(a) (i) Silver atom has completely filled d-orbitals in its ground state, it is still considered to be a transition element. Justify the statement. (ii) Why are $E^\circ_{M^{2+}/M}$ values of Mn and Zn more negative than expected ? (iii) Why do transition metals form alloys ? OR

Question

(a) (i) Silver atom has completely filled d-orbitals in its ground state, it is still considered to be a transition element. Justify the statement. (ii) Why are $E^\circ_{M^{2+}/M}$ values of Mn and Zn more negative than expected ? (iii) Why do transition metals form alloys ? OR (b) Answer the following questions on the basis of the figure given below : The figure is a plot of enthalpy of atomisation ($\Delta_aH^\circ$ / kJ mol$^{-1}$, y-axis) against atomic number (x-axis) showing three curves labelled Series 1, Series 2 and Series 3 (the 3d, 4d and 5d transition series); each curve rises to a maximum near the middle of its series and falls towards the ends, with Series 3 generally the highest and Series 1 the lowest. (i) Which element in 3d series has lowest enthalpy of atomisation ? (ii) Why do metals of the second and third series have greater enthalpies of atomisation ? (iii) Why are enthalpies of atomisation of transition metals quite high ?

Accepted Answer

> [!TLDR]
> **Part (a):** Silver is a transition element because $\ce{Ag^{2+}}$ has an incomplete $4d^9$ shell; the $E^\circ_{M^{2+}/M}$ values of Mn and Zn are more negative than expected because of the extra stability of the half-filled $3d^5$ and fully-filled $3d^{10}$ configurations of their $M^{2+}$ ions; transition metals form alloys because their similar atomic sizes let atoms substitute in the lattice. **Part (b) (OR):** in the 3d series **Zn** has the lowest enthalpy of atomisation (filled $3d^{10}$, weak bonding); 4d/5d metals have higher values because their diffuse orbitals overlap more strongly; transition metals in general have high enthalpies of atomisation because many unpaired $d$-electrons give strong metallic bonding.

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## Part (a)

### (i) Silver as a transition element

By the IUPAC definition, a transition element has an incomplete $d$-subshell **either in its ground state or in one of its common oxidation states**. Silver's ground state is $[	ext{Kr}]\,4d^{10}5s^1$ and its usual $+1$ ion $\ce{Ag^+}$ is $[	ext{Kr}]\,4d^{10}$ — both filled. However, silver also exists in the **$+2$ state** (for example in $\ce{AgF2}$), where the configuration is $[	ext{Kr}]\,4d^9$, an **incomplete** $d$-subshell. Because it can form a cation with a partly filled $d$-subshell, silver is classed as a transition element.

> [!WARNING]
> A filled $d$-subshell in the ground state alone does not exclude an element — the criterion also covers any common oxidation state.

### (ii) Why $E^\circ_{M^{2+}/M}$ of Mn and Zn is more negative than expected

$E^\circ_{M^{2+}/M}$ reflects the tendency of the metal to be reduced from $M^{2+}$; a more negative value means the $M^{2+}$ ion is unusually stable and reluctant to be reduced.
- **Mn:** $\ce{Mn^{2+}}$ is $[	ext{Ar}]\,3d^5$, a **half-filled** subshell that is extra-stable through exchange energy and symmetry. Forming this stable ion is easy, so $E^\circ$ is more negative than the trend.
- **Zn:** $\ce{Zn^{2+}}$ is

d-block elements	Transition elements
Groups 3 to 12	Groups 3 to 11 (excluding Zn, Cd, Hg)
All have d electrons	Must have incomplete d sub-shell in atom or common cation

d-block elements	Transition elements
Groups 3 to 12	Groups 3 to 11 (excluding Zn, Cd, Hg)
All have d electrons	Must have incomplete d sub-shell in atom or common cation

Q.(a)

Part (a)

(i) Silver as a transition element

(ii) Why EM2+/M∘E^\circ_{M^{2+}/M}EM2+/M∘​ of Mn and Zn is more negative than expected

(iii) Why transition metals form alloys

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(ii) Why $E^\circ_{M^{2+}/M}$ of Mn and Zn is more negative than expected

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