INTRODUCTION TO D-BLOCK CHEMISTRY FOR JEE & NEET

Introduction

         Three series of elements are formed by filling the 3d, 4d and 5d-subshells of electrons. Together these comprise the d-block elements. They are often called ‘transition elements’ because their position in the periodic table is between the s-block and p-block elements. Their properties are transitional between the highly reactive metallic elements of the s-block, which typically forms ionic compounds and the elements of the p-block, which are largely covalent. In the s and p-blocks, electrons are added to the outer shell of the atom. In the d-block, electrons are added to the penultimate shell, expanding it from 8 to 18 electrons. Typically the transition elements have an incompletely filled d-level. Group 12 (the zinc group) has a d¹⁰ configuration and since the d-subshell is complete, compounds of these elements are not typical and show some differences from the others. The transition elements make up three complete rows of ten elements and an incomplete fourth row.

         The transition elements are defined as those elements, which have partly filled
d-orbitals, as elements and in any of their important compounds.

         The general electronic configuration of the d-block elements can be represented as, (n-1)d^1-9ns^1-2

         Depending on the subshell getting filled up the transition elements form three series.
The first transition series contain the elements from Sc (Z = 21) to Zn (Z = 30) and the 3d-orbital gets filled up in this series. In the second series, the 4d-orbital gets filled up from Y (Z = 39) to Cd (Z = 48). The 5d-orbital from La (Z = 57) to Hg (Z = 80) gets filled up for the elements of the third series. The fourth series starting with Ac is incomplete.

Group No.	3	4	5	6	7	8	9	10	11	12
At. No. (Z) Symbol	21 Sc	22 Ti	23 V	24 Cr	25 Mn	26 Fe	27 Co	28 Ni	29 Cu	30 Zn
At. No. (Z) Symbol	39 Y	40 Zr	41 Nb	42 Mo	43 Tc	44 Ru	45 Rh	46 Pd	47 Ag	48 Cd
At. No. (Z) Symbol	57 La *	72 Hf	73 Ta	74 W	75 Re	76 Os	77 Ir	78 Pt	79 Au	80 Hg
At. No. (Z) Symbol	89 Ac *

                                                14 Lanthanide elements

                                    14 Actinide elements

         Unlike the s and p-block elements of the same period, the d-block elements do not show much variation in properties, both chemical and physical. This is because these elements differ only in the number of electrons in the penultimate d-shell. The number of electrons in the valence shell remains the same, ns², for most of the elements.

1.1    mETALLIC CHARACTER

         In the d-block elements, the penultimate shell of electrons is expanding. Thus, they have many physical and chemical properties in common. Thus, all the transition elements are metals. They are good conductors of heat and electricity, have a metallic luster and are hard, strong and ductile. They also form alloys with other metals. Copper exceptionally is both soft and ductile and relatively noble.

1.2    VARIABLE OXIDATION STATES

         One of the most striking features of the transition elements is that the elements usually exist in several different oxidation states and the oxidation states change in units of one.
For example: Fe³⁺ and Fe²⁺, Cu²⁺ and Cu⁺ etc.

         The oxidation states shown by the transition elements may be related to their electronic configurations. Calcium, the s-block element preceding the first row of transition elements, has the electronic configuration:

                       Ca (Z = 20):      1s²2s²2p⁶3s²3p⁶4s² :  [Ar]4s²

         It might be expected that the next ten transition elements would have this electronic arrangement with from one to ten d-electrons added in a regular way: 3d¹, 3d², 3d³…3d¹⁰. This is true except in the cases of Cr and Cu. In these two cases, one of the s-electrons moves into the d-subshell, because of the additional stability of the exactly half-filled or completely filled d-orbital. Since the energies of (n-1) d and ns-orbitals are nearly equal, the transition elements exhibit variable oxidation states. The oxidation states of the d-block elements are listed below.

Electronic configuration	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
	3d14s2	3d24s2	3d34s2	3d44s2 3d54s1	3d54s2	3d64s2	3d74s2	3d84s2	3d94s2 3d104s1	3d104s2
Oxidation states				I					I
	II	II	II	II	II	II	II	II	II	II
	III	III	III	III	III	III	III	III	III
		IV	IV	IV	IV	IV	IV	IV
			V	V	V	V	V
				VI	VI	VI
					VII

         Thus, Sc could have an oxidation state of (II) if both s-electrons are used for bonding and (III) when two s and one d-electrons are involved. Ti has an oxidation state (II) when both s-electrons are used for bonding, (III) when two s and one d-electrons are used and (IV) when two s and two d-electrons are used. Similarly, V shows oxidation numbers (II), (III), (IV) and (V). In the case of Cr, by using the single s-electron for bonding, we get an oxidation number of (I); hence by using varying number of d-electrons, oxidation states of (II), (III), (IV), (V) and (VI) are possible. Mn has oxidation states (II), (III), (IV), (V), (VI) and (VII). Among these first five elements, the correlation between electronic configuration and minimum and maximum oxidation states is simple and straight forward. In the highest oxidation states of these first five elements, all of the s and d-electrons are being used for bonding. Thus, the properties depend only on the size and valency.

         Once the d⁵ configuration is exceeded, i.e. in the last five elements, the tendency for all the d-electrons to participate in bonding decreases. Thus, Fe has a maximum oxidation state of (VI). However, the second and third elements in this group attain a maximum oxidation state of (VIII), as in RuO₄ and OsO₄. This difference between Fe and the other two elements (Ru and Os) is attributed to the increased size and decreased attraction with the nucleus.

         The oxidation number of all elements in the elemental state is zero. In addition, several of the elements have zero-valent and other low-valent states in complexes. Low oxidation states occur particularly with p-bonding ligands such as carbon monoxide and dipyridyl.

         Some other important features about the oxidation states of transition elements can be outlined as:

         1.     In group 8 (the iron group), the second and third row elements show a maximum oxidation state of (VIII) compared with (VI) for Fe.

         2.     The electronic configurations of the atoms in the second and third rows do not always follow the pattern of the first row. The configurations of group 10 elements (the nickel group) are:

                 Ni (Z = 28) :           3d⁸4s²

                 Pd (Z = 46) :           4d¹⁰5s⁰

                 Pt (Z = 78)  :           5d⁹6s¹

         3.     Since a full shell of electrons is a stable arrangement, the place where this occurs is of importance in the transition series. The d-levels are complete at copper, palladium and gold in their respective series.

                 Ni  : 3d⁸4s²             Cu :     3d¹⁰4s¹            Zn :     3d¹⁰4s²

                 Pd : 4d¹⁰5s⁰             Ag :     4d¹⁰5s¹            Cd :     4d¹⁰5s²

                 Pt:  5d⁹6s¹               Au :     5d¹⁰6s¹            Hg :     5d¹⁰6s²

         4.     Even though the ground state of the atom has a d¹⁰ configuration, Pd and the coinage metals Cu, Ag and Au behave as typical transition elements. This is because in their most common oxidation states, Cu(II) has a d⁹ configuration and Pd(II) and Au(III) have d⁸ configurations, that is they have an incompletely filled d-level. However, in zinc, cadmium and mercury, the ions Zn²⁺, Cd²⁺ and Hg²⁺ have a d¹⁰ configuration. Because of this, these elements do not show the properties characteristic of transition elements.

         Compounds are regarded as stable if they exist at room temperature, are not oxidized by the air, are not hydrolysed by water vapour and do not disproportionate or decompose at normal temperatures. Within each of the transition metals of groups 3-12, there is a difference in stability of the various oxidation states that exist. In general, the second and third row elements exhibit higher co-ordination numbers and their higher oxidation states are more stable than the corresponding first row elements. Stable oxidation states form oxides, fluorides, chlorides, bromides and iodides. Strongly reducing states probably do not form fluorides and/or oxides, but may well form the heavier halides. Conversely, strongly oxidizing states form oxides and fluorides, but not iodides.

                 Oxides and halides of some elements of the first row:

	Cr	Mn	Fe
II   O	CrO	MnO	FeO
F	CrF2	MnF2	FeF2
Cl	CrCl2	MnCl2	FeCl2
Br	CrBr2	MnBr2	FeBr2
I	CrI2	MnI2	FeI2
III   O	Cr2O3	Mn2O3	Fe2O3
F	CrF3	MnF3	FeF3
Cl	CrCl3	-	FeCl3
Br	CrBr3	-	FeBr3
I	CrI3	-	-
IV   O	CrO2	MnO2	-
F	CrF4	MnF4	-
Cl	CrCl4	-	-
Br	CrBr4	-	-
I	CrI4	-	-
V     F	CrF5	-	-
VI    O	CrO3	-	-
F	CrF6	-	-
VII  O	-	Mn2O7	-

1.3    ATOMIC AND IONIC RADII

         The covalent radii of the elements decrease from left to right across a row in the transition series, until near the end when the size increases slightly. On passing from left to right, extra protons are placed in the nucleus and extra orbital electrons are added. The orbital electrons shield the nuclear charge incompletely (d-electrons shield less efficiently than p-electrons, which in turn shield less effectively than s-electrons). Because of this poor screening by d-electrons, the nuclear charge attracts all of the electrons more strongly, hence a contraction in size occurs.

         The elements in the first group in the d-block (group 3) show the expected increase in size Sc  ¾®  Y  ¾®  La.  However, in the subsequent groups (4-12) there is an increase in radius of 0.1 ¾® 0.2 Å between the first and second member, but hardly any increase between the second and third elements. This trend is shown both in the covalent radii and in the ionic radii. Interposed between lanthanum and hafnium are the 14 lanthanide elements, in which the antipenultimate 4f-subshell of electrons is filled.

         There is a gradual decrease in size of the 14 lanthanide elements from cerium to lutetium. This is called the “lanthanide contraction”. The lanthanide contraction cancels almost exactly the normal size increase on descending a group of transition elements. Therefore, the second and third row transition elements have similar radii. As a result they also have similar lattice energies, solvation energies and ionization energies. Thus, the differences in properties between the first row and second row elements are much greater than the differences between the second and third row elements. The effects of the lanthanide contraction are less pronounced towards the right of the d-block. However, the effect still shows to a lesser degree in the p-block elements that follow.

         Covalent radii of the transition elements (in Å)

K 1.57	Ca 1.74	Sc 1.44	Ti 1.32	V 1.22	Cr 1.17	Mn 1.17	Fe 1.17	Co 1.16	Ni 1.15	Cu 1.17	Zn 1.25
Rb 2.16	Sr 1.91	Y 1.62	Zr 1.45	Nb 1.34	Mo 1.29	Tc -	Ru 1.24	Rh 1.25	Pd 1.28	Ag 1.34	Cd 1.41
Cs 2.35	Ba 1.98	La * 1.69	Hf 1.44	Ta 1.34	W 1.30	Re 1.28	Os 1.26	Ir 1.26	Pt 1.29	Au 1.34	Hg 1.44

                                         14 Lanthanide elements

1.4    DENSITY

         The atomic volumes of the transition elements are low compared with elements in neighbouring groups 1 and 2. This is because the increased nuclear charge is poorly screened and so attracts all the electrons more strongly. In addition, the extra electrons added occupy inner orbitals. Consequently, the densities of the transition metals are high. Practically, most of the elements have a density greater than 5 g cm^-3. (The only exceptions are Sc: 3.0 g cm^-3 and Y and Ti: 4.5 g cm^-3). The densities of the second row elements are high and third row values are even higher. The two elements with the highest densities are osmium: 22.57 g cm^-3 and iridium: 22.61 g cm^-3. Thus, iridium is the heaviest element among all the elements of the periodic table.

1.5    MELTING AND BOILING POINTS

         The melting and boiling points of the transition elements are generally very high. Transition elements typically melt above 1000°C. Ten elements melt above 2000°C and three melt above 3000°C (Ta: 3000°C, W: 3410°C and Re: 3180°C). There are a few exceptions.
For example, La and Ag melts just under 1000°C (920°C and 960°C respectively). Other notable exceptions are Zn (420°C), Cd(320°C) and Hg, which is liquid at room temperature and melts at -38°C. The last three do not behave as typical transition elements,  because the d-subshell is complete and d-electrons do not participate in metallic bonding.

1.6    IONIZATIONENERGY

         In a period, the first ionization energy gradually increases from left to right. This is mainly due to increase in nuclear charge. Generally, the ionization energies of transition elements are intermediate between those of s and p-block elements. The first ionization potential of the
5d-elements are higher than those of 3d and 4d-elements due to the poor shielding by 4f-electrons.

         From 3d  ¾®  4d series, general trend is observed but not from 4d ¾® 5d series because of incorporation of the 14 lanthanides elements between La and Hf. Third period of transition elements have the highest ionisation energy. This reflects the fact that increase in radius due to addition of extra shell is compensated by the decrease in radius due to lanthanide contraction.

         As the radius of 4d and 5d-elements more or less remains the same, due to which Z_eff of elements of 5d series is higher, which results in high ionization energy of the 5d-elements of transition series.

         The ionisation energy values (in kJ/mole) of the transitions elements are given in
the table below:

3d series	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
First I. E	631	656	650	652	717	762	758	736	745	906
4d series	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd
First I. E	616	674	664	685	703	711	720	804	731	876
5d series	La	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg
First I. E	541	760	760	770	759	840	900	870	889	1007

1.7    RECTIVITY OF METALS

         Many of the metals are sufficiently electropositive to react with mineral acids, liberating H₂. A few have low standard electrode potentials and remain unreactive or noble. Noble character is favoured by high enthalpies of sublimation, high ionization energies and low enthalpies of solvation. The high melting points indicate high heats of sublimation. The smaller atoms have higher ionization energies, but this is offset by small ions having high solvation energies. This tendency to noble character is most pronounced for the platinum metals (Ru, Rh, Pd, Os, Ir, Pt) and gold.

1.8    FORMATION OF COMPLEX COMPOUNDS

         The transition elements have characteristic tendency to form co-ordination compounds with Lewis bases, that is with groups that are able to donate an electron pair. These groups are called ligands. A ligand may be a neutral molecule such as NH₃, or an ion such as Cl^- or CN^-. Cobalt forms more complexes than any other element and forms more compounds than any other element after carbon.

                       Co³⁺  +  6NH₃     [Co(NH₃)₆]³⁺

                       Fe²⁺  +  6CN^-   [Fe(CN)₆]^4-

         This ability to form complexes is in marked contrast to the s- and p-block elements, which form only a few complexes. The reason transition elements are so good at forming complexes is that they have small, highly charged ions and have vacant low energy orbitals to accept lone pairs of electrons donated by other groups or ligands.

1.9    COLOUR OF COMPOUNDS

         Many ionic and covalent compounds of transition elements are coloured. In contrast compounds of the s- and p-block elements are almost always white. When light passes through a material, it is deprived of those wavelengths that are absorbed. If wavelength of the absorption occurs in the visible region of the spectrum, the transmitted light is coloured with the complementary colour to the colour of the light absorbed. Absorption in the visible and UV regions of the spectrum is caused by changes in electronic energy. Thus, the spectra are sometimes called electronic spectra.

         Colour may arise from an entirely different cause in ions with incomplete d or f-subshells. This source of colour is very important in most of the transition metal ions.

         In a free isolated gaseous ion, the five d-orbitals are degenerate that is they are identical in energy. In actual practice, the ion will be surrounded by solvent molecules if it is in solution, by other ligands if it is in a complex, or by other ions if it is in a crystal lattice. The surrounding groups affect the energy of some d-orbitals more than others. Thus, the d-orbitals are no longer degenerate and at their simplest they form two groups of orbitals of different energy. Thus, in transition element ions with a partly filled d-subshell it is possible to promote electrons from one d-level to another d-level of higher energy. This corresponds to a fairly small energy difference and so light it absorbed in the visible region. The colour of a transition metal complex is dependent on how big the energy difference is between the two d-levels. This in turn depends on the nature of the ligand and on the type of complex formed. Thus, the octahedral complex [Ni(NH₃)₆]²⁺ is blue, [Ni(H₂O)₆]²⁺ is green and [Ni(NO₂)₆]^4- is brown-red. The colour changes with the ligand used. The colour also depends on the number of ligands and the shape of the complex formed.

         The source of colour in the lanthanides and the actinides is very similar, arising from
f ® f transitions. With the lanthanides, the 4f-orbitals are deeply embedded inside the atom and are well-shielded by the 5s and 5p-electrons. The f-electrons are practically unaffected by complex formation. Hence, the colour remains almost constant for the particular ion regardless of the ligand.

         Some compounds of the transition metal are white, for example Cu₂Cl₂, ZnSO₄ and TiO₂. In these compounds, it is not possible to promote electrons within the d-level. Cu⁺ and Zn²⁺ has a d¹⁰ configuration and the d-level is completely filled. Ti⁴⁺ has a d⁰ configuration and the d-level is empty. In the series Sc(III), Ti(IV), V(V), Cr(VI) and Mn(VII), these ions may all be considered to have an empty d-subshell; hence d-d spectra are impossible and they should be colourless. However, as the oxidation state increases, these states become increasingly covalent. Rather than forming highly charged simple ions, they form oxoions like TiO²⁺, , ,  and .   is pale yellow, but  is strongly yellow coloured and  has an intense purple colour in solution, though the solid is almost black. The colour arises by charge transfer mechanism.
In , an electron is momentarily transferred from O to the metal, thus momentarily changing O^2- to O^- and reducing the oxidation state of the metal from Mn(VII) to Mn(VI). Charge transfer requires the energy levels on the two different atoms to be fairly close. Charge transfer always produces more intense colours than the colours generated due to d-d transitions. Charge transfer is also possible between metal-ion and metal-ion as seen in prussian blue, Fe₄[Fe(CN)₆]₃.

         The s and p-block elements do not have a partially filled d-subshell, so there cannot be any d-d-transitions. The energy required to promote an s or p-electron to a higher energy level is much greater and corresponds to ultraviolet light being absorbed. Thus, compounds of s and p-block elements are typically not coloured.

1.10  MAGNETIC PROPERTIES

         Compounds of the transition elements exhibit characteristic magnetic behaviour. Those, which are attracted by a magnetic field, are termed as paramagnetic. Those, which are repelled by a magnetic field, are called diamagnetic. Paramagnetic species have unpaired electrons in their electronic configuration. Diamagnetic substances are those in which electrons are fully paired. In a simple situation, where one may consider aquocomplex ions, we have the following formulation.

Metal ion	Electronic configuration	No. of unpaired e-’s	Metal ion	Electronic configuration	No. of unpaired e-’s
Sc3+	3d0	No unpaired electrons	Ti3+	3d1	1 unpaired electron
V3+	3d2	2 unpaired electrons	Cr3+	3d3	3 unpaired electrons
Mn3+	3d4	4 unpaired electrons	Fe3+	3d5	5 unpaired electrons
Mn2+	3d5	5 unpaired electrons	Fe2+	3d6	4 unpaired electrons
Co2+	3d7	3 unpaired electrons	Ni2+	3d8	2 unpaired electrons
Cu2+	3d9	1 unpaired electron	Cu+	3d10	No unpaired electrons
Zn2+	3d10	No unpaired electrons

         Unpaired electrons in any species have, each, a spin angular momentum, which can be vectorially added to yield a resultant spin angular momentum. This gives rise to a magnetic moment. Actually, there are two contributions to the magnetic moment i.e., the magnetic moment due to orbital angular momentum and the spin magnetic moment. In many situations, the environment in which a species is located has the effect of quenching out the orbital contribution.          

         Thus, in such cases, only the spin magnetic moment is measured; in units of Bohr magneton. The spin magnetic moment is given by  in BM, where n is the number of unpaired electrons.

Note: Bohr magneton has the value; BM =  where e = magnitude of electronic charge,
m_o = rest mass of the electron and c = speed of light in vacuum. Typical values are Ti³⁺, 3d¹,  BM = 1.73 BM and this agrees with the measured value. In many cases, the observed and calculated values in the spin magnetic moment are in fair agreement. In fact, determination of spin magnetic moment helps us to know the number of unpaired electrons in the complex/complex ion, which leads us to the bonding and structure elucidation of the complex/complex ion.