
This series starts from Actinium(Z=89) and is incomplete. It will include elements from atomic number 104 onwards (only some of these are isolated).
Electronic configurations
The first transition series consists of elements from Scandium (Z=21) to Zinc (Z=30). In Scandium, the 3d-orbital starts filling up and its electronic configuration is [Ar]4s23d1. As we move from scandium onwards, 3d-orbitals get filled up more and more till the last element, zinc, in which the 3d-orbitals are completely filled, i.e., [Ar]4s23d10.
Element | Symbol | Atomic number | Outer Electronic configuration |
Scandium | Sc | 21 | 3d14s2 |
Titanium | Ti | 22 | 3d24s2 |
Vanadium | V | 23 | 3d34s2 |
Chromium | Cr | 24 | 3d54s1 |
Manganese | Mn | 25 | 3d54s2 |
Iron | Fe | 26 | 3d64s2 |
Cobalt | Co | 27 | 3d74s2 |
Nickel | Ni | 28 | 3d84s2 |
Copper | Cu | 29 | 3d104s1 |
Zinc | Zn | 30 | 3d104s2 |
It may be noted that the configuration of chromium and copper are anomalous. The half-filled and completely filled electronic configurations (i.e., d5 and d10) have extra stability, one of the 4s-electron goes to nearby 3d-orbital , so that 3d-orbitals become half-filled in case of chromium and completely filled in the case of copper respectively. Therefore the electronic configuration of chromium is [Ar]3d54s1 rather than [Ar]3d44s2 while that of Cu is [Ar]3d104s1 instead of [Ar]3d94s2.
Second transition Series
The second transition series consist of elements from Yttrium Y (Z= 39) to cadmium(Z=48) i.e., Yttrium(Z=39), zirconium (Z=40), niobium(Z=41), molybdenum (Z=42), technetium (Tc = 43), ruthenium (Z = 44), Rhodium (Z=45) , Palladium (Z = 46), silver (Z = 47) and cadmium (Z= 48) . This series involves the filling of 4d-orbitals (TABLE)
Electronic configuration of Second Transition Series
Element | Symbol | Atomic number | Outer Electronic configuration |
Yttrium | Y | 39 | 4d15s2 |
Zirconium | Zr | 40 | 4d25s2 |
Niobium | Nb | 41 | 4d45s1 |
Molybdenum | Mo | 42 | 4d55s1 |
Technetium | Tc | 43 | 4d55s2 |
Ruthenium | Ru | 44 | 4d75s1 |
Rhodium | Rh | 45 | 4d85s1 |
Palladium | Pd | 46 | 4d105s0 |
Silver | Ag | 47 | 4d105s1 |
Cadmium | Cd | 48 | 4d105s2 |
Third transition Series
The series consists of elements Lanthanum( Z = 57) and from Hafnium (Z=72) to mercury (Z=80) i.e., Lanthanum (La), Hafnium(Hf), Tantalam(Ta), Tungsten (W), Rhenium(Re), Osmium(Os), Iridium(Ir), Platinum(Pt), Gold(Au) and Mercury(Hg). In between Lanthanum and Hafnium, there are fourteen elements called Lanthanides which involve the filling of 4f-orbitals and do not belong to this series. The elements of this series involve the gradual filling of 5d-orbitals.
Element | Symbol | Atomic number | Outer Electronic configuration |
Lanthanum | La | 57 | [Xe] 5d16s2 |
Hafnium | Hf | 72 | [Xe] 4f145d26s2 |
Tantalum | Ta | 73 | [Xe] 4f145d36s2 |
Tungsten | W | 74 | [Xe] 4f145d46s2 |
Rhenium | Tc | 75 | [Xe] 4f145d56s2 |
Osmium | Os | 76 | [Xe] 4f145d66s2 |
Irridium | Ir | 77 | [Xe] 4f145d76s2 |
Platinum | Pt | 78 | [Xe] 4f145d96s1 |
Gold | Au | 79 | [Xe] 4f145d106s1 |
Mercury | Hg | 80 | [Xe] 4f145d106s2 |
It may be noted in the second and third transition series, there are many anomalous configurations in comparison to those of the first transition series. These are attributed to factors like nuclear-electron and electron-electron forces.
Fourth transition series
It involves the filling of 6d-subshell starting from actinium (Z=89) ; which has the configuration 6d17s2. This series is incomplete.
Element | Symbol | Atomic number | Outer Electronic configuration |
Actinium | Ac | 89 | [Rn] 6d17s2 |
Rutherfordium | Rf | 104 | [Rn] 5f146d27s2 |
Dubinium | Db | 105 | [Rn] 5f146d37s2 |
Seaborgium | Sg | 106 | [Rn] 5f146d47s2 |
Bohrium | Bh | 107 | [Rn] 5f145d57s2 |
Hassnium | Hs | 108 | [Rn] 5f145d67s2 |
Meitnerium | Mt | 109 | [Rn] 5f146d77s2 |
Dasmstarium | Ds | 110 | [Rn] 5f146d87s2 |
Uuu | 111 | ||
Uub | 112 |
It would be of interest to consider the change in the relative energies of various sub-shells with change in atomic number. The relative energies of subshells of third , fourth and fifth energy shells are given in Fig. As is clear from the figure , upto Z = 20, 4s-sub-shell has lower energy than 3d-subshell. Therefore calcium(Z=20) has the electronic configuration 1s22s22p63s23p64s2. Beyond this there is a sharp decrease in energy of 3d-subshell and its energy becomes less than the 4s-subshell.
Fig 1 Energies of orbitals of third, fourth and fifth energy levels as a function of atomic number.
Starting with scandium(Z=21) upto zinc(Z = 30) , the 3d orbitals are filled and the d-electrons become more effective in shielding the 4s-electrons from attractive force of the nucleus. Therefore, 3d-subshells get pulled lower than 4s.
The transition elements with partly filled d-orbitals exhibit certain characteristic properties. For example, they display a variety of oxidation states, form coloured ions and enter into complex formation with different anions and neutral molecules. These metals and their compounds also exhibit catalytic property and many of them show paramagnetic behaviour. There are greater horizontal similarities in the properties of transition elements in contrast to the main group elements.
Ionisation of atoms
It is observed that after calcium, the energy of 4s-subshell is more than 3d-subshell (Fig above) . Therefore , when electrons are to be removed from these atoms, the electrons will go from 4s-orbital rather than 3d-orbital , though the former was filled earlier. For example.
Fe ( Z = 26) : [Ar ] 3d6 4s2
Fe3+ : [Ar] 3d5
Ni (Z = 28) : [Ar] 3d84s2 Ni2+ : [Ar] 3d8
The same trend is repeated for 5s, 4d and 6s, 5d orbitals in second and third transition series. In lanthanum, the energies of 4f, 5d and 6s-orbitals lie very close to one another and one electron goes to 5d-orbital before 4f-orbitals. Its configuration is 5d16s2. In other elements 4f-orbitals are stabilised relative to 5d and electron go into 4f-orbitals.
Problem
01. On what ground can you say that scandium ( Z = 21) is a transition element but zinc ( Z = 30) is not ?
PROPERTIES OF TRANSITION ELEMENTS
Physical Properties
Nearly all the transition elements display typical metallic properties such as high tensile strength, ductility, malleability , high thermal and electrical conductivity and metallic lustre. Except mercury which is a liquid at room temperature , other elements have typical metallic structures.
Their melting and boiling points are high. The Fig 2 shows the melting points of 3d. 4d and 5d transition metals.
Fig 2 Trends in melting points of transition elements The high melting points are attributed to strong interatomic bonding , which involves the participation of both ns and (n-1)d electrons. In any row the melting points of these metals rise to maximum at d6 except for anomalous values of Mn and Tc and fall regularly as the atomic number increases.
With the exception of zinc, cadmium and mercury , the transition elements are much harder and less volatile. They exhibit high enthalpies of atomisation (Fig 3 Page 4) ; the maxima at about the middle of each series indicates that one unpaired electron per d-orbital is particularly favourable for strong interatomic interaction. In general, , greater the number of valence electrons , stronger is the resultant bonding. Since the enthalpy of atomisation is an important factor in determining the standard electrode potential of a metal, those with highest enthalpy of atomisation (highest boiling point) tend to be noble metals.
The metals of the second (4d) and third (5d) series have greater enthalpies of atomisation than the corresponding elements of the first series ; this is an important factor for their having much more frequent metal-metal bonding in their compounds.
The decrease in metallic radius coupled with increase in atomic mass results in a general increase in density of these elements. Thus , from titanium to copper the increase in the density may be noted (TABLE).
Table 1 Electronic configuration and some properties of first series of transition elements
Element | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn |
Atomic number | 21 | 22 | 23 | 24 | 25 | 26 | 27 | 28 | 29 | 30 |
Ele. configuration M M+ M2+ M3+ | 3d14s2 3d14s1 3d1 [Ar] | 3d24s2 3d24s1 3d2 3d1 | 3d34s2 3d34s1 3d3 3d2 | 3d54s1 3d5 3d4 3d3 | 3d54s2 3d54s1 3d5 3d4 | 3d64s2 3d64s1 3d6 3d5 | 3d74s2 3d74s1 3d7 3d6 | 3d84s2 3d84s1 3d8 3d7 | 3d104s1 3d10 3d9 3d8 | 3d104s2 3d104s1 3d10 3d9 |
Enthalpy of atomisation kJ/mol | 326 | 473 | 575 | 397 | 281 | 416 | 425 | 430 | 339 | 126 |
Ionisation enthalpy I II III | 631 1235 2393 | 656 1309 2657 | 650 1414 2833 | 653 1592 2990 | 717 1509 3260 | 762 1561 2962 | 758 1644 3243 | 736 1752 3402 | 745 1958 3556 | 906 1734 3829 |
r pm M M2+ M3+ | 164 - 73 | 147 - 67 | 135 79 64 | 129 82 62 | 137 82 62 | 126 77 65 | 125 74 61 | 125 70 60 | 128 73 - | 137 75 - |
E° (V) M2+/M M3+/M2+ | - - | - - | - 1.18 - 0.26 | - 0.90 -0.41 | -1.18 +1.57 | - 0.44 +0.77 | - 0.28 +1.97 | - 0.25 - | + 0.34 - | - 0.76 - |
Density (g/cm3) | 3.43 | 4.1 | 6.07 | 7.18 | 7.21 | 7.8 | 8.7 | 8.9 | 8.9 | 7.1 |
Fig 3 Trends in enthalpies of atomisation of transition elements.
Variation in Atomic and Ionic sizes
In general , ions of the same charge in a given series show progressive decrease in radius with increasing atomic number. This is because the extra electron enters a d-orbital each time the nuclear charge increases by unity. The shielding effect of a d-electron is small, the net electrostatic attraction between the the nuclear charge and outermost electron increases and the ionic radius decreases. The atomic radius also decreases in the same way. The variation within a series is quite small. An interesting point emerges when sizes of elements in one series are compared with the sizes of the corresponding elements in other series. The data in Fig 4 shows that an increase from the first (3d) to second (4d) series of the elements but the radii of the third (5d) series are
Fig 4 Trends in atomic radii of transition elements
virtually the same as those of the corresponding members of the second series. This phenomenon is associated with the intervention of the 4f-orbitals, which must be filled before the 5d-series of elements begin. The filling of 4f before 5d orbital results in a regular decrease in atomic radii called lanthanoid contraction which essentially compensate for the expected increase in atomic size with increasing atomic number. The net result of lanthanoid contraction is that the second and third d-series exhibit similar radii (e.g., Zr 160 pm , Hf 159 pm) and have very similar physical and chemical properties.
Ionisation Enthalpies
With increasing nuclear charge, which accompanies the filling of the inner d-orbitals , there is an increase in ionization enthalpy along each series of the transition elements from left to right , but many small variations occur. TABLE 1 gives the ionisation enthalpies of the first row elements. These values show that although the first ionisation enthalpy in general, increases , the increase in the second and third ionization enthalpies for the successive elements are not of the same magnitude. However, the trend is similar for the second ionization enthalpies , which for the most part increase smoothly as the atomic number increases, the exceptions are chromium and copper for which these values are notably larger than those of their neighbours. These exceptions are attributed to the extra stability of half-filled or completely filled set of d-orbitals in chromium and copper. The third ionization enthalpies are quite high and there is a marked break between the values for Mn2+ and Fe2+. Also , the high values for copper, nickel and zinc indicate why it is difficult to obtain oxidation state greater than two for these elements.
Problem
02. Why do transition elements exhibit higher enthalpies of atomization ?
Oxidation states
The transition elements exhibit a large number of oxidation states. With the exception of few elements, most of these show variable oxidation states. These different oxidation states are related to the electronic configuration of their atoms. For example, the oxidation states exhibited by the transition elements of the first series are listed in TABLE.
Different oxidation states of First transition series
Element | Outer electronic configuration | Oxidation states | |
Sc | 3d14s2 | +2,+3 | |
Ti | 3d24s2 | +2,+3,+4 | |
V | 3d34s2 | +2,+3,+4,+5 | |
Cr | 3d54s1 | (+1),+2,+3,(+4),(+5),+6 | |
Mn | 3d54s2 | +2,+3,+4,(+5),+6,+7 | |
Fe | 3d64s2 | +2,+3,(+4),(+5),(+6) | |
Co | 3d74s2 | +2,+3,(+4) | |
Ni | 3d84s2 | +2,+3,+4 | |
Cu | 3d104s1 | +1,+2 | |
Zn | 3d104s2 | +2 | |
*Oxidation states with in the brackets are unstable.
Explanation
The existence of the transition elements in different oxidation states means that their atoms can lose different number of electrons. This is due to the participation of inner (n-1) d-electrons in addition to outer ns-electron because, the energies of the ns and (n-1) d-subshells are almost equal. For example, scandium has the electronic configuration of 3d14s2. It exhibits an oxidation state of +2 when it uses both of its two 4s-electrons for bonding. It can also show oxidation state of +3 when it uses its two s-electrons and one d-electron.
Variable oxidation states of second and third transition series
The elements of the second and third transition series also exhibit variable oxidation states as given in the following TABLE.
Different oxidation states of elements of second and third transition series.
Second Transition series
Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd |
+3 | +3 +4 +2 | +2 +3 +4 +5 | +2 +3 +4 +5 +6 | +2 +4 +5 +7 | +2 +3 +4 +5 +6 +7 +8 | +2 +3 +4 +6 | +2 +3 +4 | +1 +2 +3 | +2 |
Third Transition series
La | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg |
+3 | +3 +4 | +2 +3 +4 +5 | +2 +3 +4 +5 +6 | -1 +1 +2 +3 +4 +5 +6 +7 | +2 +3 +4 +6 +8 | +2 +3 +4 +6 | +2 +3 +4 +5 +6 | +1 +3 | +1 +2 |
*The oxidation states in italics are uncommon.
Important conclusion regarding oxidation states of transition elements
The examination of common oxidation states shown by different transition metals reveals the following facts :
(i) The variable oxidation states of transition metals are due to participation of inner (n -1) d and outer n s-electrons. The lowest oxidation state corresponds to the number of ns-electrons. For example, in the first transition series , the lowest oxidation states of Cr (3d54s1) and Cu(3d104s1) are +1 while for others , it is +2 (3d104s2)
(ii) Except scandium, the most common oxidation state of the first row transition elements is +2 which arises due to loss of two 4s-electrons. This means that after scandium 3d-orbitals become more stable and therefore , are lower in energy than the 4s-orbitals. As a result, electrons first removed from 4s-orbitals.
(iii) For the first five elements, the minimum oxidation state is equal to the number of electrons in the s-orbitals and the other oxidation states are given by the sum of outer s - and some or all d-electrons. The highest oxidation state is equal to the sum of the outer s (ns ) and (n -1) d-electrons. For the remaining five elements , the minimum oxidation state is given by the electrons in s-orbital while the maximum oxidation state is not related to their electronic configurations. The highest oxidation state shown by any transition metal is +8.
(iv) In the +2 and +3 oxidation states, the bonds formed are mostly ionic. In the compounds of higher oxidation states (generally formed wth oxygen and fluorine), the bonds are essentially covalent. Thus the bonds in +2 and +3 oxidation states are generally formed by the loss of two or three electrons respectively, while the bonds in higher oxidation states are formed by sharing of d-electrons. For example MnO4- (Mn in +7 ) state all the bonds are covalent.
(v) Within a group, the maximum oxidation state increases with atomic number. For example, iron (Group 8) shows common oxidation states of +2 and +3 but ruthenium and osmium in the same group form compounds in the +4 , +6 and +8 oxidation states.
(vi) Transition metals also form compounds in low oxidation states such as +1 and 0 or negative. The common examples are [Ni(CO)4], [Fe(CO)5] in which nickel and iron are in zero oxidation state.
Formation of coloured ions
Most of the compounds of the transition metals are coloured in the solid form or solution form. This is in contrast to the compounds of s- and p-block elements which are usually white.
The colour of the compounds of transition metals may be attributed to the presence of (n-1) d-subshell. In the case of compounds of transition metals, energies of five d-orbitals in the same sub-shell do not remain equal. Under the influence of approaching ions towards the central metal ion, the d-orbitals of the central metal split into different energy levels. This phenomenon is called crystal field splitting. For example, when six ions or molecules approach the metal ion(called octahedral field) , the five d-orbitals split into two sets : one set consisting of two d-orbitals(dx2- y2, dz2) of higher energy and the other set consisting of three d-orbitals (dxy, dyz, dxz) of lower energy.
In the case of transition metal ions, the electrons can be easily promoted from one energy level to another in the same d-subshell. These are called d-d transitions. The amount of energy required to excite some of the electrons to higher energy states with in the same d-subshell corresponds to energy of certain colours of visible light. Therefore , when white light falls on transition metal compound, some of its energy corresponding to a certain colour is absorbed and the electron gets raised from lower energy set of orbitals to higher energy set of orbitals as shown below:
The other colours constituting white light are transmitted and the compound appears coloured. The observed colour of the substance is always complimentary colour of the colour which is absorbed by the substance. For example, Ti3+ compounds contain one electron in d-subshell (d1). It absorbs green and yellow portions from the visible light and blue and red portions are emitted. Therefore Ti3+ ions appear purple. Similarly, hydrated cupric compounds absorb radiations corresponding to red light and the transmitted colour is greenish blue(which is complimentary colour to red colour). Thus cupric compounds have greenish-blue colour.
The frequency of the light absorbed is determined by the nature of the ligand. In aqueous solutions where water molecules are the ligands, the colour of the ions obseved are listed in TABLE.
The colours of some of the first row transition metal ions (aquated)
Configuration | Example | Colour |
3d0 | Sc3+ | Colouless |
3d0 | Ti4+ | Colourless |
3d1 | Ti3+ | Purple |
3d1 | V4+ | Blue |
3d2 | V3+ | Green |
3d3 | V2+ | Violet |
3d3 | Cr3+ | Violet |
3d4 | Mn3+ | Violet |
3d4 | Cr2+ | Blue |
3d5 | Mn2+ | Pink |
3d5 | Fe3+ | Yellow |
3d6 | Fe2+ | Green |
3d7 | Co2+ | Pink |
3d8 | Ni2+ | Green |
3d9 | Cu2+ | Blue |
3d10 | Zn2+ | Colourless |
It may be noted that the transition metal ions containing completely filled d-orbitals (d10) such as Zn2+, Cd2+, Hg2+, Cu+ etc. are generally white. For example ZnCl2 is white. Similarly Sc3+ and Ti4+ are white because they have completely empty d-orbitals.
Chemical Reactivity and E° values
Transition metals vary widely in their chemical reactivity. Many of them are suffiently electropositive to dissolve in mineral acids, although a few are ‘noble’ , that is they are unaffected by simple acids.
The metals of the first series with the exception of copper are relatively more reactive and are oxidised by 1 M H+ , though the actual rate at which these metals react with oxidising agents like hydrogen ion (H+) is some times slow. For example, titanium and vanadium , in practice , are passive to dilute non-oxidising acids at room temperature. The E° values for M2+/M (Table 1) indicate a decreasing tendency to form divalent cations across the series. The general trend towards less negative E° values is related to the general increase in sum of first and second ionisation enthalpies. The E° values of for Mn, Ni and Zn are more negative than expected from the general trend. Whereas the stabilities of half-filled d-subshell (d5 )in Mn2+ and completely filled d-subshell (d10) in zinc are related to their E° values ; for nickel, E°value is related to the highest negative enthalpy of hydration.
An examination of the E° values for the redox couple M3+/ M2+ (Table 1) shows that Mn3+and Co3+ ions are the strongest oxidising agents in aqueous solutions. The ions Ti2+, V2+ and Cr2+ are strong reducing agents and will liberate hydrogen from a dilute acid, e.g.
2 Cr2+(aq) + 2 H+(aq) ® 2 Cr3+(aq) + H2(g)
Magnetic Properties
The magnetic properties of a compound is a measure of the number of unpaired electrons in it. There are two main types of substances :
i) Paramagnetic substances : The substances which are attracted by magnetic field are called paramagnetic substances and this character arises due to the presence of unpaired electrons in the atomic orbitals.
ii) Diamagnetic substances : The substances which are repelled by magnetic field are called diamagnetic substances and this character arises due to the presence of paired electrons in the atomic orbitals.
Most of the compounds of transition elements are paramagnetic in nature and are attracted by the magnetic field.
The transition elements involve the partial filling of d-subshells. Most of the transition metal ions or their compounds have unpaired electrons in d-subshell (configuration from d1 to d9) and therefore, they give rise to paramagnetic character. The magnetic character is expressed in Bohr magnetons abbreviated as B.M. The magnetic moments of some ions of the first transition series are given in the following TABLE.
Magnetic moments of ions of first transition series.
Ion | Outer configuration | Number of unpaired electrons | Magnetic moment (mB) |
Sc3+ | 3d0 | 0 | 0 |
Ti 3+ | 3d1 | 1 | 1.75 |
V3+ | 3d2 | 2 | 2.76 |
Cr3+ | 3d3 | 3 | 3.86 |
Cr2+ | 3d4 | 4 | 4.8 |
Mn2+ | 3d5 | 5 | 5.96 |
Fe2+ | 3d6 | 4 | 5.10 |
Co2+ | 3d7 | 3 | 4.4-5.2 |
Ni2+ | 3d8 | 2 | 2.9 - 3.4 |
Cu2+ | 3d9 | 1 | 1.8 - 2.2 |
Zn2+ | 3d10 | 0 | 0 |
The magnetic moments arise only from the spin of electrons. This can be calculated from the relation :
where n is the number of unpaired electrons and B.M represents Bohr magneton. It is clear from the Table that as the number of unpaired electron increases from 1 to 5, the magnetic moment and hence the paramagnetic character also increases. After d5 configuration, there is decrease in magnetic moment due to decrease in number of unpaired electrons. For example, d6 configuration has 4 unpaired electrons, d7 configuration has 3 unpaired electrons and so on.
In addition to paramagnetic and diamagnetic substances, there are a few substances such as iron metal, iron oxide which are highly magnetic (about 1000 times more than ordinary metals). These are called Ferromagnetic substances.
Problem
03. Evaluate the magnetic moment of a divalent ion in aqueous solution if its atomic number is 25.