The First Transition Series

Introduction

d-Block elements (transition elements):
- Lie between s-block and p-block elements.
- Include elements of Groups 3 to 12.
- Start/Occur in the fourth and subsequent periods.
- All contain 1 or 2 electrons in their s-orbital, and the last electron enters the last but one d-subshell i.e. $(n-1)d$ .
- In at least one of their oxidation states, an incomplete d sub-shell (i.e., 1 – 9 electrons) is present.
- Represent a gradual transition/change from most electropositive s-block elements to least electropositive p-block elements.
First Transition series: Scandium (Z = 21) to Zinc (Z = 30), 4th Period; 3d orbitals are filled.
Groups 3-10 are named after their first members, i.e., Group 3 is the Scandium Group.

General Features and Electronic Configuration of Atoms

General Electronic Configuration for all d-block elements is $(n-1)d^{1-10} ns^{1-2}$ .
In Hydrogen atom, only one electron is present, so all the s, p, d, f orbitals of the same shell have the same energy.
In the case of multi-electron atoms, the s, p, d, f orbitals in the same shell have different energies.
This difference arises due to the shielding effect of the inner electrons of the atom.
This shielding varies depending on the principal shell and the subshell in which the electron is entering.
Due to shielding, the actual nuclear charge felt by the electron decreases.
Each electron is imperfectly shielded by inner electrons from the nuclear charge.
The energy of an electron in an atom due to shielding is given as:
$E = -\frac{2 \pi^2 m e^4 (Z^*)^2}{n^2 h^2}$
where $Z^*$ is the effective nuclear charge.
The energy of the electron decreases with an increase in the value of $Z^*$ with an increase in atomic number.
The various subshells (s, p, d, f) in the same principal shell are shielded to different extents due to their different shapes.
The order of the shielding effect of electrons in different orbitals is $s \gt p \gt d \gt f$ .
From Hydrogen (Atomic No= 1) to Argon (Atomic No= 18), the electronic configuration is filled accordingly i.e., 1s, 2s, 2p, 3s, and 3p.
The next electron having (Atomic No= 19) should enter the 3d subshell, but the electron enters the 4s subshell instead of 3d.
This anomalous behavior is explained in the electronic configuration of transition elements and is seen in all the transition series.
Due to the different penetration power of the subshells in a given principal shell ( $s \gt p \gt d \gt f$ ), the orbital which penetrates more to the nucleus:
- will be less shielded from other electrons.
- will feel more nuclear charge.
- will be strongly attracted by the nucleus.
- will have lower energy.
Radial Probability Distribution Curve for 3d and 4s orbitals:
- For 3d:
  - Consists of a single peak.
  - The radius of maximum probability is closer to the nucleus.
  - Less penetration of the Inner core (Argon Core).
- For 4s:
  - Consists of four peaks.
  - The radius of maximum probability is much farther from the nucleus.
  - More penetration of the Inner core (Argon Core).
The penetration power of 4s orbitals is quite great as compared to the 3d orbital.
So, 4s orbitals are more attracted by the nucleus and are more strongly held than 3d orbitals.
Thus, 4s electrons have lesser energy than 3d electrons and are filled first.
As soon as the 4s orbital is filled, the 3d electrons penetrate the electron density of the 4s orbital.
The effective nuclear charge on 3d electrons increases and their energy drops even lower than the energy of the 4p electron.
Before filling electrons, the energy of the 4s sub-shell is lower than that of the 3d sub-shell.
- $\implies$ 4s sub-shell is filled before 3d sub-shell.
Once the 4s sub-shell is filled, its energy will increase.
- $\implies$ The lowest energy sub-shell becomes the 3d sub-shell, so the next electron is put into the 3d sub-shell.
Thus, the elements from Scandium (Z=21) to Zinc (Z=30) involve the filling of the 3d subshell and are called the First Transition Series.

Anomalous Configuration of Chromium & Copper

Why do Chromium and Copper have anomalous Configuration?
- Half-filled and fully filled d orbitals are more stable.
- The energy difference between 4s & 3d orbitals is less.
- Electrons from the filled s orbital shift to the d orbital, resulting in half and fully filled d orbitals.
- Electronic Configuration of Cr is 4s1, 3d5 instead of 4s2, 3d4.
- Electronic Configuration of Cu is 4s1, 3d10 instead of 4s2, 3d9.

Electronic Configuration of Ions

Ions are formed by the loss of electrons from the outermost orbital of atoms.
Transition Metals form ions by loss of electrons first from the 4s subshell and then from the 3d subshell.
Once electrons start entering the 3d subshell, its energy becomes less than that of the 4s subshell.
3d orbitals become inner orbitals and are more stable as compared to 4d.
Strictly speaking, Scandium (Sc) and Zinc (Zn) are not transition elements because:
- Sc forms Sc3+ ion, which has an empty d sub-shell (3d0).
- Zn forms Zn2+ ion, which has a completely filled d sub-shell (3d10).
Cu+ is not a transition metal ion as it has a completely filled d sub-shell.
Cu2+ is a transition metal ion as it has an incompletely filled d sub-shell.
Cu shows some intermediate behavior between transition and non-transition elements because of two oxidation states, Cu(I) & Cu(II).

General Features of the d-Block Elements from Sc to Zn

d-Block Elements as Metals:
- d-block elements are typical metals.
- Are hard (except Group 11 elements).
- Electropositive in nature, High heats of vaporization.
- Good conductors of heat and electricity, strong, malleable, ductile, and lustrous.
- High melting and boiling points, except Hg is a liquid at room temperature.
- Form colored compounds.
- Tend to form a large number of complexes.
- Show many oxidation states.
- Paramagnetic in nature generally.
- Form alloys with other metals (Transition elements have similar atomic radii, which make it possible for the atom of one element to replace those of another element in the formation of an alloy).
  - e.g., Mn is used for giving hardness and wearing resistance to its alloy (duralumin).
  - Cr is used for giving inertness to stainless steel.
- Show good catalytic behavior.

Atomic Radii and Ionic Radii

Observations:
- d-block metals have smaller atomic radii than s-block metals.
- The atomic radii of the d-block metals do not show much variation across the series.
- The atomic radii decrease initially, remain almost constant in the middle, and then increase at the end of the series.
The atomic size reduces at the beginning of the series because:
- Increase in effective nuclear charge with atomic numbers as electrons are being added in the $(n-1) d$ subshell.
  - $\implies$ the electron clouds are pulled closer to the nucleus, causing a reduction in atomic size.
The atomic size decreases slowly in the middle of the series because:
- When more and more electrons enter the inner 3d sub-shell.
  - $\implies$ the screening and repulsive effects of the electrons in the 3d sub-shell increase.
  - $\implies$ the effective nuclear charge increases slowly.
The atomic size increases at the end of the series because:
- The screening and repulsive effects of the 3d electrons reach a maximum.
The reasons for the trend of the ionic radii of the d-block elements are similar to those for the atomic radii.
Ionic radii generally decrease with an increase in nuclear charge.

Density

d-block metals are generally denser than the s-block metals because filling the inner 3d orbital causes an increase in nuclear charge.
The atomic volume of transition metals decreases.
$Density = \frac{Mass}{Volume}$
So, density increases as volume decreases.
This is in agreement with the general decrease in atomic radius across the series.
Therefore, transition metals have very high densities.
Density reaches a maximum value till Ni of Group 10; after that, it decreases.
Cu and Zn have fully filled 3d orbitals, which result in increased forces of repulsion between the added electrons.
These forces of repulsion are more than the attractive force due to increased nuclear charge.
The atomic volume of Cu and Zn increases, resulting in a decrease in their density.

Melting Point and Boiling Point

The melting and boiling points increase, reach a maximum, and then decrease.
Reasons:
1. d-block metal atoms are small in size and closely packed in the metallic lattice.
2. The atoms in transition elements have strong metallic bonding.
3. The strength of metallic bonding depends on the interaction of the unpaired electrons in the outermost orbital.
d-block elements have very high melting and boiling points.
From Sc to Cr, the number of unpaired electrons increases, resulting in an increase in metallic strength.
From Mn to Zn, there is a continuous decrease in the number of unpaired electrons, resulting in a decrease in metallic strength.
The melting point increases till Cr, and then it decreases.
EXCEPTIONAL CASES OF Mn & Zn:
- Mn has quite a low value of melting point as compared to others.
- Mn has a stable electronic configuration (3d5: half-filled and 4s2: fully filled).
- The 3d electrons are tightly held, resulting in their less delocalization.
- The metallic bond is weaker in Mn, resulting in a lower melting point than Cr.
- Similarly, Group 12 elements (Zn, Cd, Hg) have a fully filled electronic configuration.
- Due to the absence of unpaired electrons, no metallic bonding occurs.
- So, they have low melting points.

Ionization Energy

1st I.E. of d-block metals are greater than those of s-block elements but lesser than those of p-block elements in the same row of the Periodic Table because:
- the d-block metals are smaller in size than the s-block metals, thus they have a greater effective nuclear charge.
- the d-block metals are greater in size than the p-block metals, thus they have a lesser effective nuclear charge.
The increase in ionization energy from Sc to Zn is not regular.
The 1st I.E. of the d-block metals increases slightly and irregularly across the series because:
- Going across the first transition series, the nuclear charge of the elements increases.
- Additional electrons are being added in the inner 3d sub-shell i.e. $(n-1)d$ subshell, the screening effect increases.
- The 4s electrons are shielded more and more.
- The screening effect of the additional 3d electrons opposes the increase in nuclear charge.
- The effective nuclear charge of the elements increases very slowly across the period, as does the Ionization energy.
Successive ionization enthalpies exhibit a similar gradual increase across the first transition series.
The increase in the 3rd and 4th ionization enthalpies across the series is progressively more rapid.
The second ionization enthalpies of both Cr and Cu are higher than those of their next elements, respectively because:
- In the case of Cr, the second ionization enthalpy involves the removal of an electron from a half-filled 3d sub-shell, which has extra stability. Therefore, this second ionization enthalpy is relatively high.
- The case is similar for copper, where its second ionization enthalpy involves the removal of an electron from a fully-filled 3d sub-shell, which also has extra stability. Thus, its second ionization enthalpy is also relatively high.

Metallic Character

All transition elements are metals.
Except for Zn, Cd, and Hg, the transition elements are hard, brittle, and less volatile.
The metallic character of TM is due to their low ionization energies and the number of vacant orbitals in their outermost shell.
The hardness of a metal depends on the strength of the metallic bonds, which arises due to the overlap of unpaired electrons between different metal atoms.

Reducing Character

The reducing character of a metal depends on its tendency to be oxidized:
$M \rightarrow M^{n+} + ne^-$
The reducing potential is measured in terms of reduction potential ( $E^o$ ) of an element.
More negative the value of $E^o$ , the stronger the reducing nature.
D-block elements act as strong reducing agents.
s-block elements are stronger reducing agents than d – block elements.
d – block elements have large ionization energies and very high enthalpies of atomization, as compared to s-block elements.
These result in the reduction of electrode potential values of d – block elements.
Zn has a low electrode potential value despite having a low enthalpy of atomization due to a very high value of Ionization energy.
The d-block metals are comparatively small, and the metallic atoms are closely packed in the metallic lattice. Besides, both the 3d and 4s electrons of the d-block metals participate in metallic bonding by delocalizing into the electron sea. The strength of metallic bond in these metals is thus very strong. In the case of s-block metals, the metallic radius is larger, and most of them do not have close-packed structures. Also, as they have only one or two valence electrons per atom delocalizing into the electron sea, the metallic bond formed is weaker. Therefore, the d-block metals have a much higher melting point than the s-block metals.

Reactivity

Transition metals are very less reactive due to:
- Small atomic size
- High ionization energies
- Due to covalent bonding, they have a high heat of sublimation
- Low heat of hydration

Variable Oxidation States

d-block elements have the ability to show variable oxidation states because:
- 3d & 4s electrons have almost the same energy.
- Both 3d & 4s electrons are available for bond formation.
The stability of a particular oxidation state depends on the nature of the element with which the transition metal enters into compound formation.
Variable oxidation states shown by TM are due to the participation of the inner $(n-1) d$ and the outer ns orbitals.
Except Sc, all elements have a +2 oxidation state.
Except Cu & Zn, all elements have a +3 oxidation state.
The first five elements of the 1st transition series:
- Minimum oxidation state is equal to the electrons in the outer ns orbital.
- Maximum oxidation state is equal to the electrons in the inner $(n-1) d$ and outer ns orbital.
The highest oxidation state observed is that of Mn (+7). This corresponds to the removal of all 3d & 4s electrons.
There is a reduction in the number of oxidation states after Mn because:
- decrease in the number of unpaired electrons and an increase in nuclear charge, which holds the 3d electrons more firmly.
The relative stability of various oxidation states can be correlated with the stability of empty, half-filled, and fully-filled configurations
- e.g., Ti4+ is more stable than Ti3+ (∵ [Ar]3d0 configuration)
- Mn2+ is more stable than Mn3+ (∵ [Ar]3d5 configuration)
- Zn2+ is more stable than Zn+ (∵ [Ar]3d10 configuration)
Compounds of transition metals having +2 and +3 oxidation states generally form ionic bonds with ligands.
Compounds of transition metals showing higher oxidation states, bonds formed are covalent in nature because:
- +2 or +3 O.S. are formed due to the loss of electrons.
- Whereas in higher O.S., the loss of so many electrons is not possible, and the bonds are formed due to the sharing of d-electrons.
Transition metals also form compounds in low oxidation states like +1, 0, and negative.
- When the metal is in zero or low oxidation state, there are no forces of attraction between the metal and the ligands forming the bonds.
- As the metal forms a number of bonds with the ligands, it should result in an increase in –ve charge on the metal.
- So, compounds of TM in the lower or zero oxidation state should be unstable, but they are not.
In such cases, TM having zero or low oxidation states tend to form compounds with ligands like CO, NO, CN-. E.g., Ni(CO)4 & Fe(CO)5
- In both these compounds, the metal has a zero oxidation state.
- The vacant orbitals of the metal accept a lone pair of electrons from the ligands, forming a normal $\sigma$ – covalent bond.
- Simultaneously, electrons from the filled metal orbitals are back-donated to the vacant orbitals of the ligand, forming a $\pi$ bond ( $d \pi - p \pi$ bond).
- This $\pi$ bond reinforces the $\sigma$ bond.
- Due to back donation of electrons from metal to ligands, no accumulation of –ve charge occurs on the metal.
Factors that determine the stability of different oxidation states of transition metals are:
- Ionization energy
- Electrode potential
- Nature of solvent

Ionization Energy and Stability

Gives information about the thermodynamic stability of TM compounds in different oxidation states.
Lesser the ionization energy, the more is the stability of the compound in that particular O.S.
Consider the first four ionization energies of Ni & Pt
- The sum of the first two ionization energies of Ni is less than that of Pt, so Ni2+ complexes are more stable thermodynamically than Pt2+ complexes.
- However, if the sum of the first four ionization energies is considered, then Pt4+ complexes are more stable thermodynamically than Ni4+ complexes, as Pt4+ complexes have a lower value of ionization energies.

Electrode Potential and Stability

It also helps to predict the stability of different oxidation states.
The lower the value of $E^o_{red}$ , the greater is the stability of that oxidation state.
Comparison of $E^o_{red}$ values of Cu+ and Cu2+:
$Cu^+ (aq) + e^- \rightarrow Cu (s) ; E^o_{red} = 0.52V$
$Cu^{2+} (aq) + 2e^- \rightarrow Cu (s) ; E^o_{red} = 0.34V$
Thus, $Cu^{2+}$ O.S. is more stable than $Cu^+$ O.S. in an aqueous solution.
Electrode potential depends on many factors:
- Sublimation energy
- Ionization energy
- Hydration energy
- M (s) → M+ (aq) + e-
- The Oxidation potential involves the following steps:
  - First Step: Involves the isolation of the atom and changing it from the solid state to the gaseous state. This process is called sublimation, and the energy required for this conversion is called the Enthalpy of Sublimation.
  - The second Step involves the ionization (removal of the outermost electron) of the gaseous atom. The energy required for the removal of the outermost electron of the gaseous atom is called Ionization energy.
  - The third Step involves the hydration of the gaseous metal ion. During hydration, energy is liberated & is called Hydration energy.
The oxidation potential depends on the sum of the above three steps:
$\Delta H = \Delta H{sub} + I.E + \Delta H{hyd}$
$Stability \space of \space an \space O.S. \propto \frac{1}{\Delta H}$
The electrode potential values of the first transition series for the $M^{2+}(aq) / M(s)$ are as:
- The electrode potential values show no regular trend due to variation in the (I.E1 + I.E2) and heat of sublimation.
- Eo value of TM is low as compared to Group 2 elements $(ns^2)$ .
- TM has high values of ionization energies and very large energy of atomization as compared to Group 2 elements. These result in the decrease of $E^o$ value of TM.
- Copper has a +ve value of $E^o$ . This is because converting Cu(s) to Cu2+ (aq) requires a very high value of ionization potential, which is not compensated even on the hydration of Cu2+ ion.
- $E^o$ values of TM decrease as we move from left to right in a series, due to the increase in (I.E1 + I.E2).
- The values of $E^o$ of Mn, Ni, and Zn are more –ve than expected.
  - This is due to the stability of the half-filled d - subshell in Mn2+ (3d5) and the completely filled d - subshell in Zn2+ (3d10).
  - The high $E^o$ of Ni2+ is due to the highest –ve energy of hydration of the Ni2+ ion.

Nature of Solvent

The stability of oxidation states in a particular solvent depends on the nature of the solvent.
In that solvent, the TM may undergo oxidation or reduction depending on the conditions under which the reaction is occurring.
E.g., Cr3+ is stable in water, but Cr2+ is unstable.

Trends in M3+ / M2+ Standard Potential

Except Cu and Zn, all elements of the first transition series show +3 O.S.
- Sc3+ has a low value of $E^o$ as in +3 O.S., it has a noble gas configuration.
- The high value of Mn3+ indicates that Mn in +2 O.S. is highly stable due to the half-filled d – subshell.
- The low value of Fe3+ indicates that Fe in +3 O.S. is highly stable due to the half-filled d – subshell.
- The low value of V is due to the stability of V2+.

Formation of Colored Ions

The natural colors of precious gemstones are due to the existence of small quantities of d-block metal ions.
Most of the d-block metals form colored compounds, and most of their complexes are colored too.
The color of the compounds of TM is due to the presence of incompletely filled $(n – 1)d$ orbitals.
The unpaired electrons can be easily promoted from a lower energy level to a higher energy level in the same d – subshell.
For the d-block elements, the five 3d orbitals are degenerate (have the same energy) in gaseous ions.
However, when compound formation occurs, the 3d orbitals split into 2 sets of orbitals having different energies. This is called crystal field splitting.
The amount of energy required to excite the electrons from a lower to a higher energy level within the same d – subshell equals/corresponds to the energies of certain colors in invisible light.
When light falls on a TM compound, it absorbs visible light (corresponding to some color) of a certain wavelength, causing the excitation of d – electrons.
Light of wavelengths of other regions of the visible light spectrum will be reflected or transmitted.
The color seen is always complementary to the color of the absorbed wavelength of visible light.
d-block metal ions have specific colors.
For d-d electronic transition and absorption of visible light to occur, there must be unpaired d electrons in the transition metal atoms or ions.
Sc3+ and Zn2+ are colorless due to the empty and fully-filled 3d sub-shell, respectively.
The colors of hydrated metal ions are determined by the oxidation states of the particular d-block elements, e.g., $Fe^{2+}(aq)$ is green while $Fe^{3+}(aq)$ is yellow.
TM ions having completely filled d-orbitals are colorless.

Magnetic Properties

Most of the Transition metals are paramagnetic in nature due to the presence of unpaired electrons in d – orbitals.
Magnetic character is expressed in terms of magnetic moment.
The greater the number of unpaired electrons, the greater is the paramagnetic character and the greater is the magnetic moment.
The magnetic moment is expressed in Bohr Magneton (B.M.).
The magnetic moment is only due to the spin of the electrons.
This is confirmed due to the close agreement between theoretical and calculated values of the magnetic moment of TM ions.
The magnetic moment due to spin contribution only can be calculated by:
$u = \sqrt{n (n+2)} \space B.M.$
Where n stands for the number of unpaired electrons.
In addition to Paramagnetism and Diamagnetism, the compounds of Iron and some other metals show another kind of magnetic behavior known as Ferro magnetism.
In some materials, the permanent atomic magnetic moments have a strong tendency to align themselves even without any external field.
These materials are said to be Ferromagnetic materials.
Some of the examples of ferromagnetic materials are cobalt, iron, nickel, gadolinium, and dysprosium.
The ferromagnetic materials are those substances that exhibit strong magnetism in the same direction of the field when a magnetic field is applied to it.
In the solid state, the metal ions of ferromagnetic substances are grouped together into small regions called Domains.
They get their strong magnetic properties due to the presence of magnetic domains.
In these domains, large numbers of atom's moments ( $10^{12}$ to $10^{15}$ ) are aligned parallel so that the magnetic force within the domain is strong.
When a ferromagnetic material is in the unmagnitized state, the domains are nearly randomly organized, and the net magnetic moment s gets canceled.
When a magnetizing field is applied, the domains become aligned in the direction of the magnetic field.
Ferromagnetic substances can be classified into three types depending upon on the alignment of magnetic domains:
- Ferromagnetic substances
 - They have large, positive susceptibility to an external magnetic field.
 - The magnetic moment of the domains gets spontaneously aligned in the direction of the applied magnetic field.
 - They can retain their magnetic properties after the external field has been removed.
 - It means it becomes permanently magnetized even after the removal of the external magnetic field.
 - Example: Ni, Co, $Fe2O3$ , $CrO_2$ , etc.
- Anti-ferromagnetic substances
 - In this, the adjacent ions/ domains that behave as tiny magnets spontaneously align themselves into opposite or antiparallel arrangements throughout the material.
 - The magnetism from magnetic atoms or ions oriented in one direction is canceled out by the set of magnetic atoms or ions that are aligned in the reverse direction.
 - So that the compound exhibits almost no gross external magnetism.
 - This type of behavior is observed at low temperatures.
 - Examples of anti-ferromagnetic substances are MnO, FeO, NiO, $MnO2$ , $V2 O_3$ .
- Ferrimagnetic Substances
 - It has a population of atoms with opposing magnetic moments, as in anti-ferromagnetic substances.
 - However, in ferrimagnetic materials, the opposing magnetic moments are unequal.
 - This results in spontaneous magnetization, i.e., a net magnetic moment is observed.
 - Such compounds are weakly attracted by the magnetic field.
 - This happens when the populations consist of different materials or ions (such as Fe2+ and Fe3+ ).
 - An example of ferrimagnetic substances is $Fe3O4$ (made up of FeO & $Fe2O3$ ), and ferrites have the formula $M^{2+}Fe2O4$ ; where M can be either Mg, Cu, Zn, etc.
However, all magnetically ordered solids, i.e., ferromagnetic, anti-ferromagnetic, and ferrimagnetic, change to the paramagnetic state at high temperature.
This happens due to the randomization of their spins at high temperature.
Curie Temperature Tc is the temperature at which certain materials lose their permanent magnetic properties, to be replaced by induced magnetism.
In other words, it is the temperature above which a ferromagnetic material becomes paramagnetic.
The Néel temperature or magnetic ordering temperature, TN, is the temperature above which an anti-ferromagnetic material becomes paramagnetic.
It means the thermal energy becomes large enough to destroy the microscopic magnetic ordering within the material.
$Fe3O4$ at 5800C loses its ferrimagnetic character and becomes paramagnetic in character.

Complex Formation

Transition metals have a high tendency to form complexes.
The high tendency of transition metal atoms/ions to form complexes is due to:
- The small size of transition metal atoms/ions
- High nuclear charge
- Availability of vacant d- orbitals of suitable energies in transition metal atoms/ions to form bonds with ligands
The bonds between the central metal atom/ion are formed due to the sharing of an electron pair of the ligand with the metal.
The bonds formed are Coordinate covalent bonds, and so the complexes are called Coordinate complexes.
Coordinate complexes have different geometries (linear, tetrahedral, square planar, octahedral, etc.) depending on the number of ligands forming bonds with the central metal atom/ion.
The stability of TM complexes increases with an increase in the atomic number of the element.
The greater the oxidation state of an element, the smaller is its atomic size, and the more will be its stability.
For example: $Co^{2+}$ does not form a stable complex with ammonia (NH3), whereas $Co^{3+}$ does and forms a complex of the type $[Co(NH3)6 ]^{3+}$ .

Formation of Interstitial Compounds

Transition metals form interstitial compounds with non-metallic elements like H, B, C, and N.
Transition metals have vacant spaces in their lattices, and small atoms of H, B, C, and N get trapped in these vacant spaces.
The interstitial spaces get filled up, leading to transition metals becoming hard and rigid.
Chemical properties of the interstitial compounds resemble those of the parent compounds, but their physical properties are different.

Catalytic Property

Many transition metals (V, Cr, Mn, Fe, Co, Ni, etc.) and their compounds act as good catalysts for various reactions.
Iron – molybdenum (Fe – Mo) is used as a catalyst in the synthesis of ammonia by Haber’s process.
Nickel is used as a catalyst in hydrogenation reactions in organic chemistry.
Vanadium pentoxide is used for oxidation of SO2 to SO3 in contact process.
The transition metals form suitable reaction intermediates with the substrate.
The formation of reaction intermediates involves the use of empty d- orbitals of the TM.
The reaction intermediates give reaction paths of lower activation energies, resulting in an increase in the rate of reaction.
Consider the conversion of SO2 to SO3 in which V2O5 is used as a catalyst.
A molecule of SO2 is adsorbed on the surface of the solid catalyst V2O5.
V2O5 changes to V2O4, giving out an O atom, which is taken up by SO2 to form SO3.
$V2O5 + SO2 \rightarrow SO3 + V2O4$
Catalyst Vanadium tetraoxide
V2O4 gets converted back into V2O5 by reaction with oxygen.
$V2O4 + O2 \rightarrow V2O_5$

Alloy Formation

Transition metals form a large number of alloys.
Transition metals have approximately the same size, due to which atoms of one metal can substitute atoms of other metals in its crystal lattice.
Solid alloys are formed by cooling a mixture solution of two or more transition metals.
e.g., a solution of manganese and iron gives an alloy called manganese steel. The solid alloys are formed by cooling a mixture solution of two or more transition metals.