The d block elements are found in the third to twelfth groups of the modern periodic table. These elements have valence electrons that fall under the d orbital, and are also referred to as transition elements or transition metals. The first three rows of the d block elements correspond to the 3d, 4d, and 5d orbitals, respectively, and are discussed in more detail in the article below.

Table of Contents

Position of D Block Elements

D Block as Transition Element

Electronic Configuration of D-Block Elements

Atomic and Ionic Radii of D-Block Elements

Properties of D-Block Elements

Oxidation States of D-Block Elements

Formation of Coloured Ions by D-Block Elements

Alloy Formation in D-Block Elements

Important Compounds of D-Block Elements

D Block Elements are the elements that make up the groups 3-12 of the periodic table. They are also known as the transition metals and are characterized by having partially filled d orbitals.

Elements with electrons present in the d-orbital of the penultimate energy level (1 to 10) and in the outermost ’s’ orbital (1-2) are considered as d block elements. Even though electrons do not fill up the ’d’ orbital in the group 12 metals, their chemistry is similar to the preceding groups and thus they are classified as d block elements.

Elements in the d-block of the periodic table typically display metallic qualities such as malleability, ductility, high electrical and thermal conductivity, and good tensile strength. There are four series in the d-block, which correspond to the filling of the 3d, 4d, 5d, and 6d orbitals.

3d: Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn

4d- Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd

5d- La, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg

Incomplete: 6d

The ’d’ orbital in each series is filled with 10 elements.

Also Read

F Block Elements

Actinides

Position of D Block Elements in the Periodic Table

IUPAC defines a transition metal as an element whose atom or its cations has a partially filled d sub-shell. D block elements, which occupy columns 3 to 12, may have atoms of elements with completely filled d orbitals.

What Makes D Block Elements Transition Elements?

Transition elements occupy groups 4–11, while Scandium and Yttrium of group 3, which have a partially filled d subshell in the metallic state, are also considered transition elements. Elements like Zn, Cd and Hg of the 12 columns of the d block, however, have a completely filled d-orbital and hence are not considered as transition elements.

All the transition metals are d block elements, but not all d block elements are transition elements.

Properties of Transition Metals

Electrons added to the ’d’ sub-orbitals that lie between the (n+1)s and (n+1)p sub-orbitals.

Placed between s and p block elements in the periodic table

Properties of s and p-block Elements

Electronic Configuration of D Block Elements

The electronic configuration of D block Elements is generally (n-1)d 1-10ns 1-2. This configuration can lead to stability when the d and s orbitals are half-filled or completely filled. An example of this is chromium, which has the electronic configuration of 3d54s1, with half-filled d and s orbitals. Copper also has a similar electronic configuration, which is 3d104s1 and not 3d94s2. Here is a resource to learn more about electron configurations.

These metals, such as Zinc, Mercury, Cadmium, and Copernicium, are not considered as transition elements due to their completely filled d orbitals in their ground states and in their general oxidation states. In contrast, the other d block elements do not exhibit completely filled orbitals.

The electronic configuration for period 4, transition elements is [Ar] 4s^1-2 3d^1-10

The electronic configuration for period 5, transition elements is: (Kr) 5s^1-2 4d^1-10

The electronic configuration for period 6, transition elements is [Xe] 4s\textsuperscript{1-2} 3d\textsuperscript{1-10}

Along the period, electrons are added to the 3$^{d}$ subshell in accordance with the Aufbau Principle and Hund’s Rule of Multiplicity, from left to right.

| 1st Transition Series | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn |

| 4s23d1 | 4s23d2 | 4s23d3 | 4s23d5 | 4s23d5 | 4s23d6 | 4s23d7 | 4s23d8 | 4s23d10 | 4s23d10 |

| 2nd Transition Series | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd |

| 5s24d1 | 5s24d2 | 5s14d4 | 5s14d5 | 5s24d5 | 5s14d7 | 5s14d8 | 5s04d10 | 5s14d10 | 5s24d10 |

| 3rd Transition Series | La | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg |

| 6s25d1 | 6s25d2 | 6s25d3 | 6s25d4 | 6s25d5 | 6s25d6 | 6s25d7 | 6s15d9 | 6s15d10 | 6s25d10 |

Anomalies can be seen in all series, and can be explained by the following considerations.

The Energy Gap Between the ns and (n-1)d Orbitals

Pairing energy for the electrons in s-orbital

Stability of half-filled orbitals compared to partly filled orbitals.

Chromium has a 4s13d5 electron configuration, as opposed to the 4s23d4 configuration, and copper has a 4s13d10 configuration, instead of the 4s23d9 configuration. These anomalies in the first transition series can be explained by the stability of half-filled orbitals compared to the partly filled orbitals.

In the second series of transition metals, starting from niobium, electrons tend to occupy d-orbitals rather than being shared in s-orbitals. The electron can either opt for sharing in the s-orbital or be excited to the d-orbital, depending on the repulsive energy it needs to overcome and the energy gap between the s and d-orbitals.

In the second series, s and d-orbitals have almost the same energy, which leads to electrons preferring the d-orbitals. As a result, s-orbitals of Niobium mostly contain only one electron. On the other hand, the third series transition metals have more paired s configurations, even at the expense of half-filled orbitals (e.g. Tungsten - 6s25d4). This series follows the filling up of the 4f orbitals and the resulting lanthanide contraction.

The smaller size of tungsten leads to a higher shielding of its d orbitals by the ‘f’ electron. This shielding increases the energy gap between the s and 5d orbitals, resulting in pairing energy being less than the excitation. Thus, the excitation of the electron does not take place in tungsten, despite the stability it could gain from having half-filled orbitals.

Atomic and Ionic Radii of D Block Elements

The Metallic Radii of the 1st, 2nd, and 3rd Row Transition Metals are shown in the image.

Atomic and Ionic Radii of Elements in the Three-Transition Series

Decreases rapidly, from column 3 to 6

Remains steady, from column 7 to 10

Increases from column 11 to 12.

The decrease in atomic radii from Sc (Group 3) to Cr (Group 6) is quite drastic, whereas from Mn (Group 7) to Zn (Group 12) the change is not so pronounced.

The larger decrease in atomic radii, from column 3 to 6 elements, is due to an increase in effective nuclear charge but poor shielding because of the smaller number of d-electrons.

In elements of column 7 to 10, the increasing effective nuclear charge is balanced by the repulsion between the shared d electrons, resulting in radii that remain the same.

In the case of elements in groups 11 and 12, the d-orbital is filled with 10 electrons which shield the electrons in the higher s-orbital. As a result, elements such as Cu and Zn in groups 11 and 12 are larger in size than the elements in the earlier groups of the block.

Since electrons occupy a higher orbital, the radii of the third series of elements are expected to be larger than those of the second series. However, the radii of both series are nearly the same. This is because in the third series of elements, the 5d orbitals are filled only after the 4f orbitals are filled, which increases the effective nuclear charge by 14 units.

The Lanthanide contraction is caused by the increase of nuclear charge, which counteracts the increase of atomic radii due to higher orbitals. Therefore, the atomic radii of the second and third series elements remain the same. For example, Niobium and Hafnium have almost the same atomic radii.

Properties of D Block Elements

Ionization Energy of D Block Elements

Ionization energy is the energy required to remove the valence electron from the atom/ion and is directly proportional to the force of attraction on the electron. Therefore, a larger nuclear charge and a smaller radius of the electron will lead to a greater ionization energy. Ionization Energy is also higher for half-filled and fully filled orbitals.

The D-block elements have an ionization energy that is larger than that of the S-block elements, and smaller than that of the P-block elements, between which they are placed. In the first series, except for chromium and copper, the first ionization energy involves removal from a filled s-orbital. Among them, the ionization energy of D-block elements increases with the increase in atomic number up to iron.

In Co and Ni, an increase in the sharing of d-electrons compensates for the increase in atomic number, resulting in a decrease in Ionization Energy. Copper and zinc, however, show an increase in IE, as s-block elements. Additionally, elements from Niobium onwards have single electrons in the s-orbital.

Hence, they display a progressive rise in IE with increasing atomic number. Palladium, however, has a filled d-shell and no electron in the s-shell. Therefore, Pd exhibits the greatest IE. Due to lanthanide contraction, the attraction of electrons by the nuclear charge is much higher and thus the IE of 5d elements is much higher than that of 4d and 3d. In the 5d series, all elements except Pt and Au have a filled s-shell.

Elements from Hafnium to Rhenium have the same Ionization Energy (IE) and the IE increases with the number of shared d-electrons, such that Iridium and Gold have the maximum IE.

Metallic Character

D-block elements exhibit typical metallic properties such as high tensile strength, malleability, ductility, electrical and thermal conductivity, metallic luster, and crystallization in bcc/ccp/hcp structures.

The hardness of metals increases with the number of unpaired electrons. Therefore, the d-block elements Cr, Mo, and W are particularly hard, while the group-12 elements Zn, Cd, and Hg are an exception. Copper is an exception to this trend, as it has a high enthalpy of atomization and low volatility.

Oxidation States of D Block Elements

The oxidation state is a hypothetical state, in which an atom appears to have lost or gained electrons relative to its valency state. This concept is still useful in understanding the properties of the atom/ion, particularly in transition elements/ions which have electrons in both s and d-orbitals.

Since the energy difference between s and d-orbitals is small, both electrons can participate in ionic and covalent bond formation, thus allowing for multiple (variable) valency states (oxidation states).

Each transition element can hence exhibit a minimum oxidation state corresponding to the number of s-electrons, and a maximum oxidation state equivalent to the total number of electrons available in both s and d-orbitals. Additionally, oxidation states in between these two extremes are possible.

| Sc | +2,+3 | +3 | +3 | Y | +2,+3 | +3 | +3 | La | +2,+3 | +3 | +3 |

Element	Oxidation States	Ti	+2, +3, +4
	+2	+4
Zr	+2, +3, +4	+2	+4
Hf	+2, +3, +4	+4	+4

| V | +2,+3,+4,+5, | +2 | +5 | | Nb | +2,+3,+4,+5, | +2 | +5 | | Ta | +2,+3,+4,+5, | +4 | +5 |

| Cr | +2,+3,+4,+5,+6, | +1 | +2 | +6 | Mo | +2,+3,+4,+5,+6, | +4 | +6 | W | +2,+3,+4,+5,+6, | +4 | +6 |

| Mn | +2,+3,+4,+5,+6,+7 | +2 | +7 | Tc | +2,+3,+4,+5,+6,+7 | +4 | +7 | Re | +2,+3,+4,+5,+6,+7 | +4 | +7 |

| Fe | +2,+3,+4,+5,+6 | +6 | Ru | +2,+3,+4,+5,+6,+7,+8 | +8 | Os | +2,+3,+4,+5,+6,+7,+8 | +8 |

| Co | +2,+3,+4 | | +2 | +4 | | Rh | +2,+3,+4 | | +3 | +4 | | Ir | +2,+3,+4 | | +4 | +4 |

Element	Charges	+2	+4
Ni	+2,+3,+4	+2	+4
Pd	+2,+3,+4	+2	+4
Pt	+2,+3,+4	+4	+4

| Cu | +1,+2 | +1 | +2 | Ag | +1,+2 | +1 | +2 | Au | +1,+2 | +1 | +2 |

| Zn | +2 | +2 | +2 | Cd | +2 | +2 | +2 | Hg | +2 | +1 | +2 | +2 |

Oxidation States: Trends and Analysis

1. The minimum Oxidation state of 1 is exhibited by Cr, Cu, Ag, Au, and Hg.

2. The oxidation state of 3d series elements is most stable in +2, 4d series in +2 and +4 and 5d series in +4. However, Cr⁶⁺ and Mn⁷⁺ (of 3d series) are not stable in their higher oxidation states. Compounds containing them, such as CrO₄^2- and MnO₄^-, are very reactive and strong oxidizing agents.

While Mo⁶⁺ and Tc⁷⁺ (of 4d) are stable in their higher OS, compounds containing them, MoO₄^2- and **TcO₄^– are unreactive and stable. Similarly, W⁶⁺ and Re⁷⁺ (of 5d) are stable in their higher OS, compounds containing them, WO₄^2- and **ReO₄^– are unreactive and stable.

The first-row transition metals generally form more stable compounds in their +2 and +3 oxidation states than their second and third-row counterparts. For instance, the most stable compounds of chromium are Cr(III), while the corresponding compounds of molybdenum and tungsten, Mo(III) and W(III), are highly reactive.

Pyrophoric elements, such as those found in each group, often burst into flames on contact with atmospheric oxygen. However, the heavier elements in each group form stable compounds in higher oxidation states that are not seen with the lightest member of the group.

3. Strongly oxidizing elements with high oxidation numbers form compounds of oxides and fluorides, but not bromides and iodides.

Vanadium can form only VO₄^-, CrO₄^2-, MnO₄^-, VF₅, VCl₅, VBr₃, VI₃ and not VBr₅, VI₅. V⁵⁺ oxidizes Br^- and I^- to Br₂ and I₂ but not fluoride due to its high electronegativity and small size.

Similarly, strongly reducing, low oxidation number elements form bromides and iodides and not oxides and fluorides.

4. Middle-order elements in each series exhibit a maximum oxidation state equal to the s and d-electrons. For example, Manganese in the 3d series has a maximum oxidation state of +7, while Ruthenium in the 4d series and Osmium in the 5d series possess a maximum oxidation state of +8.

5. Elements can exhibit oxidation states between their minimum and maximum values.

Elements in their lower oxidation states will be ionic and basic (e.g. TiO, VO, CrO, MnO, TiCl2 and VCl2), in-between state amphoteric (e.g. Ti2O3, V2O3, Mn2O3, CrO3, Cr2O3, TiCl3, VCl3) and higher oxidation state covalent and acidic (e.g. V2O5, MnO3, Mn2O7, VCl4 and VOCl3).

7. In complexes such as Ni(CO)4, Fe(CO)5, [Ag(CN)2]–, and [Ag(NH3)2]+, lower oxidation states may be stabilized by back bonding.

Lower oxidation states in these metals are stabilised by ligands such as CO, which are pi-electron donors. On the other hand, higher oxidation states are stabilised by electronegative elements like Fluorine (F) and Oxygen (O). Therefore, the high oxidation compounds of these metals are mainly fluorides and oxides.

Relative stabilities of the oxidation states are affected by multiple factors, such as the stability of the resulting orbital, electronegativity, atomization enthalpy, hydration enthalpy, etc.

Ti4+ (3d0) is more stable than Ti3+ (3d1), and Mn2+ (3d5) is more stable than Mn3+ (3d4).

The relative stabilities of Ni2+ and Pt2+ compounds, as well as of Pt4+ and Ni4+ compounds, can be explained by their respective ionization energies. For instance, Ni2+ compounds are thermodynamically more stable than Pt2+, and Pt4+ compounds are more stable than Ni4+.

Metal	IE1 + IE2 (kJ/mol)	IE3 + IE4 (kJ/mol)	Etotal (kJ/mol)
Metal	IE1 + IE2	IE3 + IE4	Etotal

| Ni | 2490 | 8800 | 11290 |

| Pt | 2660 | 6700 | 9360 |

The ionization of Ni to Ni2+ requires less energy (2490 kJ mol−1) than the energy required for the production of Pt2+ (2660 kJ mol−1). Thus, Ni2+ compounds are thermodynamically more stable than Pt2+ compounds.

The formation of Pt4+ requires less energy (9360 kJ/mol) than that required for the formation of Ni4+ (11290 kJ/mol). This makes Pt4+ compounds more stable than Ni4+ compounds, which is evidenced by the fact that the [PtCl6]2+ complex ion is known, while the corresponding ion for Nickel is not known.

9. The heavier elements in the p-block prefer lower oxidation states due to the inert pair effect, while the heavier members in a group of d-block elements tend to be more stable in higher oxidation states.

Electrode Potential of D Block Elements

The oxidation state of a cation for which ΔH_{+ lE + ΔHhyd} or E° is more negative (for less positive) will be more stable, thus allowing for the prediction of relative stabilities of transition metal ions in different oxidation states in the aqueous medium from electrode potential data.

The E° value becomes less negative along the series, indicating increased stability of the reduced state.

Transition elements have a lower E° compared to first and second group metals.

Physical Properties of D Block Elements

The Trend of Density: As we move down the transition series, the density increases initially and then decreases, which is in reverse of the trend of atomic radii.

Down the column, the density of 4d series is larger than 3d. This is due to lanthanide contraction and a larger decrease in atomic radii, resulting in the volume density of 5d series transition elements being double that of the 4d series.

In the 3d series, scandium has the lowest density while copper has the highest density. Osmium (d=22.57g cm-3) and Iridium (d=22.61g cm-3) of the 5d series have the highest density among all d block elements.

Fe ˂ Ni ˂ Cu ˂ Hg ˂ Au

How do D Block Elements Achieve High Melting and Boiling Points?

The strong covalent bonding formed by unpaired electrons and empty or partially filled d-orbitals, in addition to the metallic bonding by s-electrons, is the reason why d-block elements have higher melting and boiling points than s and p block elements. This trend continues until the d5 configuration, after which it decreases as more electrons become paired in the d-orbital.

Cr, Mo, and W possess the highest melting points in their respective series of elements.

Manganese (Mn) and Technetium (Tc) have half-filled configurations, which leads to weak metallic bonding and abnormally low melting and boiling points.

Group12 (Zn, Cd and Hg) have no unpaired d-electrons, so they are unable to form covalent bonds. As a result, their melting and boiling points will be the lowest in their series.

Mercury – the Liquid Metal: Mercury is the only metal that is liquid at room temperature. Its six valence electrons are more tightly held by the nucleus due to the lanthanide contraction, which results in less involvement of outer s-electrons in metallic bonding.

What Noble Metals are Considered Transition Elements?

In the Three Transition Series

The ionization energies of elements increase gradually across a given row.

From left to right across the 3d to 5d transition elements, density, electronegativity, electrical and thermal conductivities increase, while the enthalpies of hydration of the metal cations decrease in magnitude.

The transition metals become increasingly less reactive and more “noble” in character due to their relatively high ionization energies, increasing electronegativity and decreasing enthalpies of hydration. Metals such as Platinum (Pt) and Gold (Au) in the lower right corner of the d block are so unreactive that they are often referred to as the “noble metals”.

The Magnetic Properties of D Block Elements

Materials are classified by their interaction with the magnetic field as:

• Diamagnetic: If repelled, a diamagnetic material will be pushed away from a magnetic field.

Paramagnetic: If attracted to a magnetic field and can be magnetized temporarily.

Ferromagnetic: If it can maintain its magnetic properties even when not exposed to a magnetic field.

Paired electrons cause diamagnetism. On the other hand, unpaired electrons result in para-magnetism and when these aligned together, unpaired electrons produce ferromagnetism. D block elements and their ions exhibit this behaviour depending on the number of unpaired electrons.

Unpaired electrons contribute to both an orbital magnetic moment and a spin magnetic moment. However, for the 3d series, the orbital angular moment is negligible and the approximate spin-only magnetic moment is given by the formula:

$\mu = \sqrt{4s(s+1)} = \sqrt{n(n+1)} \quad \text{BM}$

The total spin, S, for higher d-series is equal to n, the number of unpaired electrons, multiplied by the Bohr Magneton (BM). In addition to the spin moment, the actual magnetic moment of these elements also includes components from the orbital moment. Chromium and molybdenum are two elements that possess the maximum number (6) of unpaired electrons and magnetic moment.

| Ion | Outer Configuration | Number of Unpaired Electrons | Magnetic Moment (BM) |

| Calculated | Observed |

| Sc3+ | 3d0 | 0 | 0 | 0 |

| Ti3+ | 3d1 | 1 | 1.73 | 1.75 |

| Ti²⁺ | 3d² | 2 | 2.84 | 2.86 |

| V2+ | 3d3 | 3 | 3.87 | 3.86 |

| Cr²⁺ | 3d⁴ | 4 | 4.90 | 4.80 |

| Mn²⁺ | 3d⁵ | 5 | 5.92 | 5.95 |

Element	Configuration	Charge	Ionization Energy (eV)	Electronegativity
Fe2+	3d6	+4	4.90	5.0-5.5

| Co²⁺ | 3d⁷ | 3 | 3.87 | 4.4-5.2 |

| Ni²⁺ | 3d⁸ | 2 | 2.84 | 2.9-3.4 |

| Cu²⁺ | 3d⁹ | 1 | 1.73 | 1.4 - 2.2 |

| Zn²⁺ | 3d¹⁰ | 0 | 0 | 0 |

Formation of Colored Ions by D Block Elements

Compounds of d-block elements have a variety of colours. When a frequency of light is absorbed, the light transmitted exhibits a colour complementary to the frequency absorbed. Transition element ions can absorb the frequency in the visible region to use it two ways and produce a visible colour.

1. d-d Transition

One way is the excitation of an electron to a higher energy level, which is known as d-d transition. In transition element ions, the presence of d-electron and empty d-orbital can result in colour formation. This is due to the excitation and de-excitation of valence electrons.

D-orbitals are generally degenerate, meaning they have the same energy. When ligands that can form coordinate bonds with these ions are present, the degeneracy is removed and the D-orbitals are split into two groups: eg and t2g. The energy difference (∆E) between the two groups depends on the strength of the incoming ligand.

Electrons in the lower d-orbitals can be excited into the higher d-orbitals by absorbing energy in the visible region (λ=400-700nm) and emitting a colour complementary to it.

The [Cu(H2O)6] 2+ ions absorb red radiation and appear blue-green in color, while hydrated Co2+ ions absorb radiation in the blue-green region and appear red when exposed to sunlight.

In the presence of water molecules, cupric ion changes from a colorless state to a blue color.

a) The colour of the ions varies depending on their oxidation state. For example, potassium dichromate is yellow when in the Cr6+ oxidation state, while Cr3+ and Cr2+ are generally green and blue, respectively.

The colour of the compound is dependent on the complexing or coordinating group as well. For example, Cu2+ displays a light blue colour in the presence of water as the ligand, but a deep blue colour in the presence of ammonia as the ligand.

c) Transition metal ions which have:

Cu⁺(3d¹⁰), Zn²⁺(3d¹⁰), Cd²⁺(4d¹⁰), Hg²⁺(5d¹⁰), and Zn, Cd, Hg are colourless due to completely filled d-orbitals with no vacant d-orbitals available for electron excitation.

Transition metal ions which have completely empty d-orbitals without d-electrons are also colourless. For example, Sc³⁺(3d⁰) and Ti⁴⁺(3d⁰) ions are both colorless.

L-M and M-L π-π Bonding

Metal ions may accept the p electrons donated by ligands, forming a ligand-metal or metal-ligand interaction known as dπ – pπ bonding. This type of bonding can also give compounds their colour.

Complex Formation Tendency of D Block Elements

Complex compounds are formed when a metal is bound to a number of neutral molecules or anions. D block elements, due to their small ionic size, high charge, and the availability of d orbitals for bond formation, are especially prone to forming such compounds.

Transition metals and their ions, due to their larger nuclear charge and smaller size, can attract electrons and receive lone pairs of electrons from anions and neutral molecules into their empty d-orbitals, thus forming coordinate bonding.

Examples of transition metal complexes include:

• [Co(NH3) 6] 3+
• [Cu(NH3)4] 2+
• Y(H2O) 6]2+
• [Fe(CN)6]4−
• [FeF6] 3−
• [Ni(CO)4]

Catalytic Activity of Elements

Catalysts are essential for the large-scale industrial production of numerous chemicals. Many d-block elements in their ionic form, such as metals, are commonly used as catalysts in many chemical and biological reactions.

Some very important commercial catalytic processes involving d block metals include Iron in Haber’s process to make ammonia, Vanadium Pentoxide in the manufacture of Sulphuric Acid, Zigler Natta Catalyst in Titanium Chloride polymerization and Palladium Chloride in the conversion of Ethylene to Acetaldehyde.

Most transition elements act as good catalysts due to their ability to easily form complexes with other molecules.

The presence of vacant d-orbitals.

The tendency to show different oxidation states.

The tendency to form reaction intermediates with reactants.

The presence of flaws in their crystal lattices.

By taking the reaction through a path of low activation energy, they are able to achieve their desired outcome:

Providing a large surface area for absorption and allowing enough time for a reaction to occur

May interact with the reactants through their empty orbitals.

May actively participate in redox reactions due to their multiple oxidation states.

Formation of Alloys in D Block Elements

The atomic radii of the transition elements in any series are quite similar, making it easy for them to replace each other in the lattice and form solid solutions. Alloys can be formed when the difference in atomic radii is within 15%.

Such solid solutions are called alloys. Alloys are homogeneous solid solutions of two metals or a metal with a non-metal. The alloys of transition metals are hard and high melting as compared to the host metal.

Iron is combined with other metals such as chromium, vanadium, molybdenum, tungsten, and manganese to create various steels. Some important alloys include:

• Bronze: Cu (75-90%) + Sn (10-25%)
• Chromium Steel: Cr (2-4% of Fe)
• Stainless Steel: Cr (12-14%) and Ni (2-4%) of Fe
• Solder: Pb + Sn

Interstitial Compounds of D Block Elements

Small non-metallic atoms and molecules such as hydrogen, boron, and carbon can be trapped in the voids of transition metal’s crystal lattice structure during crystal formation, forming what are known as interstitial compounds. These compounds are neither ionic nor covalent and are non-stoichiometric, such as TiH1.7 and VH0.56.

Interstitial compounds have the following properties:

Their melting points are extremely high.

They are extremely difficult.

They have comparable conductivity properties to other metals.

They are unreactive and tend to be chemically inert.

Examples of the interstitial compounds that are formed with transition metals are:

• TiC
• Mn4N
• Fe3H
• TiH2

Non-Stoichiometric Compounds

Transition metal compounds of different oxidation states can sometimes be present together, potentially formed by solid structure defects or by the prevailing conditions. However, this mixture behaves as if it were a single compound.

Non-stoichiometric compounds, such as Fe0.94O, Fe0.84O, VSe0.98, and Se1.2, do not have a finite composition or structure, particularly when combining with elements from group 16 (O, S, Se, Te).

Important Compounds of D Block Elements

The D Block Elements form some compounds of vital industrial importance. Some such compounds include:

Potassium Dichromate (K2Cr2O7)

This compound is considered to be of great significance in the leather industry, and is also utilized as an oxidant in the majority of azo compound preparation processes.

The structure of the dichromate ion is composed of two tetrahedra that share a single corner, with a Chromium-Oxygen-Chromium bond angle of 1260. Potassium dichromate is a powerful oxidizing agent, and is also employed as a primary standard in volumetric analysis.

Potassium Permanganate (KMnO4)

KMnO₄ has an intense purple colour. It exhibits both diamagnetic and weak paramagnetic properties, which depend on the temperature.

The permanganate ion exhibits diamagnetism due to the lack of unpaired electrons in it. Additionally, potassium permanganate is used as an oxidant in the preparation of various products in organic chemistry. It is also used for bleaching cotton, silk, and wool, as well as for decolorizing oils due to its powerful oxidizing properties.

D and F Block Elements - Oxidation States

15 Essential Questions on D-Block and F-Block Elements

Alloy Formation of d- and f-block Elements

Colour of Ions in d- and f-Block Elements

Frequently Asked Questions (FAQs)

The general electron configuration of D block elements is [noble gas] ns² np^1-6

The first element of the 3d series is Scandium (Sc) with an electron configuration of [Ar] 3d1 4s2.

D block elements are metals.

One noble metal in D block elements is Gold.

Gold forms interstitial compounds with other D Block elements such as palladium and platinum.

TiH is an interstitial compound of D Block elements.— title: “D Block Elements” link: “/d-block-elements” draft: false

The D block elements can be found from the third to twelfth groups of the modern periodic table. These elements are also referred to as transition elements or transition metals. Their valence electrons fall under the d orbital. The first three rows of the D block elements, which correspond to the 3d, 4d, and 5d orbitals respectively, are discussed in the article below.

Table of Contents

Position of D Block Elements

D Block as Transition Element

Electronic Configuration of D-Block Elements

Atomic and Ionic Radii of D-Block Elements

Properties of D-Block Elements

Oxidation States of D-Block Elements

Formation of Coloured Ions by D-Block Elements

Alloy Formation in D-Block Elements

Important Compounds of D-Block Elements

D Block Elements are groups of elements in the periodic table that have their outermost electron in the d orbital.

Elements which have electrons present in the d-orbital of the penultimate energy level (1 to 10) and in the outermost ‘s’ orbital (1-2) are considered d block elements. Even though the group 12 metals do not completely fill up the ’d’ orbital, their chemistry is similar to the preceding groups and so they are also categorized as d block elements.

Elements in the d-block typically possess metallic qualities such as malleability, ductility, high electrical and thermal conductivity, and good tensile strength. These elements are divided into four series, each series corresponding to the filling of 3d, 4d, 5d, or 6d orbitals.

3d- Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn

4d: Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd

5d- La, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg

Incomplete (6d)

The ’d’ orbital in each series is filled with 10 elements.

Also Read ⇒ This Article

F Block Elements

Actinides

Position of D Block Elements in the Periodic Table

IUPAC defines a transition metal as “an element whose atom or its cations has a partially filled d sub-shell, and which occupies columns 3 to 12 of the periodic table, potentially containing atoms of elements with completely filled ‘d’ orbitals.”

What are D Block Elements and Why are They Called Transition Elements?

Transition elements occupy groups 4–11. Scandium and yttrium of group 3, having a partially filled d subshell in the metallic state, are also considered as transition elements. Elements like Zn, Cd and Hg of the 12 columns of the d block have a completely filled d-orbital and hence are not considered as transition elements.

All the transition elements are d-block elements, but not all d-block elements are transition elements.

Properties of Transition Metals

Electrons added to the ’d’ sub-orbitals that lie between the (n+1)s and (n+1)p sub-orbitals.

Placed between s and p block elements in the periodic table

Properties of s-block and p-block Elements

Electronic Configuration of D Block Elements

The electronic configuration of D block Elements generally follows the pattern of (n-1)d 1-10ns 1-2. These elements often achieve stability through the use of either half-filled or completely filled d orbitals. An example of this is chromium, which has an electronic configuration of 3d54s1, with half-filled d and s orbitals. Copper is another such example, with an electronic configuration of 3d104s1, rather than 3d94s2. More information on electron configurations can be found here.

These metals (Zinc, Mercury, Cadmium, and Copernicium) are not considered transition elements due to their completely filled d orbitals in their ground states and in their general oxidation states. In contrast, the other elements are considered d block elements.

The electronic configuration for period 4, transition elements is (Ar) 4s^1-2 3d^1-10

The electronic configuration for period 5, transition elements is [Kr] 5s[1-2] 4d[1-10]

The electronic configuration for period 6, transition elements is [Xe] 4s^1-2 3d^1-10

Along the period, electrons are added to the 3*d* subshell in accordance with the Aufbau principle and Hund’s rule of multiplicity, from left to right.

| 1st Transition Series | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn |

| 4s23d1 | 4s23d2 | 4s23d3 | 4s13d5 | 4s23d5 | 4s23d6 | 4s23d7 | 4s23d8 | 4s13d10 | 4s23d10 |

| 2nd Transition Series | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd |

| 5S24D1 | 5S24D2 | 5S14D4 | 5S14D5 | 5S24D5 | 5S14D7 | 5S14D8 | 5S04D10 | 5S14D10 | 5S24D10 |

| 3rd Transition Series | La | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg |

| 6s25d1 | 6s25d2 | 6s25d3 | 6s25d4 | 6s25d5 | 6s25d6 | 6s25d7 | 6s15d9 | 6s15d10 | 6s25d10 |

Anomalies do occur in all the series; this can be explained by taking the following considerations into account.

The Energy Gap Between the ns and (n-1)d Orbitals

Pairing energy for the electrons in s-orbital

The stability of half-filled orbitals compared to partly filled orbitals.

Chromium has a 4*s*13*d*5 electron configuration instead of the 4*s*23*d*4 configuration, and copper has a 4*s*13*d*10 configuration instead of the 4*s*23*d*9. These anomalies in the first transition series can be explained by the stability of half-filled orbitals compared to the partly filled orbitals.

In the second series transition metals, starting from niobium, electrons appear to be more likely to occupy the d orbitals rather than the s orbitals. The electron can either opt to share the s orbital or be excited to the d orbital, depending on the repulsive energy it has to overcome and the energy gap between the s and d orbitals.

In the second series, the s and d-orbitals have almost the same energy, which leads to electrons occupying the d-orbital. As a result, the s-orbital in Niobium mostly has only one electron. On the other hand, the third series transition metals have more paired s configurations, even at the expense of half-filled orbitals (e.g. Tungsten- 6s25d4). This series comes after the filling up of 4f orbitals and the resulting lanthanide contraction.

The reduced size of tungsten results in the high shielding of d orbitals by the ‘f’ electron. This shielding increases the energy gap between the s and 5d orbitals, leading to a pairing energy that is less than the excitation. As a result, excitation of the electron does not take place in tungsten, even though half-filled orbitals would make it possible.

Atomic and Ionic Radii of D Block Elements

Metallic Radii of the First, Second, and Third Row Transition Metals

Atomic and Ionic Radii of Elements in the Three Transition Series

Decreases rapidly, from column 3 to 6

Remains steady, from column 7 to 10

Increases from column 11 to 12.

The decrease in atomic radii from Scandium (Sc) to Chromium (Cr) (Group 3 to 6) is quite significant, whereas the decrease from Manganese (Mn), Iron (Fe), Cobalt (Co), and Nickel (Ni) (Group 7, 8, 9, and 10) is almost the same. After this, the atomic radii increase in Copper (Cu) and Zinc (Zn).

The larger decrease in atomic radii, from column 3 to 6 elements, is due to the increase in effective nuclear charge, but poor shielding because of the smaller number of d-electrons.

The increasing effective nuclear charge in elements of column 7 to 10 is counteracted by the repulsion between the shared d electrons, thus maintaining the same radii.

In the case of the elements in groups 11 and 12, the d orbital is filled with 10 electrons, shielding the electrons in the higher s-orbital. As a result, elements such as Cu and Zn in these groups are larger in size than their predecessors in the same block.

Since electrons occupy a higher orbital, the radii of the third series of elements are expected to be greater than the radii of the second series of elements. However, the radii of both series are almost the same. This is because in the third series of elements, the 5d orbitals are filled only after the 4f orbitals, which increases the effective nuclear charge by 14 units.

The larger shrinkage of radii known as Lanthanide contraction is caused by the higher nuclear charge. This increase in radii due to the higher orbital is effectively neutralized by the increase in the nuclear effective charge. Therefore, the atomic radii of second and third series elements are the same, for example, Niobium and Hafnium have almost the same atomic radii.

Properties of D-Block Elements

The Ionization Energy of D Block Elements

Ionization energy is the energy required to remove a valence electron from an atom/ion and is directly related to the force of attraction on the electron. Therefore, the larger the nuclear charge and the smaller the radius of the electron, the higher the ionization energy (IE). Ionization Energy is also higher for half-filled and fully filled orbitals.

The d block elements have an Ionization Energy that is larger than the s-block elements and smaller than the p-block elements, between which they are placed. In the first series, except for chromium and copper, the first Ionization Energy involves removal from a filled s-orbital. Among them, the Ionization Energy of the d block elements increases with the increase in atomic number up to Fe.

The Ionization Energy (IE) of Co and Ni decreases as the number of shared d-electrons increases, compensating for the increase in atomic number. Meanwhile, Copper and Zinc, which are s-block elements, show an increasing IE. In the second series, elements from Niobium onwards have single electrons in their s-orbital.

Hence, they display a steady increase in IE with increasing atomic number. Palladium, however, has a full d-shell and no electrons in the s-shell, leading to the maximum IE for Pd. This is due to lanthanide contraction, where the attraction of electrons to the nuclear charge is much stronger, resulting in higher IE values for the 5d series elements than those of 4d and 3d. All elements in the 5d series, with the exception of Pt and Au, have a filled s-shell.

Elements from Hafnium to Rhenium have the same Ionization Energy (IE) and after IE increases with the number of shared d-electrons such that Iridium and Gold have the maximum IE.

Metallic Character

D-block elements exhibit typical metallic properties such as high tensile strength, malleability, ductility, electrical and thermal conductivity, metallic lustre, and crystallize in bcc/ccp/hcp structures.

The hardness of metals increases with the number of unpaired electrons, making Cr, Mo, and W very hard among d block elements. However, Copper is an exception, as it has high enthalpy of atomization and low volatility. The group-12 elements (Zn, Cd and Hg) also show an exception in this regard.

Oxidation States of D Block Elements

The oxidation state is a hypothetical state which suggests that an atom has released or gained more electrons than its valency state. This is useful in understanding the properties of the atom/ion, especially in the case of transition elements/ions which may have electrons in both s and d-orbitals.

Since the energy difference between s and d-orbitals is small, both electrons can participate in both ionic and covalent bond formation, thus exhibiting multiple (variable) valency states (oxidation states).

Each transition element can hence exhibit a minimum oxidation state corresponding to the number of s-electrons and a maximum oxidation state equivalent to the total number of electrons available in both s and d-orbitals. Additionally, oxidation states in between these two extremes are also possible.

| Sc | +2,+3 | +3 | +3 | Y | +2,+3 | +3 | +3 | La | +2,+3 | +3 | +3 | | Ti | +2,+3,+4, | +2 | +4 | Zr | +2,+3,+4, | +2 | +4 | Hf | +2,+3,+4, | +4 | +4 | | V | +2,+3,+4,+5, | +2 | +5 | Nb | +2,+3,+4,+5, | +2 | +5 | Ta | +2,+3,+4,+5, | +4 | +5 | | Cr | +2,+3,+4,+5,+6, | +1 | +2 | +6 | Mo | +2,+3,+4,+5,+6, | +4 | +6 | W | +2,+3,+4,+5,+6, | +4 | +6 | | Mn | +2,+3,+4,+5,+6,+7 | +2 | +7 | Tc | +2,+3,+4,+5,+6,+7 | +4 | +7 | Re | +2,+3,+4,+5,+6,+7 | +4 | +7 | | Fe | +2,+3,+4,+5,+6, | +2 | +6 | Ru | +2,+3,+4,+5,+6,+7,+8 | +4 | +8 | Os | +2,+3,+4,+5,+6,+7,+8 | +4 | +8 | | Co | +2,+3,+4, | +2 | +4 | Rh | +2,+3,+4, | +3 | +4 | Ir | +2,+3,+4, | +4 | +4 | | Ni | +2,+3,+4, | +2 | +4 | Pd | +2,+3,+4, | +2 | +4 | Pt | +2,+3,+4, | +4 | +4 | | Cu | +1,+2, | +1 | +2 | +2 | Ag | +1,+2, | +1 | +2 | Au | +1,+2, | +1 | +2 | | Zn | +2, | +2 | +2 | Cd | +2, | +2 | +2 | Hg | +2, | +1 | +2 | +2 |

Oxidation States: Trends and Patterns

1. The minimum Oxidation state of Cr, Cu, Ag, Au and Hg is 1.

2. The oxidation state of 3d series elements is most stable at +2, 4d series in +2 and +4, and 5d series in +4. However, Cr6+ and Mn7+ (both of 3d) are not stable in their higher oxidation states. Compounds containing them, such as CrO42- and MnO4–, are very reactive and strong oxidizing agents.

While Mo⁶⁺ and Tc⁷⁺ (of 4d) are stable in their higher OS, compounds containing them, MoO₄^2- and TcO₄^- are unreactive and stable. Similarly, W⁶⁺ and Re⁷⁺ (of 5d) are stable in their higher OS, and compounds containing them, WO₄^2- and ReO₄^- are unreactive and stable.

The cations of the second and third-row transition metals in lower oxidation states (+2 and +3) are much more prone to oxidation than the corresponding ions of the first-row transition metals. For instance, the most stable compounds of chromium are of Cr(III), while the corresponding Mo(III) and W(III) compounds are highly reactive.

Pyrophoric elements, such as those in each group, are prone to bursting into flames upon contact with atmospheric oxygen. However, the heavier elements within each group form stable compounds in higher oxidation states that are distinct from the oxidation states of the lightest element in the group.

3. Strongly oxidizing, high oxidation number elements form compounds of oxides and fluorides, but not bromides and iodides.

Vanadium forms only VO4–, CrO42-, MnO4–, VF5, VCl5, VBr3, VI3 and not VBr5, VI5. V5+ oxidizes Br– and I– to Br2 and I2 but not fluoride because of its high electronegativity and small size.

Similarly, strongly reducing, low oxidation number elements form bromides and iodides, rather than oxides and fluorides.

4. Middle-order elements in each series exhibit a maximum oxidation state equal to the s and d-electrons. For example, Manganese in the 3d series has a maximum oxidation state of +7, while Ruthenium in the 4d series and Osmium in the 5d series possess a maximum oxidation state of +8.

5. Oxidation states of elements may range from the minimum to the maximum.

6. Elements with lower oxidation states will be ionic and basic (e.g. TiO, VO, CrO, MnO, TiCl2 and VCl2), while those with an in-between state are amphoteric (e.g. Ti2O3, V2O3, Mn2O3, CrO3, Cr2O3, TiCl3, and VCl3). Elements with higher oxidation states will be covalent and acidic (e.g. V2O5, MnO3, Mn2O7, VCl4 and VOCl3).

7. The lower oxidation states of Ni(CO)4, Fe(CO)5, [Ag(CN)2]–, and [Ag(NH3)2]+ may be stabilized by back bonding in complexes.

Lower oxidation states of these metals are stabilised by ligands such as CO, which act as pi-electron donors, whereas higher oxidation states are stabilised by electronegative elements like Fluorine (F) and Oxygen (O). Hence, the high oxidation compounds of these metals are primarily fluorides and oxides.

Relative stabilities of the oxidation states are influenced by multiple factors, such as the stability of the resulting orbital, ionization energy, atomization enthalpy, hydration enthalpy, etc.

Ti4+ (3d0) is more stable than Ti3+ (3d1), and Mn2+ (3d5) is more stable than Mn3+ (3d4).

The relative stabilities of Ni2+ and Pt2+ compounds, as well as Pt4+ and Ni4+ compounds, can be explained by the respective ionization energies. For instance, Ni2+ compounds are more thermodynamically stable than Pt2+, while Pt4+ compounds are more stable than Ni4+.

Metal	IE1 + IE2 (kJ/mol)	IE3 + IE4 (kJ/mol)	Etotal (kJ/mol)

| Ni | 2490 | 8800 | 11290 |

| Pt | 2660 | 6700 | 9360 |

The ionization of Ni to Ni2+ requires less energy (2490 kJ mol−1) than the energy required for the production of Pt2+ (2660 kJ mol−1). Therefore, Ni2+ compounds are thermodynamically more stable than Pt2+ compounds.

The formation of Pt4+ requires 9360 kJ/mol, which is lesser than the energy required for the formation of Ni4+ (11290 kJ/mol). Consequently, Pt4+ compounds are more stable than Ni4+ compounds. This is further corroborated by the fact that [PtCl6]2+ complex ion is known, while the corresponding ion for nickel is not known.

9. The heavier elements in the p-block tend to prefer lower oxidation states due to the inert pair effect, while the heavier members of the d block elements are more stable in higher oxidation states.

Electrode Potential of D Block Elements

The oxidation state of a cation for which ΔH(ΔHsub + lE + ΔHhyd) or E° is more negative (for less positive) will be more stable, and the relative stabilities of transition metal ions in different oxidation states in the aqueous medium can be predicted from the electrode potential data.

The E° value becomes less negative along the series, indicating greater stability of the reduced state.

Transition elements have a lower E° compared to first and second group metals.

Physical Properties of D Block Elements

The Trend in Density: The density of the transition series will increase at first, remain almost the same, and then decrease along the period, which is the opposite of the atomic radii.

The column density of the 4d series is larger than that of the 3d series due to lanthanide contraction and a larger decrease in atomic radii, resulting in the volume density of the 5d series transition elements being twice that of the 4d series.

In the 3d series, scandium has the lowest density and copper highest density. Osmium (d=22.57g cm-3) and Iridium (d=22.61g cm-3) of 5d series have the highest density among all d block elements.

Fe < Ni < Cu, Fe < Cu < Au, Fe < Hg < Au.

How Do D Block Elements Have Such High Melting and Boiling Points?

The strong covalent and metallic bonding formed by unpaired electrons and empty or partially filled d-orbitals in d-block elements give them high melting and boiling points compared to s and p block elements. This trend continues up to the d5 configuration and then decreases as more electrons become paired in the d-orbital.

Cr, Mo, and W possess the highest melting points in their respective series of elements.

Manganese (Mn) and Technetium (Tc) possess half-filled configurations, which results in weak metallic bonding and abnormally low melting and boiling points.

Group12, Zn, Cd and Hg have no unpaired d-electrons and therefore do not form covalent bonds. As a result, their melting and boiling points are the lowest in their respective series.

Mercury - the Liquid Metal: Mercury is the only metal that is liquid at room temperature. Its 6s valence electrons are more strongly attracted to the nucleus due to the lanthanide contraction, resulting in less involvement of the outer s-electrons in metallic bonding.

Which Transition Elements are Considered Noble Metals?

In the three transition series

The ionization energies of elements increase gradually along a given row.

From left to right in the 3d to 5d transition series, density, electronegativity, electrical and thermal conductivities all increase, while the enthalpies of hydration of the metal cations decrease in magnitude.

The transition metals become increasingly less reactive and more “noble” as they move down the d-block. This is due to their high ionization energies, increasing electronegativity, and decreasing enthalpies of hydration. Metals such as Platinum (Pt) and Gold (Au) in the lower right corner of the d-block are so unreactive that they are often referred to as the “noble metals”.

Magnetic Properties of D-Block Elements

Materials are classified by their interaction with the magnetic field as:

• Diamagnetic: If repelled by a magnetic field.

Paramagnetic: If a material is attracted to a magnetic field and becomes magnetized in the presence of the field.

Ferromagnetic: If it can maintain its magnetism even when not in a magnetic field.

Paired electrons are responsible for diamagnetism. On the other hand, unpaired electrons lead to para-magnetism, and when aligned together, they cause ferromagnetism. D block elements and their ions may also show this behaviour, depending on the number of unpaired electrons.

Unpaired electrons contribute to both an orbital magnetic moment and a spin magnetic moment. However, for the 3d series, the orbital angular moment is negligible and the approximate spin-only magnetic moment is given by the following formula:

¹√[4s (s + 1)] = ¹√[n (n + 1)] BM

The magnetic moment of an atom is given by S x n x Bohr Magneton (BM), where S is the total spin and n is the number of unpaired electrons. For higher d-series, the actual magnetic moment includes components from the orbital moment in addition to the spin moment. Chromium and molybdenum possess the maximum number (6) of unpaired electrons and magnetic moment.

| Element | Outer Configuration | No. of Unpaired Electrons | Magnetic Moment (BM) |

| Calculated | Observed |

| Sc3+ | 3d0 | 0 | 0 | 0 |

| Ti3+ | 3d1 | 1 | 1.73 | 1.75 |

| Ti²⁺ | 3d² | 2 | 2.84 | 2.86 |

| V2+ | 3d3 | 3 | 3.87 | 3.86 |

| Cr²⁺ | 3d⁴ | 4 | 4.90 | 4.80 |

| Mn²⁺ | 3d⁵ | 5 | 5.92 | 5.95 |

| Fe²⁺ | 3d⁶ | 4 | 4.90 | 5.0-5.5 |

| Co²⁺ | 3d⁷ | 3 | 3.87 | 4.4-5.2 |

| Ni²⁺ | 3d⁸ | 2 | 2.84 | 2.9-3.4 |

| Cu²⁺ | 3d⁹ | 1 | 1.73 | 1.4-2.2 |

| Zn²⁺ | 3d¹⁰ | 0 | 0 | 0 |

Formation of Colored Ions by D Block Elements

Compounds of d block elements have a variety of colours. When a frequency of light is absorbed, the light transmitted displays a colour that is complementary to the frequency absorbed. Transition element ions can absorb frequencies in the visible region and use them to produce visible colours.

1. d-d Transition

The excitation of an electron to a higher energy level, the presence of d-electrons and empty d-orbitals in transition element ions, and the valence electron excitation and de-excitation are all processes that lead to the formation of colour, known as d-d transition.

D-orbitals are typically degenerate and have the same energy. When ligands that can form coordinate bonds with these ions are present, the degeneracy is removed and the orbitals are split into two groups: eg and t2g. The energy difference (∆E) is dependent on the strength of the incoming ligand.

Electrons in the lower d-orbitals can be excited into the higher d-orbitals by absorbing energy in the visible region (λ=400-700nm) and emitting a colour complementary to it.

The [Cu(H2O)6]²⁺ ions absorb red radiation and appear blue-green in color, while hydrated Co²⁺ ions absorb radiation in the blue-green region, and thus appear red in sunlight.

In the presence of water molecules, cupric ion is colorless but turns blue in color.

a) The color of the ions is dependent on their oxidation state. Potassium Dichromate (Cr6+) is yellow in color, while Cr3+ and Cr2+ are usually green and blue, respectively.

b) The colour of a compound is also dependent on the complexing or coordinating group. For example, Cu2+ exhibits a light blue colour in the presence of water as a ligand, but a deep blue colour in the presence of ammonia as the ligand.

c) Transition metal ions which have:

Cu⁺(3d¹⁰), Zn²⁺(3d¹⁰), Cd²⁺(4d¹⁰), and Hg²⁺(5d¹⁰) ions, which have completely filled d-orbitals and no vacant d-orbitals for excitation of electrons, are colourless. Zn, Cd, and Hg are also colourless.

Transition metal ions which have completely empty d-orbitals without d-electrons are also colorless. For example, Sc³⁺(3d⁰) and Ti⁴⁺(3d⁰) ions are colorless.

L-M-π and M-L-π Bonding

The interaction between ligands and metal ions, known as ligand-metal or metal-ligand (dπ–pπ) bonding, involves the donation of p electrons from the ligands into the empty d orbitals of the metal ions, which can give the compounds color.

The Tendency of D Block Elements to Form Complexes

Complex compounds are compounds composed of a metal bound to multiple neutral molecules or anions. The d block elements in the periodic table are especially prone to forming complex compounds due to their small ionic size, high charge, and the presence of d orbitals that facilitate chemical bonding.

Transition metals and their ions, due to their larger nuclear charge and smaller size, can attract electrons and receive lone pairs of electrons from anions and neutral molecules into their empty d-orbitals, thus forming coordinate bonding.

Examples of transition metal complexes include:

• [Co(NH3)6]3+
• [Cu(NH3)4]2+
• [Y(H2O)6]2+
• [Fe(CN)6]4−
• [FeF6]3−
• [Ni(CO)4]

Catalytic Activity of Elements

Catalysts are essential for the industrial-scale production of many chemicals. Many d-block elements, in their ionic form, are used as catalysts in numerous chemical and biological reactions.

Some very important commercial catalytic processes involving d block metals include:

• Iron in Haber’s process to make ammonia
• Vanadium pentoxide in the manufacture of sulphuric acid
• Titanium chloride as Zigler Natta catalyst in polymerization
• Palladium chloride in the conversion of ethylene to acetaldehyde

Most transition elements act as good catalysts due to their ability to form multiple oxidation states.

The presence of vacant d-orbitals.

The tendency to exhibit different oxidation states.

The tendency to form reaction intermediates with reactants.

The presence of flaws in their crystal lattices.

They take the reaction through a path of low activation energy by utilizing catalysts:

Giving a wide area for absorption and allowing enough time to respond

May interact with the reactants through their vacant orbitals.

May actively interact through their multiple oxidation states via redox reaction.

Formation of Alloys in D Block Elements

The atomic radii of the transition elements in any series are relatively similar, making it easy for them to replace each other in the lattice and form solid solutions. Atoms with a difference in radii of up to 15% can form alloys.

Alloys are homogeneous solid solutions of two metals or a metal with a non-metal. These solid solutions are known as alloys and they tend to be harder and have higher melting points than the host metal. Transition metals are particularly prone to forming alloys.

Iron is often alloyed with other metals, such as chromium, vanadium, molybdenum, tungsten, and manganese, to form various steels. Some important alloys include:

Bronze: Cu (75-90%) + Sn (10-25%)
Chromium Steel: Cr (2-4% of Fe)
Stainless Steel: Cr (12-14%) and Ni (2-4%) of Fe
Solder: Pb + Sn

Interstitial Compounds of D Block Elements

Voids in the crystal lattice structure of transition metals can be filled with small non-metallic atoms and molecules such as hydrogen, boron, and carbon. These compounds, known as interstitial compounds, are neither ionic nor covalent and are non-stoichiometric, as seen in compounds such as TiH1.7 and VH0.56.

Interstitial compounds have the following properties:

Their melting points are extremely high.

They are extremely difficult.

They have similar conductivity properties when compared to other metals.

They are unreactive and tend to be chemically inert.

Examples of the interstitial compounds that are formed with transition metals are:

• TiC
• Mn4N
• Fe3H
• TiH2

Non-Stoichiometric Compounds

Transition metal compounds of different oxidation states may sometimes coexist. They can be formed due to solid structure defects or due to the prevailing conditions. However, this mixture behaves as if it is a single compound.

Non-stoichiometric Compounds with Group 16 Elements

Examples:

• Fe0.94O
• Fe0.84O
• VSe0.98
• Se1.2

Important Compounds of D Block Elements

The D Block Elements form some compounds of vital industrial importance. Some such compounds include:

Potassium Dichromate (K2Cr2O7)

This compound is of great significance in the leather industry and is also used as an oxidant in the preparation of most azo compounds.

The structure of the dichromate ion consists of two tetrahedra that share a single corner with a Chromium-Oxygen-Chromium bond angle of 1260. Potassium dichromate is a powerful oxidizing agent and is also used as a primary standard in the process of volumetric analysis.

Potassium Permanganate (KMnO4)

The physical appearance of KMnO4 has an intense purple colour. It exhibits both diamagnetic and weak paramagnetic properties, which are dependent on the temperature.

The permanganate ion exhibits diamagnetism because it does not have any unpaired electrons. Potassium permanganate is a useful oxidant in organic chemistry and is used in bleaching cotton, silk, and wool. It is also used to decolourize oils due to its strong oxidizing ability.

D and F Block Elements – Oxidation States

15 Essential Questions Regarding D-Block and F-Block Elements

Alloy Formation of d-block and f-block Elements

![Alloy Formation of d-f Block Elements]()

Colour of Ions in d- and f-Block Elements

Frequently Asked Questions (FAQs)

The general electron configuration of D block elements is [noble gas] nd^1-10 ns^2-8 np^1-6.

The first element of the 3d series is Scandium (Sc), with an electron configuration of [Ar] 3d1 4s2.

D block elements are metals.

One noble metal in D block elements is gold.

Gold is a noble metal in the D block elements. Examples of interstitial compounds of D Block elements include Iron Carbide (FeC) and Iron Nitride (FeN).

TiH is an interstitial compound of D Block elements.