Chapter 8 The d-and f-Block Elements

The transition elements are those elements having a partially filled d or f subshell in any common oxidation state. The term "transition elements" most commonly refers to the d-block transition elements. Ag has a completely filled 4d orbital (4d10 5s1) in its ground state. Now, silver displays two oxidation states (+1 and +2). In the +1 oxidation state, an electron is removed from the s-orbital. However, in the +2 oxidation state, an electron is removed from the d-orbital. Thus, the d-orbital now becomes incomplete (4d9). Hence, it is a transition element.

Sc & zn belongs to 3^rd group pf periodic table. The extent of metallic bonding an element undergoes decides the enthalpy of atomization. The more extensive the metallic bonding of an element, the more will be its enthalpy of atomization. In all transition metals (except Zn, electronic configuration: 3d10 4s2), there are some unpaired electrons that account for their stronger metallic bonding. Due to the absence of these unpaired electrons, the inter-atomic electronic bonding is the weakest in Zn and as a result, it has the least enthalpy of atomization.

Mn (Z = 25) = 3d⁵ 4s²

 Mn⁺² is the most stable ion for manganese, the d-orbital can be made to remove 0 to 7 electrons. Compounds of manganese therefore range from Mn(0) as Mn_(s), Mn(II) as MnO, Mn(II,III) as Mn₃O₄, Mn(IV) as MnO₂, or manganese dioxide, Mn(VII) in the permanganate ion MnO₄^-, and so on. Mn has the maximum number of unpaired electrons present in the d-subshell (5 electrons). Hence, Mn exhibits the largest number of oxidation states, ranging from +2 to +7.

The E^θ(M²⁺/M) value of a metal depends on the energy changes involved in the following:

1. Sublimation: The energy required for converting one mole of an atom from the solid state to the gaseous state.

2. Ionization: The energy required to take out electrons from one mole of atoms in the gaseous state.

3. Hydration: The energy released when one mole of ions are hydrated.

Now, copper has a high energy of atomization and low hydration energy. Hence, the Eθ(M2+/M) value for copper is positive.

Ionization enthalpies are found to increase in the given series due to a continuous filling of the inner d-orbitals. The irregular variations of ionization enthalpies can be attributed to the extra stability of configurations such as d⁰, d^5, d¹⁰. Since these states are exceptionally stable, their ionization enthalpies are very high.

In case of first ionization energy, Cr has low ionization energy. This is because after losing one electron, it attains the stable configuration (3d⁵). On the other hand, Zn has exceptionally high first ionization energy as an electron has to be removed from stable and fully-filled orbitals (3d¹⁰ 4s²).

Second ionization energies are higher than the first since it becomes difficult to remove an electron when an electron has already been taken out. Also, elements like Cr and Cu have exceptionally high second ionization energies as after losing the first electron, they have attained the stable configuration (Cr⁺: 3d⁵ and Cu⁺: 3d¹⁰). Hence, taking out one electron more from this stable configuration will require a lot of energy.

The oxidation state of an element is related to the number of electrons that an atom loses, gains, or appears to use when joining with another atom in compounds. It also determines the ability of an atom to oxidize (to lose electrons) or to reduce (to gain electrons) other atoms or species. Oxidation results in an increase in the oxidation state. Reduction results in a decrease in the oxidation state.  If an atom is reduced, it has a higher number of valence shell electrons, and therefore a higher oxidation state, and is a strong oxidant. For example, oxygen (O) and fluorine (F) are very strong oxidants.Both oxide and fluoride ions are highly electronegative and have a very small size. Due to these properties, they are able to oxidize the metal to its highest oxidation state.

The following reactions are involved when Cr²⁺ and Fe²⁺ act as reducing agents.

Cr²⁺  Cr³⁺ Fe²⁺  Fe³⁺

The  value is - 0.41 V and  is +0.77 V. This means that Cr²⁺ can be easily oxidized to Cr³⁺, but Fe²⁺ does not get oxidized to Fe³⁺ easily. Therefore, Cr²⁺ is a better reducing agent that Fe³⁺.

Z = 27

[Ar] 3d⁷ 4s²

M²⁺ = [Ar] 3d⁷

3d⁷ =

i.e., 3 unpaired electrons

n = 3

In an aqueous medium, Cu²⁺ is more stable than Cu⁺. This is because although energy is required to remove one electron from Cu⁺ to Cu²⁺, high hydration energy of Cu²⁺ compensates for it. Therefore, Cu⁺ ion in an aqueous solution is unstable. It disproportionates to give Cu²⁺ and Cu.

In actinoids, 5f orbitals are filled. These 5f orbitals have a poorer shielding effect than 4f orbitals (in lanthanoids). Thus, the effective nuclear charge experienced by electrons in valence shells in case of actinoids is much more than that experienced by lanthanoids. Hence, the size contraction in actinoids is greater as compared to that in lanthanoids.

(i) Cr³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³

Or, [Ar]¹⁸3d³

(ii) Pm³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰5s² 5p⁶ 4f⁴

Or, [Xe]⁵⁴ 3d³

(iii) Cu⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰

Or, [Ar]¹⁸ 3d¹⁰

(iv) Ce⁴⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁶

Or, [Xe]⁵⁴

(v) Co²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷

Or, [Ar]¹⁸3d⁷

(vi) Lu²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁶ 4f¹⁴ 5d¹

Or, [Xe]⁵⁴2f¹⁴3d³

(vii) Mn²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵

Or, [Ar]¹⁸ 3d⁵

Electronic configuration of Mn²⁺ is [Ar]¹⁸ 3d⁵.

Electronic configuration of Fe²⁺ is [Ar]¹⁸ 3d⁶.

It is known that half-filled and fully-filled orbitals are more stable. Therefore, Mn in (+2) state has a stable d⁵ configuration. This is the reason Mn²⁺shows resistance to oxidation to Mn³⁺. Also, Fe²⁺has 3d⁶ configuration and by losing one electron, its configuration changes to a more stable 3d⁵ configuration. Therefore, Fe²⁺ easily gets oxidized to Fe⁺³ oxidation state.

The oxidation states displayed by the first half of the first row of transition metals are given in the table below. 

	Sc	Ti	V	Cr	Mn
		+2	+2	+2	+2
	+3	+3	+3	+3	+3
Oxidation state		+4	+4	+4	+4
			+5	+5	+6
				+6	+7

It can be easily observed that except Sc, all others metals display +2 oxidation state. Also, on moving from Sc to Mn, the atomic number increases from 21 to 25. This means the number of electrons in the 3d-orbital also increases from 1 to 5.

Sc (+2) = d¹

Ti (+2) = d²

V (+2) = d³

Cr (+2) = d⁴

Mn (+2) = d⁵

+2 oxidation state is attained by the loss of the two 4 selectrons by these metals. Since the number of d electrons in (+2) state also increases from Ti(+2) to Mn(+2), the stability of +2 state increases (as d-orbital is becoming more and more half-filled). Mn (+2) has d⁵ electrons (that is half-filled d shell, which is highly stable).

The elements in the first-half of the transition series exhibit many oxidation states with Mn exhibiting maximum number of oxidation states (+2 to +7). The stability of +2 oxidation state increases with the increase in atomic number. This happens as more electrons are getting filled in the d-orbital. However, Sc does not show +2 oxidation state. Its electronic configuration is 4s²3d¹. It loses all the three electrons to form Sc³⁺. +3 oxidation state of Sc is very stable as by losing all three electrons, it attains stable noble gas configuration, [Ar]. Ti (+ 4) and V (+5) are very stable for the same reason. For Mn, +2 oxidation state is very stable as after losing two electrons, its d-orbital is exactly half-filled, [Ar] 3d⁵.

	Electronic configuration in ground state	Stable oxidation states
(i)	3d3 (Vanadium)	+2, +3, +4 and +5
(ii)	3d5 (Chromium)	+3, +4, +6
(iii)	3d5 (Manganese)	+2, +4, +6, +7
(iv)	3d8 (Cobalt)	+2, +3
(v)	3d4	There is no3d4 configuration in ground state.

 

(i) Vanadate, VO^-₃

Oxidation state of V is + 5.

 

(ii) Chromate, CrO^2-₄

Oxidation state of Cr is + 6.

 

(iii) Permanganate, MnO^-₄

Oxidation state of Mn is + 7.

As we move along the lanthanoid series, the atomic number increases gradually by one. This means that the number of electrons and protons present in an atom also increases by one. As electrons are being added to the same shell, the effective nuclear charge increases. This happens because the increase in nuclear attraction due to the addition of proton is more pronounced than the increase in the interelectronic repulsions due to the addition of electron. Also, with the increase in atomic number, the number of electrons in the 4f orbital also increases. The 4f electrons have poor shielding effect. Therefore, the effective nuclear charge experienced by the outer electrons increases. Consequently, the attraction of the nucleus for the outermost electrons increases. This results in a steady decrease in the size of lanthanoids with the increase in the atomic number. This is termed as lanthanoid contraction.

Consequences of lanthanoid contraction

(i) There is similarity in the properties of second and third transition series.

(ii) Separation of lanthanoids is possible due to lanthanide contraction.

(iii) It is due to lanthanide contraction that there is variation in the basic strength of lanthanide hydroxides (Basic strength decreases from La (OH)₃ to Lu (OH)₃).

Transition elements are those elements in which the atoms or ions (in stable oxidation state) contain partially filled d-orbital. These elements lie in the d-block and show a transition of properties between s-block and p-block. Therefore, these are called transition elements.

Elements such as Zn, Cd, and Hg cannot be classified as transition elements because these have completely filled d-subshell.

Transition metals have a partially filled d-orbital. Therefore, the electronic configuration of transition elements is (n - 1)d^1-10 ns^0-2.

The non-transition elements either do not have a d-orbital or have a fully filled d-orbital. Therefore, the electronic configuration of non-transition elements is ns^1-2 or ns²np^1-6.

In the lanthanide series, +3 oxidation state is most common i.e., Ln (III) compounds are predominant. However, +2 and +4 oxidation states can also be found in the solution or in solid compounds.

(i) Transition metals show paramagnetic behaviour. Paramagnetism arises due to the presence of unpaired electrons with each electron having a magnetic moment associated with its spin angular momentum and orbital angular momentum. However, in the first transition series, the orbital angular momentum is quenched. Therefore, the resulting paramagnetism is only because of the unpaired electron.

(ii) Transition elements have high effective nuclear charge and a large number of valence electrons. Therefore, they form very strong metallic bonds. As a result, the enthalpy of atomization of transition metals is high.

(iii) Most of the complexes of transition metals are coloured. This is because of the absorption of radiation from visible light region to promote an electron from one of the d-orbitals to another. In the presence of ligands, the d-orbitals split up into two sets of orbitals having different energies. Therefore, the transition of electrons can take place from one set to another. The energy required for these transitions is quite small and falls in the visible region of radiation. The ions of transition metals absorb the radiation of a particular wavelength and the rest is reflected, imparting colour to the solution.

(iv) The catalytic activity of the transition elements can be explained by two basic facts.

(a) Owing to their ability to show variable oxidation states and form complexes, transition metals form unstable intermediate compounds. Thus, they provide a new path with lower activation energy, E_a, for the reaction.

(b) Transition metals also provide a suitable surface for the reactions to occur.

Transition metals are large in size and contain lots of interstitial sites. Transition elements can trap atoms of other elements (that have small atomic size), such as H, C, N, in the interstitial sites of their crystal lattices. The resulting compounds are called interstitial compounds.

In transition elements, the oxidation state can vary from +1 to the highest oxidation state by removing all its valence electrons. Also, in transition elements, the oxidation states differ by 1 (Fe²⁺and Fe³⁺; Cu⁺and Cu²⁺). In non-transition elements, the oxidation states differ by 2, for example, +2 and +4 or +3 and +5, etc.

Potassium dichromate is prepared from chromite ore (FeCr₂O₄) in the following steps.

Step (1):

Preparation of sodium chromate

4FeCr₂O₄     +   16NaOH   +   7O₂    →      8NaCrO₄    +   2Fe₂O₃  +  8H₂O

 

Step (2):

Conversion of sodium chromate into sodium dichromate

2Na₂CrO₄ + conc.H₂SO₄    →   Na₂Cr₂O₇  +  Na₂SO₄  +  H₂O

 

Step(3): Conversion of sodium dichromate to potassium dichromate

Na₂Cr₂O₇    +   2KCl       →       K₂Cr₂O₇   +   2NaCl

 

Potassium dichromate being less soluble than sodium chloride is obtained in the form of orange coloured crystals and can be removed by filtration.

The dichromate (Cr₂O^2-₇) ion exists in equilibrium with chromate (CrO^2-₄) ion at pH 4. However, by changing the pH, they can be interconverted.

K₂Cr₂O₇ acts as a very strong oxidising agent in the acidic medium.

K₂Cr₂O₇    +    4H₂SO₄   →    K₂SO₄   + Cr(SO₄)₃  +  4H₂O   +  3[O]   

K₂Cr₂O₇ takes up electrons to get reduced and acts as an oxidising agent. The reaction of K₂Cr₂O₇  with other iodide, iron (II) solution, and H2S are given below.

Potassium permanganate can be prepared from pyrolusite (MnO₂). The ore is fused with KOH in the presence of either atmospheric oxygen or an oxidising agent, such as KNO₃or KClO₄, to give K₂MnO₄.

 

The green mass can be extracted with water and then oxidized either electrolytically or by passing chlorine/ozone into the solution.

Electrolytic oxidation

K₂MnO₄   ↔    2K⁺   +    MnO^2-₄

H₂O     ↔   H⁺   +  OH^-

At anode, manganate ions are oxidized to permanganate ions.

MnO^2-₄       ↔     MnO^-₄    +   e^-

Green                 Purple 

Oxidation by chlorine

 

(i) The E^⊖ value for Fe³⁺/Fe^{2+ ­­} is higher than that for Cr³⁺/Cr²⁺and lower than that for Mn³⁺/Mn²⁺. So, the reduction of Fe³⁺to Fe²⁺is easier than the reduction of Mn³⁺to Mn²⁺, but not as easy as the reduction of Cr³⁺ to Cr²⁺. Hence, Fe³⁺is more stable than Mn³⁺, but less stable than Cr³⁺. These metal ions can be arranged in the increasing order of their stability as: Mn³⁺< Fe³⁺< Cr³⁺

 

(ii) The reduction potentials for the given pairs increase in the following order.

Mn²⁺ / Mn < Cr²⁺ / Cr < Fe²⁺ /Fe

So, the oxidation of Fe to Fe²⁺is not as easy as the oxidation of Cr to Cr²⁺and the oxidation of Mn to Mn²⁺. Thus, these metals can be arranged in the increasing order of their ability to get oxidised as: Fe < Cr < Mn value for Fe3+/ Fe2+­­ is higher than that for Cr3+/Cr2+and lower than that for Mn3+/Mn2+. So, the reduction of Fe3+to Fe2+is easier than the reduction of Mn3+to Mn2+, but not as easy as the reduction of Cr3+ to Cr2+. Hence, Fe3+is more stable than Mn3+, but less stable than Cr3+. These metal ions can be arranged in the increasing order of their stability as: Mn3+< Fe3+< Cr3+ (ii) The reduction potentials for the given pairs increase in the following order. Mn2+ / Mn < Cr2+ / Cr < Fe2+ /Fe So, the oxidation of Fe to Fe2+is not as easy as the oxidation of Cr to Cr2+and the oxidation of Mn to Mn2+. Thus, these metals can be arranged in the increasing order of their ability to get oxidised as: Fe < Cr < Mn

Only the ions that have electrons in d-orbital and in which d-d transition is possible will be coloured. The ions in which d-orbitals are empty or completely filled will be colourless as no d-d transition is possible in those configurations. 

Element	Atomic Number	Ionic State	Electronic configuration in ionic state
Ti	22	Ti3+	[Ar] 3d1
V	23	V3+	[Ar] 3d2
Cu	29	Cu+	[Ar] 3d10
Sc	21	Sc3+	[Ar]
Mn	25	Mn2+	[Ar] 3d5
Fe	26	Fe3+	[Ar] 3d5
Co	27	Co2+	[Ar] 3d7

From the above table, it can be easily observed that only Sc³⁺ has an empty d-orbital and Cu⁺ has completely filled d-orbitals. All other ions, except Sc³⁺ and Cu⁺, will be coloured in aqueous solution because of d - d transitions.

Sc			+3
Ti	+1	+2	+3	+4
V	+1	+2	+3	+4	+5
Cr	+1	+2	+3	+4	+5	+6
Mn	+1	+2	+3	+4	+5	+6	+7
Fe	+1	+2	+3	+4	+5	+6
Co	+1	+2	+3	+4	+5
Ni	+1	+2	+3	+4
Cu	+1	+2	+3
Zn		+2

From the above table, it is evident that the maximum number of oxidation states is shown by Mn, varying from +2 to +7. The number of oxidation states increases on moving from Sc to Mn. On moving from Mn to Zn, the number of oxidation states decreases due to a decrease in the number of available unpaired electrons. The relative stability of the +2 oxidation state increases on moving from top to bottom. This is because on moving from top to bottom, it becomes more and more difficult to remove the third electron from the d-orbital.

(i) Electronic configuration

The general electronic configuration for lanthanoids is [Xe]⁵⁴4f^0-145d^0-16s² and that for actinoids is [Rn]⁸⁶5f^1-146d^0-17s². Unlike 4forbitals, 5f orbitals are not deeply buried and participate in bonding to a greater extent.

(ii) Oxidation states

The principal oxidation state of lanthanoids is (+3). However, sometimes we also encounter oxidation states of + 2 and + 4. This is because of extra stability of fully-filled and half-filled orbitals. Actinoids exhibit a greater range of oxidation states. This is because the 5f, 6d, and 7s levels are of comparable energies. Again, (+3) is the principal oxidation state for actinoids. Actinoids such as lanthanoids have more compounds in +3 state than in +4 state.

(iii) Atomic and lonic sizes

Similar to lanthanoids, actinoids also exhibit actinoid contraction (overall decrease in atomic and ionic radii). The contraction is greater due to the poor shielding effect of 5f orbitals.

(iv) Chemical reactivity

In the lanthanide series, the earlier members of the series are more reactive. They have reactivity that is comparable to Ca. With an increase in the atomic number, the lanthanides start behaving similar to Al. Actinoids, on the other hand, are highly reactive metals, especially when they are finely divided. When they are added to boiling water, they give a mixture of oxide and hydride. Actinoids combine with most of the non-metals at moderate temperatures. Alkalies have no action on these actinoids. In case of acids, they are slightly affected by nitric acid (because of the formation of a protective oxide layer).

(i) Cr²⁺ is strongly reducing in nature. It has a d⁴ configuration. While acting as a reducing agent, it gets oxidized to Cr³⁺ (electronic configuration, d³). This d³ configuration can be written as t³_2g configuration, which is a more stable configuration. In the case of Mn³⁺(d⁴), it acts as an oxidizing agent and gets reduced to Mn²⁺(d⁵). This has an exactly half-filled d-orbital and is highly stable.

(ii) Co (II) is stable in aqueous solutions. However, in the presence of strong field complexing reagents, it is oxidized to Co (III). Although the 3^rdionization energy for Co is high, but the higher amount of crystal field stabilization energy (CFSE) released in the presence of strong field ligands overcomes this ionization energy.

(iii) The ions in d¹ configuration tend to lose one more electron to get into stable d⁰ configuration. Also, the hydration or lattice energy is more than sufficient to remove the only electron present in the d-orbital of these ions. Therefore, they act as reducing agents

It is found that sometimes a relatively less stable oxidation state undergoes an oxidation - reduction reaction in which it is simultaneously oxidised and reduced. This is called disproportionation.

For example,

In the first transition series, Cu exhibits +1 oxidation state very frequently. It is because Cu (+1) has an electronic configuration of [Ar] 3d¹⁰. The completely filled d-orbital makes it highly stable.

	Gaseous ions	Number of unpaired electrons
(i)	Mn3+ , [Ar] 3d4	4
(ii)	Cr3+ , [Ar] 3d3	3
(iii)	V3+ , [Ar] 3d2	2
(iv)	Ti3+ , [Ar] 3d1	1

Cr³⁺ is the most stable in aqueous solutions owing to a t³_2g configuration.

(i) In the case of a lower oxide of a transition metal, the metal atom has a low oxidation state. This means that some of the valence electrons of the metal atom are not involved in bonding. As a result, it can donate electrons and behave as a base.

On the other hand, in the case of a higher oxide of a transition metal, the metal atom has a high oxidation state. This means that the valence electrons are involved in bonding and they are unavailable. There is also a high effective nuclear charge.

As a result, it can accept electrons and behave as an acid.

For example,Mn^∥O is basic and Mn₂ ^vii O₇ is acidic.

(ii) Oxygen and fluorine act as strong oxidising agents because of their high electronegativities and small sizes. Hence, they bring out the highest oxidation states from the transition metals. In other words, a transition metal exhibits higher oxidation states in oxides and fluorides. For example, in OsF₆ and V₂O₅, the oxidation states of Os and V are +6 and +5 respectively.

(iii) Oxygen is a strong oxidising agent due to its high electronegativity and small size. So, oxo-anions of a metal have the highest oxidation state. For example, in MnO^-₄, the oxidation state of Mn is +7.

(i) Potassium dichromate (K₂Cr₂O₇) is prepared from chromite ore (FeCr₂O₄) in the following steps.

Potassium chloride being less soluble than sodium chloride is obtained in the form of orange coloured crystals and can be removed by filtration. The dichromate (Cr₂O²_-7) ion exists in equilibrium with chromate (CrO^2-₄) ion at pH 4. However, by changing the pH, they can be interconverted.

(ii) Potassium permanganate (KMnO₄) can be prepared from pyrolusite (MnO2). The ore is fused with KOH in the presence of either atmospheric oxygen or an oxidising agent, such as KNO₃or KClO₄, to give K₂MnO₄.

The green mass can be extracted with water and then oxidized either electrolytically or by passing chlorine/ozone into the solution.

Electrolytic oxidation

 

An alloy is a solid solution of two or more elements in a metallic matrix. It can either be a partial solid solution or a complete solid solution. Alloys are usually found to possess different physical properties than those of the component elements.

An important alloy of lanthanoids is Mischmetal. It contains lanthanoids (94- 95%), iron (5%), and traces of S, C, Si, Ca, and Al.

Uses

(1) Mischmetal is used in cigarettes and gas lighters.

(2) It is used in flame throwing tanks.

(3) It is used in tracer bullets and shells.

Inner transition metals are those elements in which the last electron enters the f-orbital. The elements in which the 4f and the 5f orbitals are progressively filled are called f-block elements. Among the given atomic numbers, the atomic numbers of the inner transition elements are 59, 95, and 102.

Lanthanoids primarily show three oxidation states (+2, +3, +4). Among these oxidation states, +3 state is the most common. Lanthanoids display a limited number of oxidation states because the energy difference between 4f, 5d, and 6s orbitals is quite large. On the other hand, the energy difference between 5f, 6d, and 7s orbitals is very less. Hence, actinoids display a large number of oxidation states. For example, uranium and plutonium display +3, +4, +5, and +6 oxidation states while neptunium displays +3, +4, +5, and +7. The most common oxidation state in case of actinoids is also +3.

The last element in the actinoid series is lawrencium, Lr. Its atomic number is 103 and its electronic configuration is [Rn]5 f¹⁴6d¹7s². The most common oxidation state displayed by it is +3; because after losing 3 electrons it attains stable f¹⁴ configuration.

The lanthanides that exhibit +2 and +4 states are shown in the given table. The atomic numbers of the elements are given in the parenthesis.

+2	+4
Nd(60)	Ce(58)
Sm(62)	Pr(59)
Eu(63)	Nd(60)
Tm(69)	Tb(65)
Yb(70)	Dy(66)

 

Ce after forming Ce⁴⁺ attains a stable electronic configuration of [Xe].

Tb after forming Tb⁴⁺ attains a stable electronic configuration of [Xe] 4f^7.

Eu after forming Eu²⁺ attains a stable electronic configuration of [Xe] 4f^7.

Yb after forming Yb²⁺ attains a stable electronic configuration of [Xe] 4f^14.

Electronic configuration

The general electronic configuration for lanthanoids is [Xe]⁵⁴4f^0-145d^0-16s² and that for actinoids is [Rn]⁸⁶5f^1-146d^0-17s². Unlike 4 forbitals, 5f orbitals are not deeply buried and participate in bonding to a greater extent.

 

Oxidation states

The principal oxidation state of lanthanoids is (+3). However, sometimes we also encounter oxidation states of + 2 and + 4. This is because of extra stability of fully-filled and half-filled orbitals. Actinoids exhibit a greater range of oxidation states. This is because the 5f, 6d, and 7slevels are of comparable energies. Again, (+3) is the principal oxidation state for actinoids. Actinoids such as lanthanoids have more compounds in +3 state than in +4 state.

 

Chemical reactivity

In the lanthanide series, the earlier members of the series are more reactive. They have reactivity that is comparable to Ca. With an increase in the atomic number, the lanthanides start behaving similar to Al. Actinoids, on the other hand, are highly reactive metals, especially when they are finely divided. When they are added to boiling water, they give a mixture of oxide and hydride. Actinoids combine with most of the non-metals at moderate temperatures. Alkalies have no action on these actinoids. In case of acids, they are slightly affected by nitric acid (because of the formation of a protective oxide layer).

Atomic number	Electronic configuration
61	[Xe]54 4f5 5d0 6s2
91	[Rn]86 5f2 6d1 7s2
101	[Rn]86 5f13 5d0 7s2
109	[Rn]86 5f14 6d7 7s2

 

(i) In the 1^st, 2^nd and 3^rd transition series, the 3d, 4d and 5dorbitals are respectively filled.

We know that elements in the same vertical column generally have similar electronic configurations.

In the first transition series, two elements show unusual electronic configurations:

Cr(24) = 3d⁵ 4s¹

Cu(29) = 3d¹⁰ 4s¹

Similarly, there are exceptions in the second transition series. These are:

Mo(42) = 4d⁵ 5s¹

Tc(43) = 4d⁶ 5s¹

Ru(44) = 4d⁷ 5s¹

Rh(45) = 4d⁸ 5s¹

Pd(46) = 4d¹⁰ 5s⁰

Ag(47) = 4d¹⁰ 5s¹

There are some exceptions in the third transition series as well. These are:

W(74) = 5d⁴ 6s²

Pt(78) = 5d⁹ 6s¹

Au(79) = 5d¹⁰ 6s¹

As a result of these exceptions, it happens many times that the electronic configurations of the elements present in the same group are dissimilar.

 

(ii) In each of the three transition series the number of oxidation states shown by the elements is the maximum in the middle and the minimum at the extreme ends.

However, +2 and +3 oxidation states are quite stable for all elements present in the first transition series. All metals present in the first transition series form stable compounds in the +2 and +3 oxidation states. The stability of the +2 and +3 oxidation states decreases in the second and the third transition series, wherein higher oxidation states are more important.

For example  are stable complexes, but no such complexes are known for the second and third transition series such as Mo, W, Rh, In. They form complexes in which their oxidation states are high. For example: WCl₆, ReF₇, RuO₄, etc.

 

(iii) In each of the three transition series, the first ionisation enthalpy increases from left to right. However, there are some exceptions. The first ionisation enthalpies of the third transition series are higher than those of the first and second transition series. This occurs due to the poor shielding effect of 4felectrons in the third transition series.

Certain elements in the second transition series have higher first ionisation enthalpies than elements corresponding to the same vertical column in the first transition series. There are also elements in the 2^nd transition series whose first ionisation enthalpies are lower than those of the elements corresponding to the same vertical column in the 1^st transition series.

 

(iv) Atomic size generally decreases from left to right across a period. Now, among the three transition series, atomic sizes of the elements in the second transition series are greater than those of the elements corresponding to the same vertical column in the first transition series. However, the atomic sizes of the elements in the third transition series are virtually the same as those of the corresponding members in the second transition series. This is due to lanthanoid contraction.

Metal ion	Number of d-electrons	Filling of d-orbitals
Ti2+	2	t22g
V2+	3	t32g
Cr3+	3	t32g
Mn2+	5	t32g e2g
Fe2+	6	t42g e2g
Fe3+	5	t32g e2g
CO2+	7	t52g e2g
Ni2+	8	t62g e2g
Cu2+	9	t62g e3g

 

The properties of the elements of the first transition series differ from those of the heavier transition elements in many ways.

(i) The atomic sizes of the elements of the first transition series are smaller than those of the heavier elements (elements of 2^nd and 3^rd transition series).

However, the atomic sizes of the elements in the third transition series are virtually the same as those of the corresponding members in the second transition series. This is due to lanthanoid contraction.

(ii) +2 and +3 oxidation states are more common for elements in the first transition series, while higher oxidation states are more common for the heavier elements.

(iii) The enthalpies of atomisation of the elements in the first transition series are lower than those of the corresponding elements in the second and third transition series.

(iv) The melting and boiling points of the first transition series are lower than those of the heavier transition elements. This is because of the occurrence of stronger metallic bonding (M-M bonding).

(v) The elements of the first transition series form low-spin or high-spin complexes depending upon the strength of the ligand field. However, the heavier transition elements form only low-spin complexes, irrespective of the strength of the ligand field.

(i) K₄[Mn(CN)₆]

For in transition metals, the magnetic moment is calculated from the spin-only formula. Therefore,

We can see from the above calculation that the given value is closest to n = 1. Also, in this complex, Mn is in the +2 oxidation state. This means that Mn has 5 electrons in the d-orbital.

Hence, we can say that CN ^- is a strong field ligand that causes the pairing of electrons.

 

(ii) [Fe(H₂O)₆]²⁺

We can see from the above calculation that the given value is closest to n = 4. Also, in this complex, Fe is in the +2 oxidation state. This means that Fe has 6 electrons in the d-orbital.

Hence, we can say that H₂O is a weak field ligand and does not cause the pairing of electrons.

(iii) K₂[MnCl₄]

We can see from the above calculation that the given value is closest to n = 5. Also, in this complex, Mn is in the +2 oxidation state. This means that Mn has 5 electrons in the d-orbital.

Hence, we can say that Cl ^- is a weak field ligand and does not cause the pairing of electrons.