D Block Elements

Step 1: Definition.
Transition elements are d-block elements that have partially filled d-orbitals in their ground state or in one of their oxidation states.
Step 2: Characteristics.
Some important characteristics of transition elements are: \begin{enumerate} \item
Variable oxidation states: Transition metals commonly exhibit more than one oxidation state (e.g., Fe and Fe ). \item
Formation of coloured compounds: Due to d–d electronic transitions, many transition metal salts and complexes are coloured. \item
Catalytic properties: They act as good catalysts because of variable oxidation states and ability to provide active sites (e.g., V O in contact process). \item
Formation of complex compounds: Transition metals form coordination complexes with ligands due to availability of vacant d-orbitals (e.g., [Fe(CN) ] ). \end{enumerate} Conclusion:
Thus, transition elements show variable oxidation states, coloured ions, catalytic activity, and complex formation as their main characteristics.

Step 1: Definition of Transition Elements.
Transition elements are the elements found in the d-block of the periodic table. They are characterized by having partially filled d-orbitals in at least one of their oxidation states. These elements tend to form colored compounds, exhibit variable oxidation states, and are often good conductors of electricity.
Step 2: Electronic Configuration of the First Transition Series.
The elements of the first transition series are from Scandium (Sc) to Zinc (Zn). Their electronic configurations are as follows: - Scandium (Sc): [Ar] 3d 4s - Titanium (Ti): [Ar] 3d 4s - Vanadium (V): [Ar] 3d 4s - Chromium (Cr): [Ar] 3d 4s - Manganese (Mn): [Ar] 3d 4s - Iron (Fe): [Ar] 3d 4s - Cobalt (Co): [Ar] 3d 4s - Nickel (Ni): [Ar] 3d 4s - Copper (Cu): [Ar] 3d 4s - Zinc (Zn): [Ar] 3d 4s
Step 3: Paramagnetism of Transition Metals.
Transition metals are often paramagnetic because they have unpaired electrons in their d-orbitals, which make them attracted to an external magnetic field. The paramagnetism in transition metals is due to the presence of these unpaired electrons in the metal's electronic configuration. The number of unpaired electrons determines the strength of the paramagnetism.

(i) Coloured compounds.
Transition metals have partially filled d-orbitals. When light falls, electrons undergo d–d transitions, absorbing certain wavelengths, and the complementary colour is observed. Example: (violet), (blue).
(ii) Paramagnetism.
Transition metals and their ions often contain unpaired d-electrons. These unpaired electrons produce magnetic moments, leading to paramagnetic behaviour. Example: Fe , Mn .

Step 1: Why ions show colours.
The formation of coloured ions, especially in transition metals, is due to the presence of unpaired d-electrons. These electrons undergo d–d electronic transitions when light is absorbed, producing colour.

Step 2: Analyzing options.
- (A) Unpaired electrons: Correct. They are responsible for colour due to electronic transitions.
- (B) Paired electrons: Do not cause colour because no electronic transition occurs.
- (C) Non-bonded electrons: They may cause other properties but not coloured ions.
- (D) None of the above: Incorrect, since unpaired electrons are responsible.

Step 3: Conclusion.
Therefore, coloured ions are formed due to the presence of unpaired electrons.

Final Answer:

Step 1: Industrial preparation of Ammonia.
Ammonia is prepared industrially by the
Haber process. It involves direct combination of nitrogen and hydrogen gases under specific conditions. Step 2: Flow diagram of Haber process.

Flow Diagram: Step 3: Reaction of Ammonia with Copper ion ( ).
Ammonia reacts with copper(II) ions to form a deep blue complex: Step 4: Reaction of Ammonia with Silver ion ( ).
Ammonia reacts with silver ions to form a soluble complex: Conclusion:
Ammonia is prepared industrially by the Haber process under high temperature, pressure, and Fe catalyst. It also forms characteristic coordination complexes with (deep blue) and (diammine silver complex).

Step 1: Electronic configuration of Cr (Z = 24).
The expected configuration is: But due to extra stability of half-filled orbitals, the actual configuration is: Step 2: Electronic configuration of Cu (Z = 29).
The expected configuration is: But due to extra stability of fully filled orbitals, the actual configuration is: Step 3: Two main characteristics of transition elements.
\begin{enumerate} \item
Variable oxidation states: Transition metals exhibit multiple oxidation states due to the involvement of both and electrons in bonding. \item
Formation of coloured compounds: Due to the presence of partially filled -orbitals, they show electronic transitions which give rise to characteristic colours. \end{enumerate} Conclusion:
Chromium and copper show exceptional electronic configurations due to stability of half-filled and fully-filled -orbitals. Transition elements are unique for their variable oxidation states and coloured compounds.

(i) Variation of atomic radii across a series:
Across a transition series, atomic radii decrease slightly from left to right. This happens because as we move across the period, the nuclear charge increases due to the addition of protons.
However, the added electrons enter the penultimate orbital, providing poor shielding of the outer electrons. Thus, the increased nuclear charge pulls electrons closer, reducing atomic size.
After the mid-series, the effect of electron-electron repulsion in the -orbitals balances the nuclear pull, so the decrease becomes very small. (ii) Paramagnetic nature of transition metal ions:
Transition metal ions generally contain unpaired electrons in their orbitals.
Paramagnetism arises due to the presence of these unpaired electrons, which have magnetic moments associated with their spins. Hence, most transition metal ions exhibit paramagnetism. (iii) Variable oxidation states in transition elements:
Transition elements show variable oxidation states because the energy difference between the and orbitals is very small.
As a result, both and electrons can take part in bonding.
Thus, transition metals can exhibit multiple oxidation states (e.g., Fe shows +2 and +3, Mn shows +2 to +7). Conclusion:
The unique electronic configuration of transition metals ( ) explains their small variation in radii, paramagnetism, and variable oxidation states.