The d- and f - Block Elements

 The d- and f - Block Elements


The d-block of the periodic table contains the elements of groups 3-12 in which the d orbitals are progressively filled in each of the four long periods. The f-block consists of elements in which 4 f and 5 f orbitals are progressively filled. They are placed in a separate panel at the bottom of the periodic table. The names transition metals and inner transition metals are often used to refer to the elements of d-and f-blocks respectively. 

There are mainly four series of the transition metals

➤ 3d series (Sc to Zn)

 4d series (Y to Cd)

 5d series (La and Hf to Hg) 

 6d series which has Ac and elements from Rf to Cn. 

The two series of the inner transition metals; 4f (Ce to Lu) and 5f (Th to Lr) are known as lanthanoids and actinoids respectively. 

πŸ‘‰ Originally the name transition metals was derived from the fact that their chemical properties were transitional between those of s and p-block elements. 

πŸ‘‰ According to IUPAC, transition metals are defined as metals which have incomplete d subshell either in neutral atom or in their ions.

πŸ‘‰ Zinc, cadmium and mercury of group 12 have full `d^10` configuration in their ground state as well as in their common oxidation states and hence, are not regarded as transition metals.

Various precious metals such as silver, gold and platinum and industrially important metals like iron, copper and titanium belong to the transition metals series.

Position in the Periodic Table

The d–block occupies the large middle section of the periodic table flanked between s– and p– blocks in the periodic table. The d–orbitals of the penultimate energy level of atoms receive electrons giving rise to four rows of the transition metals, i.e., 3d, 4d, 5d and 6d.

 Electronic Configurations of the d-Block Elements Elements

πŸ‘‰ In general the electronic configuration of outer orbitals of these elements is `(n-1)d^ {1–10}` `ns^{ 1–2}`. The `(n–1)` stands for the inner d orbitals which may have one to ten electrons and the outermost ns orbital may have one or two electrons. 

This generalisation has several exceptions because of very little energy difference between `(n-1)d` and `ns` orbitals. Furthermore, half and completely filled sets of orbitals are relatively more stable. A consequence of this factor is reflected in the electronic configurations of Cr and Cu in the 3d series.

For example, consider the case of Cr, which has `3d^5 4s^1`  configuration instead of `3d^4 4s^2` ; the energy gap between the two sets (3d and 4s) of orbitals is small enough to prevent electron entering the 3d orbitals. Similarly in case of Cu, the configuration is `3d^10 4s^1`  and not `3d ^9 4s^2` .

πŸ‘‰ The electronic configurations of outer orbitals of Zn, Cd, Hg and Cn are represented by the general formula `(n-1)d^10 ns^2` . The orbitals in these elements are completely filled in the ground state as well as in their common oxidation states. Therefore, they are not regarded as transition elements.

πŸ‘‰  With partly filled d orbitals these elements exhibit certain characteristic properties such as display of a variety of oxidation states, formation of coloured ions and entering into complex formation with a variety of ligands.

πŸ‘‰ The transition metals and their compounds also exhibit catalytic property and paramagnetic behaviour.

πŸ‘‰ There are greater similarities in the properties of the transition elements of a horizontal row in contrast to the non-transition elements. However, some group similarities also exist.

Melting and Boiling Points

Transition metals have very high melting and boiling points. It is clear from the figure that the melting points of these metals rise to a maximum value and then decrease with an increase in atomic number except for anomalous values of Mn and Tc. 

The high melting and boiling points of these metals are due to strong metallic bonds between the atoms of these elements. This is also evident from the fact that these metals have high enthalpies of atomization. The metallic bond is formed due to the interaction of electrons in the outermost orbitals. The strength of bonding is roughly related to the number of unpaired electrons. In general, the greater the number of valence electrons, the stronger the metallic bonding, and consequently, melting points are high.

Therefore, as we move along a particular series, the metallic strength increases upto the middle with increasing availability of unpaired electrons upto `d^5` configuration (e.g. Sc has 1, Ti has 2, V has 3, Cr has 5 unpaired electrons) and then decreases with decreasing availability of unpaired d-electrons (e.g. Fe has 4, Co has 3 unpaired electrons and so on). Therefore, the melting points decrease after the middle because of increase of pairing of electrons. The elements of group 12 (zinc, cadmium and mercury) are quite soft with low melting points. Mercury is a liquid at room temperature and melts at -38°C. These three elements behave typically because there are no unpaired electrons available for metallic bonding and, therefore, their melting points are low.

Why do the transition elements exhibit higher enthalpies of atomisation?

 Because of large number of unpaired electrons in their atoms they have stronger interatomic interaction and hence stronger bonding between atoms resulting in higher enthalpies of atomisation. 

Variation in Atomic and Ionic Sizes of Transition Metals

An interesting point emerges when atomic sizes of one series are compared with those of the corresponding elements in the other series. The curves in Fig. show an increase from the first `(3d)` to the second `(4d)` series of the elements but the radii of the third `(5d)` series are 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 compensates for the expected increase in atomic size with increasing atomic number. The net result of the lanthanoid contraction is that the second and the third `d` series exhibit similar radii (e.g., Zr 160 pm, Hf 159 pm).
However, the shielding of one `4f` electron by another is less than that of one d electron by another, and as the nuclear charge increases along the series, there is fairly regular decrease in the size of the entire `4f^n` orbitals.


πŸ‘‰ The decrease in metallic radius coupled with increase in atomic mass results in a general increase in the density of these elements. Thus, from titanium (Z = 22) to copper (Z = 29) the significant increase in the density.

Oxidation States

One of the notable features of a transition elements is the great variety of oxidation states.

πŸ‘‰ The elements which give the greatest number of oxidation states occur in or near the middle of the series. Manganese, for example, exhibits all the oxidation states from `+2` to `+7`.

πŸ‘‰ The lesser number of oxidation states at the extreme ends stems from either too few electrons to lose or share (Sc, Ti) or too many d electrons (hence fewer orbitals available in which to share electrons with others) for higher valence (Cu, Zn). 

πŸ‘‰ Early in the series scandium(II) is virtually unknown and titanium (IV) is more stable than Ti(III) or Ti(II). At the other end, the only oxidation state of zinc is `+2` (no d electrons are involved).

πŸ‘‰ The maximum oxidation states of reasonable stability correspond in value to the sum of the s and d electrons upto manganese `(Ti^{iv}, V^v, Cr^{vi}, Mn^{vii})` followed by a rather abrupt decrease in stability of higher oxidation states, so that the typical species to follow are `(Fe^{ii,iii}, Co^{ii,iii}, Ni^{ii}, Cu^{i,ii}, Zn^{ii})`.

πŸ‘‰ The variability of oxidation states, a characteristic of transition elements, arises out of incomplete filling of d orbitals in such a way that their oxidation states differ from each other by unity, e.g., `V^{ii}, V^{iii}, V^{iv}, V^{v}`. This is in contrast with the variability of oxidation states of non transition elements where oxidation states normally differ by a unit of two.

Name a transition element which does not exhibit variable oxidation states. 

Scandium (Z = 21) does not exhibit variable oxidation states.

Formation of Coloured Ions

πŸ‘‰ When an electron from a lower energy d orbital is excited to a higher energy d orbital, the energy of excitation corresponds to the frequency of light absorbed . This frequency generally lies in the visible region. The colour observed corresponds to the complementary colour of the light absorbed.

πŸ‘‰ The frequency of the light absorbed is determined by the nature of the ligand.

πŸ‘‰ In aqueous solutions where water molecules are the ligands, the colours of the ions observed are listed in Table

Magnetic Properties

When a magnetic field is applied to substances, mainly two types of magnetic behaviour are observed: diamagnetism and paramagnetism

πŸ‘‰Diamagnetic substances are repelled by the applied field while the paramagnetic substances are attracted.

πŸ‘‰ Substances which are attracted very strongly are said to be ferromagnetic. In fact, ferromagnetism is an extreme form of paramagnetism.

πŸ‘‰ Paramagnetism arises from the presence of unpaired electrons, each such electron having a magnetic moment associated with its spin angular momentum and orbital angular momentum.
For the compounds of the first series of transition metals, the contribution of the orbital angular momentum is effectively quenched and hence is of no significance. For these, the magnetic moment is determined by the number of unpaired electrons and is calculated by using the ‘spin-only’ formula, i.e.,

`\mu = sqrt{ n(n + 2)}`

where `n` is the number of unpaired electrons and `Β΅` is the magnetic moment in units of Bohr magneton (BM). A single unpaired electron has a magnetic moment of 1.73 Bohr magnetons (BM).

The magnetic moment increases with the increasing number of unpaired electrons. Thus, the observed magnetic moment gives a useful indication about the number of unpaired electrons present in the atom, molecule or ion. The magnetic moments calculated from the ‘spin-only’ formula and those derived experimentally for some ions of the first row transition elements are given in Table

 Formation of Complex Compounds

πŸ‘‰Complex compounds are those in which the metal ions bind a number of anions or neutral molecules giving complex species with characteristic properties.

A few examples are: `[Fe (CN)_6 ]^{3-}`,  `[Fe (CN)_6 ]^{4-}`, `[Cu(NH_3)_4]^{2+}`, `[PtCl_4]^{2-}`

πŸ‘‰ The transition metals form a large number of complex compounds. This is due to  their high ionic charges and the availability of d orbitals for bond formation

Catalytic Properties

πŸ‘‰πŸ‘‰    catalyst is a substance that speeds up a chemical reaction without being consumed in the process. It lowers the activation energy required for the reaction to occur, making it happen more quickly.

πŸ‘‰ The transition metals and their compounds are known for their catalytic activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes. 


1. Vanadium(V) oxide (in Contact Process) πŸ‘‰ formation of sulphuric acid , 

2. Finely divided iron (in Haber’s Process) πŸ‘‰  ammonia

3.  Nickel (in Catalytic Hydrogenation) πŸ‘‰ unsaturated fat to saturated fat. 

πŸ‘‰ Catalysts at a solid surface involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst (first row transition metals utilise 3d and 4s electrons for bonding). This has the effect of increasing the concentration of the reactants at the catalyst surface and also weakening of the bonds in the reacting molecules (the activation energy is lowering).

 πŸ‘‰ The transition metal ions can change their oxidation states, they become more effective as catalysts.


`2I^{-} +  S_2O_8^{2-} \rightarrow I_2 + 2SO_4^{2-}`

For example, iron(III) catalyses the reaction between iodide and persulphate ions

`2Fe^{3+} + 2I^{-} \rightarrow  2Fe^{2+} + I_2`

`2Fe^{2+} + S_2O_8^{2-}  \rightarrow  2Fe^{3+} +   2SO_4^{2-}`

Formation of Interstitial Compounds

πŸ‘‰πŸ‘‰Interstitial compounds are those which are formed when small atoms like H, C or N are trapped inside the crystal lattices of metals.

πŸ‘‰ They are usually non-stoichiometric and are neither typically ionic nor covalent, for example, `TiC, Mn_4N, Fe_3H, VH_0.56, TiH_1.7`, etc.

πŸ‘‰ The formulas quoted do not, of course, correspond to any normal oxidation state of the metal.

 Because of the nature of their composition, these compounds are referred to as interstitial compounds. The principal physical and chemical characteristics of these compounds are as follows: 
(i) They have high melting points, higher than those of pure metals. 
(ii) They are very hard, some borides approach diamonds in hardness. 
(iii) They retain metallic conductivity. 
(iv) They are chemically inert. 

Alloy Formation

πŸ‘‰πŸ‘‰ An alloy is a blend of different metals, or a metal mixed with other elements, resulting in a new material with improved properties.
Alloys may be homogeneous solid solutions in which the atoms of one metal are distributed randomly among the atoms of the other.

πŸ‘‰ Alloys are formed by atoms with metallic radii that are within about 15 per cent of each other.

πŸ‘‰ Because of similar radii and other characteristics of transition metals the atoms of one metal can substitute the atoms of other atoms. 

πŸ‘‰ The alloys so formed are hard and have often high melting points


The best known are ferrous alloys: chromium, vanadium, tungsten, molybdenum and manganese are used for the production of a variety of steels and stainless steel. Alloys of transition metals with non-transition metals such as brass (copper-zinc) and bronze (copper-tin), are also of considerable industrial importance

Oxides of Metals

πŸ‘‰These oxides are generally formed by the reaction of metals with oxygen at high temperatures.

πŸ‘‰ All the metals except scandium form `MO` oxides which are ionic.

πŸ‘‰ The highest oxidation number in the oxides, coincides with the group number and is attained in `Sc_2O_3`  to `Mn_2O_7` .

πŸ‘‰ Beyond group 7, no higher oxides of iron above `Fe_2O_3`  are known. 

 πŸ‘‰ Besides the oxides `(V_2O_5, V_2O_4, TiO_2)`, the oxocations stabilise `V^v` as `VO_2^+`, `V^{iv}` as `VO^{2+}`, `Ti^{iv}` as `TiO^{2+}`.

 πŸ‘‰πŸ‘‰ Oxoanions are negatively charged ions composed of oxygen and other elements, forming essential compounds in chemistry with distinct structures and properties.

πŸ‘‰πŸ‘‰ Oxocations are positively charged ions (cations) formed by combining oxygen with other elements, influencing chemical reactivity and bonding in compounds.

πŸ‘‰ As the oxidation number of a metal increases, its ionic character decreases. In the case of `Mn, Mn_2O_7`  is a covalent green oil. . Even `CrO_3` and `V_2O_`5`  have low melting points.

O.N ↑  I.C ↓ Cov.C ↑

O.N ↑ M.P ↓

πŸ‘‰ In these higher oxides, the acidic character is predominant.   O.N ↑ Ac ↑

Thus, `Mn_2O_7` gives `HMnO_4` and `CrO_3` gives `H_2CrO_4` and `H_2Cr_2O_7`. 

`Mn_2O_7 + H_2O \rightarrow 2HMnO_4`

`CrO_3 + H_2O \rightarrow H_2CrO_4` 

πŸ‘‰ `V_2O_5` is, however, amphoteric though mainly acidic and it gives `VO_4^{3-}` as well as `VO_2^{+}`  salts. 

πŸ‘‰ In vanadium there is gradual change from the basic `V_2O_3`  to less basic `V_2O_4`  and to amphoteric `V_2O_5`, `V_2O_4`  dissolves in acids to give `VO^{2+}` salts. 

πŸ‘‰ `V_2O_5` reacts with alkalies as well as acids to give `VO_4^{3-}` and `VO_4^{+}` respectively.

πŸ‘‰ The well characterised `CrO` is basic but `Cr_2O_3`  is amphoteric.

Potassium dichromate `K_2Cr_2O_7`

Potassium dichromate is a very important chemical used in leather industry and as an oxidant for preparation of many azo compounds. 

πŸ‘‰ Dichromates are generally prepared from chromate, which in turn are obtained by the fusion of chromite ore (`FeCr_2O_4` ) with sodium or potassium carbonate in free access of air. The reaction with sodium carbonate occurs as follows:

`4 FeCr_2O_4 + 8 Na_2CO_3 + 7 O_2 → 8 Na_2CrO_4 + 2 Fe_2O_3 + 8 CO_2`

πŸ‘‰ The yellow solution of sodium chromate is filtered and acidified with sulphuric acid to give a solution from which orange sodium dichromate, `Na_2Cr_2O_7 . 2H_2O` can be crystallised.

`2Na_2CrO_4 + 2 H^+ → Na_2Cr_2O_7 + 2 Na^+ + H_2O`

πŸ‘‰ Sodium dichromate is more soluble than potassium dichromate. The latter is therefore, prepared by treating the solution of sodium dichromate with potassium chloride

`Na_2Cr_2O_7 + 2 KCl → K_2Cr_2O_7 + 2 NaCl`

Orange crystals of potassium dichromate crystallise out. 

πŸ‘‰ The chromates and dichromates are interconvertible in aqueous solution depending upon pH of the solution. The oxidation state of chromium in chromate and dichromate is the same.

`2 CrO_4^{2–} + 2H^+ → Cr_2O_7^{2–} + H_2O` (pH↓)

`Cr_2O_7^{ 2–} + 2 OH^{ -} → 2 CrO_4 ^{2–} + H_2O` (pH ↑)

The structures of chromate ion, `CrO_4^{2–}` and the dichromate ion, `Cr_2O_7^{2–}`

πŸ‘‰The chromate ion is tetrahedral whereas the dichromate ion consists of two tetrahedra sharing one corner with Cr–O–Cr bond angle of `126°`.

πŸ‘‰ Sodium and potassium dichromates are strong oxidising agents; the sodium salt has a greater solubility in water and is extensively used as an oxidising agent in organic chemistry. Potassium dichromate is used as a primary standard in volumetric analysis. In acidic solution, its oxidising action can be represented as follows:

`Cr_2O_7^{2–} + 14H^+ + 6e^– → 2Cr_3^+ + 7H_2O` `(E^⊝ = 1.33V)`

πŸ‘‰ Acidified potassium dichromate will oxidise iodides to iodine, sulphides to sulphur, tin(II) to tin(IV) and iron(II) salts to iron(III). The half-reactions are noted below:

πŸ‘‰ The full ionic equation may be obtained by adding the half-reaction for potassium dichromate to the half-reaction for the reducing agent, for e.g.,

`Cr_2O_7^{ 2–} + 14 H^+ + 6 Fe^{2+} → 2 Cr^{3+} + 6 Fe^{3+} + 7 H_2O`

Potassium permanganate` KMnO_4`

πŸ‘‰ Potassium permanganate is prepared by fusion of `MnO_2` with an alkali metal hydroxide and an oxidising agent like `KNO_3`. This produces the dark green `K_2MnO_4` which disproportionates in a neutral or acidic solution to give permanganate.

`2MnO_2 + 4KOH + O_2 → 2K_2MnO_4 + 2H_2O`

`3MnO_4^{ 2–} + 4H^+ → 2MnO_4^– + MnO_2 + 2H_2O`

πŸ‘‰ Commercially it is prepared by the alkaline oxidative fusion of `MnO_2` followed by the electrolytic oxidation of manganate (Vl)

πŸ‘‰ In the laboratory, a manganese (II) ion salt is oxidised by peroxodisulphate to permanganate.

`2Mn^{2+} + 5S_2O_8^{ 2–} + 8H_2O → 2MnO_4^ – + 10SO_4^ {2–} + 16H^+`

πŸ‘‰ Potassium permanganate forms dark purple (almost black) crystals which are isostructural with those of  `KClO_4` . The salt is not very soluble in water (6.4 g/100 g of water at 293 K), but when heated it decomposes at 513 K. 

`2KMnO_4 → K_2MnO_4 + MnO_2 + O_2`

πŸ‘‰ It has two physical properties of considerable interest: its intense colour and its diamagnetism along with temperature-dependent weak paramagnetism.

πŸ‘‰ The manganate and permanganate ions are tetrahedral; the `Ο€-` bonding takes place by overlap of `p` orbitals of oxygen with `d` orbitals of manganese. The green manganate is paramagnetic because of one unpaired electron but the permanganate is diamagnetic due to the absence of unpaired electron.

πŸ‘‰ Acidified permanganate solution oxidises oxalates to carbon dioxide, iron(II) to iron(III), nitrites to nitrates and iodides to free iodine. The half-reactions of reductants are:

πŸ‘‰ If we represent the reduction of permanganate to manganate, manganese dioxide and manganese(II) salt by half-reactions,

We can very well see that the hydrogen ion concentration of the solution plays an important part in influencing the reaction.

πŸ‘‰ A few important oxidising reactions of `KMnO_4`  are given below: 

1. In acid solutions:

2. In neutral or faintly alkaline solutions:


πŸ‘‰The f-block consists of the two series, lanthanoids (the fourteen elements following lanthanum) and actinoids (the fourteen elements following actinium). 

πŸ‘‰ Because lanthanum closely resembles the lanthanoids, it is usually included in any discussion of the lanthanoids for which the general symbol Ln is often used. Similarly, a discussion of the actinoids includes actinium besides the fourteen elements constituting the series.

πŸ‘‰ The lanthanoids resemble one another more closely than do the members of ordinary transition elements in any series. They have only one stable oxidation state and their chemistry provides an excellent opportunity to examine the effect of small changes in size and nuclear charge along a series of otherwise similar elements.

πŸ‘‰ The chemistry of the actinoids is, on the other hand, much more complicated. The complication arises partly owing to the occurrence of a wide range of oxidation states in these elements and partly because their radioactivity creates special problems in their study.

πŸ‘‰ The Lanthanoids

The names, symbols, electronic configurations of atomic and some ionic states and atomic and ionic radii of lanthanum and lanthanoids (for which the general symbol Ln is used) are given in Table.

 πŸ‘‰ Electronic Configurations   `(n-2)f^{1-14}  (n-1)d^{0-1}  ns^2`

It may be noted that atoms of these elements have electronic configuration with `6s^2` common but with variable occupancy of `4f` level. However, the electronic configurations of all the tripositive ions (the most stable oxidation state of all the lanthanoids) are of the form `4f^n`  (n = 1 to 14 with increasing atomic number).

πŸ‘‰ Atomic and Ionic Sizes

The overall decrease in atomic and ionic radii from lanthanum to lutetium (the lanthanoid contraction) is a unique feature in the chemistry of the lanthanoids. It has far reaching consequences in the chemistry of the third transition series of the elements. 

The decrease in atomic radii (derived from the structures of metals) is not quite regular as it is regular in `M ^{3+}` ions (Fig.). 

This contraction is, of course, similar to that observed in an ordinary transition series and is attributed to the same cause, the imperfect shielding of one electron by another in the same sub-shell. 

 Oxidation States

πŸ‘‰All lanthanoids exhibit a common stable oxidation state of `+3`.  However, occasionally `+2` and `+4` ions in solution or in solid compounds are also obtained.  This irregularity (as in ionisation enthalpies) arises mainly from the extra stability of empty, half-filled or filled  f subshell. 

 πŸ‘‰The formation of `Ce^{IV}` is favoured by its noble gas configuration, but it is a strong oxidant reverting to the common `+3` state. 
The `E^o`  value for `Ce^{4+}`/ `Ce^{3+}` is `+ 1.74 V` which suggests that it can oxidise water.

`Ce^{4+} + e^{-} \rightarrow  Ce^{3+}`

However, the reaction rate is very slow and hence `Ce(IV)` is a good analytical reagent.

πŸ‘‰ Pr, Nd, Tb and Dy also exhibit `+4` state but only in oxides, `MO_2` .

πŸ‘‰ `Eu^{2+}` is formed by losing the two `s` electrons and its `f^7`  configuration accounts for the formation of this ion. However, `Eu^{2+}` is a strong reducing agent changing to the common `+3` state. 

`Eu^{2+}  \rightarrow  Eu^{3+} + e^{-}`

πŸ‘‰ Similarly `Yb^{2+}` which has `f^{14}` configuration is a reductant.

`Yb^{2+}  \rightarrow  Yb^{3+} + e^{-}`

πŸ‘‰`Tb^{IV}` has half-filled `f`-orbitals and is an oxidant. 

`Tb^{4+} + e^{-} \rightarrow  Tb^{3+}`

πŸ‘‰The behaviour of samarium is very much like europium, exhibiting both `+2` and `+3` oxidation states.

⇒ Oxidising agents   `Ce^{4+}, Tb^{4+}`

`Ce^{4+} + e^{-} \rightarrow  Ce^{3+}`

`Tb^{4+} + e^{-} \rightarrow  Tb^{3+}`

⇒ Reducing agents   `Eu^{2+}, Yb^{2+}`

`Eu^{2+}  \rightarrow  Eu^{3+} + e^{-}`

`Yb^{2+}  \rightarrow  Yb^{3+} + e^{-}`

General Characteristics

πŸ‘‰All the lanthanoids are silvery white soft metals and tarnish rapidly in air. The hardness increases with increasing atomic number, samarium being steel hard. Their melting points range between 1000 to 1200 K but samarium melts at 1623 K. They have typical metallic structure and are good conductors of heat and electricity. Density and other properties change smoothly except for Eu and Yb and occasionally for Sm and Tm.

πŸ‘‰Many trivalent lanthanoid ions are coloured both in the solid state and in aqueous solutions. Colour of these ions may be attributed to the presence of `f` electrons. Neither `La^{3+}` nor `Lu^{3+}` ion shows any colour but the rest do so. However, absorption bands are narrow, probably because of the excitation within `f` level. 

πŸ‘‰The lanthanoid ions other than the `f^0`  type `(La^{3+} &  Ce^{4+})` and the `f^14` type `(Yb^{2+} &  Lu^{3+})` are all paramagnetic.

πŸ‘‰The first ionisation enthalpies of the lanthanoids are around `600 kJ mol^{–1}`, the second about `1200 kJ mol^{–1}` comparable with those of calcium. 
The variation of third ionisation enthalpies show some stabilities of  empty `(f^0)`, half-filled `(f^7)` and completely filled `(f^14)`orbitals `f ` level. This is indicated from the abnormally low value of the third ionisation enthalpy of lanthanum `(4f^0)`, gadolinium `(4f^7)` and lutetium `(4f^14)`.

πŸ‘‰In their chemical behaviour, in general, the earlier members of the series are quite reactive similar to calcium but, with increasing atomic number, they behave more like aluminium. Values for E V for the half-reaction:

`Ln^{3+}(aq) + 3e^– → Ln(s)`

are in the range of –2.2 to –2.4 V except for Eu for which the value is – 2.0 V. This is, of course, a small variation.

Chemical Reactions

πŸ‘‰The metals combine with hydrogen when gently heated in the gas.

πŸ‘‰The carbides, `Ln_3C, Ln_2C_3`  and `LnC_2` are formed when the metals are heated with carbon.

πŸ‘‰ They liberate hydrogen from dilute acids and burn in halogens to form halides.

πŸ‘‰They form oxides `M_2O_3` and hydroxides `M(OH)_3` . The hydroxides are definite compounds, not just hydrated oxides. They are basic like alkaline earth metal oxides and hydroxides. 

Their general reactions are depicted in Fig.


The Actinoids

πŸ‘‰The actinoids include the fourteen elements from Th to Lr.

πŸ‘‰The actinoids are radioactive elements and the earlier members have relatively long half-lives, the latter ones have half-life values ranging from a day to 3 minutes for lawrencium (Z =103). The latter members could be prepared only in nanogram quantities. These facts render their study more difficult.

Electronic Configurations 

πŸ‘‰All the actinoids are believed to have the electronic configuration of `7s^2` and variable occupancy of the `5f` and `6d` subshells.

πŸ‘‰The fourteen electrons are formally added to `5f`, though not in thorium (Z = 90) but from Pa onwards the `5f` orbitals are complete at element 103.

πŸ‘‰The irregularities in the electronic configurations of the actinoids, like those in the lanthanoids are related to the stabilities of the `f ^0 , f ^7`  and `f^14` occupancies of the `5f` orbitals. Thus, the configurations of Am and Cm are `[Rn] 5f ^7 7s^ 2` and `[Rn] 5f ^{7} 6d^1 7s^ 2`.

πŸ‘‰Although the `5f` orbitals resemble the `4f` orbitals in their angular part of the wave-function, they are not as buried as `4f` orbitals and hence `5f `electrons can participate in bonding to a far greater extent.

Ionic Sizes

πŸ‘‰The general trend in lanthanoids is observable in the actinoids as well. There is a gradual decrease in the size of atoms or `M^{3+}` ions across the series. This may be referred to as the actinoid contraction (like lanthanoid contraction). The contraction is, however, greater from element to element in this series resulting from poor shielding by 5f electrons.

Oxidation States

There is a greater range of oxidation states, which is in part attributed to the fact that the `5f, 6d` and `7s` levels are of comparable energies. 

πŸ‘‰The actinoids show in general `+3` oxidation state. The elements, in the first half of the series frequently exhibit higher oxidation states. For example, the maximum oxidation state increases from `+4` in Th to `+5, +6` and `+7` respectively in Pa, U and Np but decreases in succeeding elements.

πŸ‘‰The actinoids resemble the lanthanoids in having more compounds in `+3` state than in the `+4` state. However, `+3` and `+4` ions tend to hydrolyse. 

πŸ‘‰Because the distribution of oxidation states among the actinoids is so uneven and so different for the former and later elements, it is unsatisfactory to review their chemistry in terms of oxidation states.

General Characteristics and Comparison with Lanthanoids

πŸ‘‰The actinoid metals are all silvery in appearance but display a variety of structures. The structural variability is obtained due to irregularities in metallic radii which are far greater than in lanthanoids.

πŸ‘‰The actinoids are highly reactive metals, especially when finely divided.

πŸ‘‰The action of boiling water on them, for example, gives a mixture of oxide and hydride and combination with most non metals takes place at moderate temperatures.

πŸ‘‰Hydrochloric acid attacks all metals but most are slightly affected by nitric acid owing to the formation of protective oxide layers; alkalies have no action.

πŸ‘‰The magnetic properties of the actinoids are more complex than those of the lanthanoids. Although the variation in the magnetic susceptibility of the actinoids with the number of unpaired 5 f electrons is roughly parallel to the corresponding results for the lanthanoids, the latter have higher values.

πŸ‘‰It is evident from the behaviour of the actinoids that the ionisation enthalpies of the early actinoids, though not accurately known, but are lower than for the early lanthanoids. This is quite reasonable since it is to be expected that when 5f orbitals are beginning to be occupied, they will penetrate less into the inner core of electrons The 5f electrons, will therefore, be more effectively shielded from the nuclear charge than the 4f electrons of the corresponding lanthanoids. Because the outer electrons are less firmly held, they are available for bonding in the actinoids.

πŸ‘‰A comparison of the actinoids with the lanthanoids, with respect to different characteristics as discussed above, reveals that behaviour similar to that of the lanthanoids is not evident until the second half of the actinoid series. However, even the early actinoids resemble the lanthanoids in showing close similarities with each other and in gradual variation in properties which do not entail change in oxidation state. The lanthanoid and actinoid contractions, have extended effects on the sizes, and therefore, the properties of the elements succeeding them in their respective periods. The lanthanoid contraction is more important because the chemistry of elements succeeding the actinoids are much less known at the present time.


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