Electronic Configuration of \({\rm{d}}\) Block Elements: Transition elements are those elements that belong to the \({\rm{d – }}\)block of the periodic table. They are called transition elements because they show the gradual transition between \({\rm{s – }}\)block and \({\rm{p – }}\)block elements.

We come across many transition elements daily without actually realizing it. Many items ranging from jewellery to kitchen cutleries are all made up of transition elements. From ships to buildings, from utensils to food colours, all consist of \({\rm{d – }}\)block elements. The most abundantly found transition elements are iron and titanium. Scroll down to know more about the general electronic configuration of d block elements.

What are d-block Elements and their Electronic Configuration?

The elements that have or can readily form partially filled \({\rm{‘d’}}\) orbitals are called \({\rm{d – }}\)block elements or transition elements.

The \({\rm{d – }}\)block elements are present in the group \(3\) to \(12\) of the modern Periodic table. The \({\rm{‘d’}}\) block elements lie in the middle of the Periodic table, to the right of \({\rm{s}}\) block elements and to the left of \({\rm{p}}\) block elements. The \({\rm{d – }}\) block elements represent a change or transition in properties from the most electropositive \({\rm{s – }}\) block element to the most electronegative \({\rm{p – }}\) block elements. Transition elements usually have partly filled \(\left( {{\rm{n – 1}}} \right)\,{\rm{d – }}\)orbitals.

In total, \(40\) elements in the Periodic Table belong to the d-block.

Classification of d-block Elements

Based on whether the last electron goes to \(3\;{\rm{d}},4\;{\rm{d}},5\;{\rm{d}},\) or \(6\;{\rm{d}}\) orbitals, the d-block elements are classified into four series. These are-

3d Series or First transition series

This Series of elements involve the filling of \({\rm{3d}}\) orbitals. It starts with Scandium\(\left( {{\rm{Sc}}} \right)\) with atomic number \(21\) and goes up to Zinc \(\left( {{\rm{Zn}}} \right)\) with an atomic number of \(30.\)

4d Series or Second transition series

This Series of elements involves filling \({\rm{4d}}\) orbitals and includes \(10\) elements from yttrium\(\left( {\rm{Y}} \right)\) with atomic number \(39\) to Cadmium\(\left( {{\rm{Cd}}} \right)\) with an atomic number of \(48.\)

5d Series or Third transition series

This Series of elements involve the filling of \({\rm{5d}}\) orbitals. The first element of this Series is Lanthanum \(\left( {{\rm{La}}} \right),\) with an atomic number of \(57.\) It is followed by \(14\) elements called Lanthanides which involve the filling of \(4\;{\rm{f}}\) orbitals. The following \(9\) elements, from Hafnium \(\left( {{\rm{Hf}}} \right)\) with atomic number \(72\) to Mercury \(\left( {{\rm{Hg}}} \right)\) with atomic number \(80,\) belong to the third Series of Transition elements.

6d Series or Fourth transition series

This Series is incomplete because some of the elements of this Series are unknown. It ranges from Actinium \(\left( {{\rm{Ac}}} \right)\) with atomic number \(89,\) Rutherfordium \(\left( {{\rm{Rf}}} \right)\) with atomic number \(104,\) to an element with atomic number \(112.\)

General Electronic Configuration of d-block Elements

The general electronic configuration of d-block elements represents the arrangement of electrons in orbitals around the nucleus of an atom.

Orbitals \({\rm{s,p,d,}}\) and \({\rm{f}}\) are the four chief nuclear orbitals. The electrons fill these orbitals according to their energy levels. These four orbitals can be arranged in the increasing order of their energy levels as \({\rm{s < p < d < f,}}{\rm{.}}\)

The filling of electrons in \({\rm{s,p,d,}}\) and \({\rm{f}}\) orbitals occur according to Aufbau’s principle, which states that the orbital possessing the least energy is filled first.

The \({\rm{s}}\) orbital can get a maximum of two electrons while \({\rm{p,d,}}\) and \({\rm{f}}\) orbitals can hold \(6, 10,\) and \(14\) electrons, respectively. 

In the transition elements, the last electron usually enters the penultimate \({\rm{(n – 1)d}}\) orbitals, and that is why they are called \({\rm{d – }}\)block elements in the modern periodic table. The general valence shell configurations of every transition element are \({\rm{(n – 1)}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{n}}{{\rm{s}}^{{\rm{0,1,2}}}}\)

The \({\rm{(n – 1)}}\) indicates the inward d orbitals with one to ten electrons, and the peripheral ns orbital may have one or two electrons.

Based on the Periodic Table, the variable \({\rm{‘n’}}\) can be determined simply by locating the row number or period to which the particular element belongs. 

Moreover, half and filled orbitals are moderately more stable.

1^st Series or 3d Series of d-block Elements

The external configuration of first-line transition elements is \({\rm{[Ar]3}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{4}}{{\rm{s}}^{{\rm{1 – 2}}}}\) This Series of elements involve the filling of \({\rm{3d}}\) orbitals. It starts with Scandium\(\left( {{\rm{Sc}}} \right)\) with atomic number \(21\) and goes up to Zinc \(\left( {{\rm{Zn}}} \right),\) having an atomic number of \(30.\) The electronic configuration of different elements of the \({\rm{3d}}\) Series are shown below-

From the table above, we can figure out that chromium and copper don’t follow \({\rm{[Ar]3}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{4}}{{\rm{s}}^{{\rm{1 – 2}}}}\)

Chromium \(\left( {{\rm{Cr}}} \right)\) has an electronic configuration that is \({\rm{[Ar]3}}{{\rm{d}}^{\rm{5}}}{\rm{4}}{{\rm{s}}^{\rm{1}}}\) instead of \({\rm{[Ar]3}}{{\rm{d}}^{\rm{4}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}\).

Copper \(\left( {{\rm{Cu}}} \right)\) has an electronic configuration that is \({\rm{[Ar]3}}{{\rm{d}}^{\rm{10}}}{\rm{4}}{{\rm{s}}^{\rm{1}}}\) instead of \({\rm{[Ar]3}}{{\rm{d}}^{\rm{9}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}\) .

The reason behind this anomaly is because of the following reasons:

1. The energy gap between the two sets \({\rm{3d}}\) and \({\rm{4s}}\) of orbitals is sufficiently little to anticipate electrons entering the \({\rm{3d}}\) orbitals. It is tentatively found that half and filled orbitals are more stable.
2. Pairing energy of the electrons in the \({\rm{s}}\) orbital.
3. Stability of half-filled orbitals to the partly filled orbitals.

2^nd Series or 4d Series of d-block Elements

This Series of elements involves filling \({\rm{4d}}\) orbitals and includes \(10\) elements from yttrium\(\left( {\rm{Y}} \right)\) with atomic number \(39\) to Cadmium\(\left( {{\rm{Cd}}} \right)\) with an atomic number of \(48.\) The external configuration of the second-line transition elements as \({\rm{[Kr]4}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{5}}{{\rm{s}}^{{\rm{1 – 2}}}}\).The electronic configuration of the \({\rm{d – }}\)block elements in the second series is as follows:

3^rd Series or 5d Series of d-block Elements

This Series of elements involve the filling of \({\rm{3d}}\) orbitals. The first element of this series is Lanthanum \(\left( {{\rm{La}}} \right){\rm{,}}\) having an atomic number of \(57.\) It is followed by \(14\) elements called Lanthanides which involve the filling of \(\left( {{\rm{4f}}} \right)\) orbitals. The following \(9\) elements from Hafnium \(\left( {{\rm{Hf}}} \right)\) with atomic number \(72\) to Mercury \(\left( {{\rm{Hg}}} \right)\) with atomic number \(80\) belong to the third series transition elements. The external configuration of the second-line transition elements as \({\rm{[Xe]5}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{6}}{{\rm{s}}^{\rm{2}}}\). The electronic configuration of the \({\rm{d – }}\)block elements in the third Series is as follows:

4^th Series or 6d Series of d-block Elements

This series is incomplete because some of the elements of this Series are unknown. It includes Actinium \(\left( {{\rm{Ac}}} \right)\) with atomic number \(89,\) Rutherfordium \(\left( {{\rm{Rf}}} \right)\) with atomic number \(104,\) to an element with atomic number \(112.\) The external configuration of the second-line transition elements as \({\rm{[Rn]6}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}\).The electronic Configuration of the \({\rm{d – }}\)block elements in the fourth Series is as follows:

General Outer Electronic Configuration of d-block Elements

The general valence shell configurations of every transition element is :

\({\rm{(n – 1)}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{n}}{{\rm{s}}^{{\rm{0,1,2}}}}{\rm{.}}\)

The external configuration of first-line transition elements is \({\rm{[Ar]3}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{4}}{{\rm{s}}^{{\rm{1 – 2}}}}{\rm{.}}\) This Series of elements involve the filling of \({\rm{3d}}\) orbitals.

The external configuration of the second-line transition elements is \({\rm{[Kr]4}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{5}}{{\rm{s}}^{{\rm{1 – 2}}}}{\rm{.}}\)

The external configuration of the third-line transition elements is \({\rm{[Xe]5}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{6}}{{\rm{s}}^{\rm{2}}}{\rm{.}}\)

The external configuration of the fourth-line transition elements is \({\rm{[Rn]6}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}\)

Practice Exam Questions

Exceptional Electronic Configuration of d-block Elements

Periodicity in Properties of First Transition Series Elements

1. Atomic Radius/Ionic Radius: It is the distance between the nucleus of an atom/ion and the electron in the outer shell. The value of the cationic radius is always less than the atomic radius. The atomic radius decreases from the third to the seventh square on moving from left to right in the first transition series of the d block, after which the atomic radius remains almost the same till the 10th square. After this, there is a slight increase in the atomic radius in the 11th and 12th classes.

2. Ionisation Enthalpy/Ionisation Energy: The energy required to separate an electron present in the outermost shell of an isolated gaseous atom is called ionisation energy. Moving from left to right in the 3d series, decreasing the atomic radius till the seventh group requires more energy to remove the electrons present in their outermost shell, hence the ionization enthalpy initially increases in the 3d-series. But after the 7th class, due to a small increase in atomic radius, the value of ionization enthalpy also decreases.

3. Metallic Nature: The d block metals are present in the following crystalline arrangements:
(i) BCC
(ii) FCC
(iii) HCP
Due to the excess of effective nuclear charge and the presence of unpaired d electrons, metallic bonds are strong in them, so due to strong intermolecular metallic bonding, their melting point, boiling point and particle enthalpy values ​​are high.

As the number of unpaired d electrons increases in the 3d series, the values ​​of melting point, boiling point, hardness and particle enthalpy of these metals increase.

4. Variable Oxidation State: The d-block metals show more than one type of oxidation state, hence their oxidation state is called the variable oxidation state. Since the energy of the final (n-1)d and nS subshells in d-block metals is almost the same, the electrons of these two subshells participate in bonding, so it shows a variable oxidation state. Their highest oxidation state is equal to the sum of the total electrons present in nS and the number of unpaired electrons present in (n-1)d orbitals.

5. Catalytic Properties: Catalytic properties are found in d block metals, due to the following reasons:
(i) In these, a vacant d orbital becomes available.
(ii) It shows a variable oxidation state.
(iii) They tend to form complex compounds.
(iv) Their size is small.

All Transition Elements are d-block Elements, but all d-block Elements are not Transition Elements

\({\rm{Zn,}}\,{\rm{Cd,}}\) and Hg have their orbitals filled both in their ground state and common oxidation states. Therefore, we can represent it as \({\rm{(n – 1)}}{{\rm{d}}^{{\rm{10}}}}{\rm{n}}{{\rm{s}}^{\rm{2}}}\). Although they are placed in \({\rm{d – }}\)block due to \({\rm{d}}\) orbital, they are not considered transition elements. This is because these elements have filled d orbitals which imparts extra stability to the atom. Hence, they are not referred to as transition elements.

General Properties of d-block Elements

All transition elements exhibit similar properties because of the identical electronic configuration of their peripheral or outermost shell. On every addition of an extra electron to the \({\rm{3d}}\) shell, an effective shield is created between the nucleus and the outer \({\rm{4s}}\) shell. The peripheral shell configuration of these elements is \({\rm{n}}{{\rm{s}}^{\rm{2}}}\). The general properties of the transition elements are as follows:

1. They form stable complexes.
2. They have high melting and boiling points.
3. Contain large charge/radius ratio.
4. They form compounds which are often paramagnetic.
5. They are hard and possess high densities.
6. They form compounds with profound catalytic activity.
7. They show variable oxidation states.
8. They form coloured ions and compounds.

Metallic Nature

As there are only one or two electrons in the peripheral shell, all the transition elements exhibit properties of metals. They are highly malleable, ductile, and excellent conductors of electricity and heat. In addition, they possess metallic lustre and tensile strength. Apart from Mercury, which is fluid and delicate like alkali metals, all the transition elements are hard and brittle. Hardness increases with the number of unpaired electrons.

Melting and Boiling Points

Transition elements have high melting and boiling points due to overlapping \({\rm{(n – 1)d}}\) orbitals and the covalent bonding of the unpaired d orbital electrons. However, \({\rm{Zn,}}\,{\rm{Cd,}}\) and \({\rm{Hg}}\) have filled \({\rm{(n – 1)d}}\) orbitals, due to which they cannot form covalent bonds. Thus, they have a lower melting point than other \({\rm{d – }}\) block elements. Manganese \(\left( {{\rm{Mn}}} \right)\) and Technetium \(\left( {{\rm{Tc}}} \right)\) have half-filled configurations resulting in weak metallic bonding and abnormally low melting and boiling points.

Ionic Radii

1. Ionic radii decrease rapidly from column \(3\) to \(6\) due to an increase in the effective nuclear charge and poor shielding of the smaller number of \({\rm{d – }}\)electrons.
2. Ionic radii remain steady from columns \(7\) to \(10\) because the effective nuclear charge is balanced by the repulsion between the shared d electrons.
3. Ionic radii start increasing from column \(11\) to column \(12;\) the \({\rm{d}}\) orbital is full with ten electrons and shields the electrons present in the higher \({\rm{s – }}\)orbital. So, group \(11\) &\(12\) elements like \({\rm{Cu}}\) & \({\rm{Zn}}\) have bigger sizes than their earlier elements in the block.

Density

The density of the transition elements is relatively higher than those of the s block elements. However, the densities gradually decrease from scandium to copper because of an irregular decrease in metallic radii and a relative increase in atomic mass. 

Ionisation Energy

Ionisation Energy is the energy required to remove a valence electron from an atom or ion. The higher the nuclear charge and the smaller the radii, the higher is the ionization energy. Transition elements are less electropositive than \({\rm{s – }}\)block elements. Hence they do not form ionic compounds. The ionization potential of transition elements lies between \({\rm{s}}\) and \({\rm{p}}\) block elements. They possess high ionization energy because of their small size.The ionization potential of \({\rm{d – }}\)block elements increases from left to right. The ionization energy of \({\rm{Cr}}\) and \({\rm{Cu}}\) is higher than their neighbours because of the half-filled and filled d-orbitals.

Oxidation State

All the transition elements, apart from the first and the last, display various oxidation states. 

As the energy difference between s and d-orbital is small, both electrons can be involved in ionic and covalent bond formation and exhibit multiple(variable) valency states (oxidation states).

The elements scandium through manganese (the first half of the first transition series) show the highest oxidation state. This is because their valence shell exhibits loss of all of the electrons in both the s and d orbitals. Manganese displays the highest number of oxidation states which are \(+2, +3, +4, +5, +6,\) and \(+7.\)

Iron forms oxidation states from \(+2\) to \(+6.\) Elements in the first transition series form ions with a charge of \(+2\) or \(+3.\) The elements belonging to the second and third transition series generally are more stable in higher oxidation states than the first series elements. In general, as the atomic radius increases down a group, ions of the second and third Series become larger than the ions in the first Series.

Trick to Remember Electronic Configuration of the s, p, d, f Elements

Here we have given an easy trick to remember the general electronic configuration of s, p, d and f block elements.

? SS PS PS DPS DPS FDPS FDPS

School(S), School(S), Public(P), School(S), Public(P), School(S), Divisional(D), Public(P), School(S), Divisional(D), Public(P), School(S), Federal(F), Divisional(D), Public(P), School(S), Federal(F), Divisional(D), Public(P), Federal(F), Divisional(D)

Common Uses of d-block Elements

1. The most abundant transition element, iron, plays a vital role in building infrastructures.
2. Zinc is used as a galvanizing agent and as an anode in dry batteries.
3. Tungsten is used in making electrical fibres.
4. Iron is used as a catalyst in the manufacture of ammonia.
5. Silver bromide is used in photography.
6. Titanium is used to detect underwater sound.
7. Mercury is used in thermometers because of its high thermal coefficient of expansion.
8. Iron is present in our body as haemoglobin, which transports oxygen to the brain and muscles.

Summary

Transition elements are key to life and evolution. In this article, we studied \({\rm{d – }}\) block elements and their general electronic configuration. We also discussed the various properties and uses of these elements and their real-life applications and importance.

FAQs on Electronic Configuration of d-block Elements

Q.1: Why are \({\rm{d}}\) block elements coloured?
Ans: Most of the compounds of transition elements are coloured due to incompletely filled \({\rm{(n – 1)d}}\) orbitals. These unpaired \({\rm{d – }}\)electrons experience electronic transition starting with one d-orbital then onto the next. During this \({\rm{d – d}}\) transition, the electrons absorb radiation of certain energy and transmit the rest of the energy as coloured light.

Q.2: What is the general electronics configuration of \({\rm{D}}\) block elements?
Ans: The general electronic configuration of d-block elements is \({\rm{(n – 1)}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{n}}{{\rm{s}}^{{\rm{0,1,2}}}}\).

Q.3: \({\rm{Z}}{{\rm{n}}^{{\rm{ + 2}}}}\) salts are white while \({\rm{C}}{{\rm{u}}^{{\rm{ + 2}}}}\) salts are coloured. Why?
Ans: \({\rm{Z}}{{\rm{n}}^{{\rm{ + 2}}}}\) salts are white because it does not have an unpaired electron, whereas \({\rm{C}}{{\rm{u}}^{{\rm{ + 2}}}}\) salts are coloured because they have an unpaired electron and undergo \({\rm{d – d}}\) transition by absorbing light from the visible region and radiating blue colour.

Q.4: What is the general electronic configuration of \({\rm{D}}\) and \({\rm{F}}\) block elements?
Ans: The general electronic configuration of d-block elements is \({\rm{(n – 1)}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{n}}{{\rm{s}}^{{\rm{0,1,2}}}}\) and the general electronic configuration of f-block elements is \({\rm{(n – 2)}}{{\rm{f}}^{{\rm{1 – 14}}}}{\rm{(n – 1)}}{{\rm{d}}^{{\rm{0 – 1}}}}{\rm{n}}{{\rm{s}}^{\rm{2}}}\).

Q.5: Why is Zinc not regarded as a transition element?
Ans: Zinc is not regarded as a transition element because neither \({\rm{Zn}}\) nor \({\rm{Z}}{{\rm{n}}^{{\rm{ + 2}}}}\) ions have incompletely filled d-orbitals.

We hope this detailed article on the electronic configuration of d-block elements helps you in your preparation. If you get stuck do let us know in the comments section below. We will get back to you at the earliest.

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