Ionisation Enthalpy, Oxidation States — Explained

Detailed Explanation

The properties of ionisation enthalpy and oxidation states are central to understanding the unique chemistry of transition elements. These elements, located in the d-block of the periodic table, exhibit distinct trends and characteristics compared to s-block and p-block elements, largely due to the involvement of their (n-1)d electrons.

Ionisation Enthalpy of Transition Elements

Definition and Successive Ionisation Enthalpies:

Ionisation enthalpy (IE) is the energy required to remove an electron from a gaseous atom or ion. For transition elements, we consider the first, second, and successive ionisation enthalpies. The first ionisation enthalpy ($IE_1$) removes an electron from a neutral atom.

The second ionisation enthalpy ($IE_2$) removes an electron from the $M^+$ ion, and so on. A crucial point is that successive ionisation enthalpies always increase ($IE_1 < IE_2 < IE_3 dots$) because removing an electron from a positively charged species requires overcoming a stronger electrostatic attraction.

General Trends Across a Period (e.g., 3d Series):

As we move across a transition series from left to right (e.g., Sc to Zn), the nuclear charge increases steadily. One might expect a consistent increase in ionisation enthalpy due to this increasing nuclear attraction. However, the trend is not as smooth as in s-block or p-block elements. The $IE_1$ values generally show a slight and irregular increase across the 3d series. This irregularity is primarily due to:

1
Increasing Nuclear Charge: — As atomic number increases, the number of protons increases, pulling the electrons more strongly towards the nucleus.
2
Shielding Effect of d-electrons: — The added d-electrons effectively shield the outer ns electrons from the increasing nuclear charge. This shielding is not perfectly efficient, but it partially counteracts the effect of increased nuclear charge.
3
Electron-electron Repulsion: — As more electrons are added to the d-subshell, electron-electron repulsions increase, which can slightly reduce the energy required to remove an electron.

Specific Irregularities in 3d Series $IE_1$:

Chromium (Cr) and Copper (Cu): — These elements show slightly higher $IE_1$ values than their neighbors. For Cr ($[Ar]3d^54s^1$), the $4s^1$ electron is removed, leaving a stable half-filled $3d^5$ configuration. For Cu ($[Ar]3d^{10}4s^1$), the $4s^1$ electron is removed, leaving a stable fully-filled $3d^{10}$ configuration. The stability associated with these configurations makes it slightly harder to remove the first electron.
Manganese (Mn) and Zinc (Zn): — These elements have relatively high $IE_2$ values. For Mn ($[Ar]3d^54s^2$), removing the first $4s$ electron gives $Mn^+$ ($[Ar]3d^54s^1$). Removing the second electron from $4s^1$ is relatively easy. However, $IE_3$ for Mn is very high because it involves removing an electron from the stable half-filled $3d^5$ configuration of $Mn^{2+}$ ($[Ar]3d^5$). For Zn ($[Ar]3d^{10}4s^2$), removing the first $4s$ electron gives $Zn^+$ ($[Ar]3d^{10}4s^1$). Removing the second $4s$ electron is relatively easy. However, $IE_3$ for Zn is extremely high as it would involve breaking the stable fully-filled $3d^{10}$ configuration of $Zn^{2+}$ ($[Ar]3d^{10}$). This explains why $Zn^{2+}$ is the most common and stable oxidation state for zinc.

Trends Down a Group:

Moving down a group (e.g., from 3d to 4d to 5d series), atomic size generally increases, and the outermost electrons are further from the nucleus. This typically leads to a decrease in ionisation enthalpy.

However, for transition elements, the trend is complicated by the lanthanoid contraction. The poor shielding by 4f electrons in the lanthanoids causes a significant increase in effective nuclear charge for the subsequent 5d elements.

As a result, the 5d elements often have ionisation enthalpies comparable to, or even slightly higher than, their 4d counterparts, despite being larger atoms. For example, $IE_1$ for Hf is similar to Zr.

Oxidation States of Transition Elements

Variable Oxidation States:

The most distinctive feature of transition elements is their ability to exhibit multiple oxidation states. This contrasts sharply with s-block elements (typically fixed $+1$ or $+2$) and p-block elements (which show fewer variations, often separated by two units). The primary reason for this variability is the very small energy difference between the $(n-1)d$ orbitals and the $ns$ orbitals. This allows electrons from both these subshells to participate in bond formation.

Common Oxidation States in 3d Series:

Scandium (Sc): — Only exhibits $+3$ oxidation state ($[Ar]3d^14s^2 \rightarrow Sc^{3+}$ with $[Ar]$ configuration). All three valence electrons are lost.
Titanium (Ti): — Shows $+2, +3, +4$. $+4$ is most stable ($[Ar]3d^24s^2 \rightarrow Ti^{4+}$ with $[Ar]$ configuration).
Vanadium (V): — Shows $+2, +3, +4, +5$. $+5$ is most stable in compounds like $V_2O_5$ ($[Ar]3d^34s^2 \rightarrow V^{5+}$ with $[Ar]$ configuration).
Chromium (Cr): — Shows $+2, +3, +6$. $+3$ is common, $+6$ is seen in chromates and dichromates ($[Ar]3d^54s^1$).
Manganese (Mn): — Exhibits the widest range, from $+2$ to $+7$. $+2$ and $+7$ are particularly stable. ($[Ar]3d^54s^2$).
Iron (Fe): — Shows $+2, +3$. $+3$ is generally more stable than $+2$ in many compounds ($[Ar]3d^64s^2$).
Cobalt (Co): — Shows $+2, +3$. $+2$ is more common in simple salts, $+3$ is stable in complexes ($[Ar]3d^74s^2$).
Nickel (Ni): — Primarily $+2$, sometimes $+3$ or $+4$ in complexes ($[Ar]3d^84s^2$).
Copper (Cu): — Shows $+1, +2$. $+2$ is more common and stable in aqueous solutions ($[Ar]3d^{10}4s^1$).
Zinc (Zn): — Only exhibits $+2$ oxidation state ($[Ar]3d^{10}4s^2 \rightarrow Zn^{2+}$ with $[Ar]3d^{10}$ configuration).

Stability of Oxidation States:

Lower Oxidation States (+2, +3): — These are common for most transition metals. The $+2$ state typically arises from the loss of the two $ns$ electrons. The $+3$ state involves the loss of two $ns$ electrons and one $(n-1)d$ electron.
Higher Oxidation States: — The highest oxidation state generally increases from group 3 to group 7 (Sc to Mn) and then decreases. The maximum oxidation state often corresponds to the sum of $ns$ and $(n-1)d$ electrons, especially for elements up to Mn. For example, Mn ($3d^54s^2$) can show $+7$ (in $MnO_4^-$).
Factors Affecting Stability:

* Electronic Configuration: Half-filled ($d^5$) and fully-filled ($d^{10}$) d-orbitals confer extra stability. For example, $Mn^{2+}$ ($3d^5$) is very stable. $Zn^{2+}$ ($3d^{10}$) is very stable.

* Electronegativity of Ligands/Oxidizing Agents: Higher oxidation states are more stable when the metal is bonded to highly electronegative elements like oxygen or fluorine (e.g., $MnO_4^-$, $CrO_4^{2-}$).

* Acidic/Basic Character of Oxides: Oxides in lower oxidation states are generally basic (e.g., $MnO$), while those in higher oxidation states are acidic (e.g., $Mn_2O_7$). Intermediate oxidation states can be amphoteric (e.

g., $Cr_2O_3$). * Aqueous Stability: The stability of an oxidation state in aqueous solution is influenced by hydration enthalpy and lattice enthalpy (for solid compounds).

Disproportionation:

Some oxidation states can undergo disproportionation, where an element in an intermediate oxidation state simultaneously oxidizes and reduces itself. For example, $Cu^+$ in aqueous solution is unstable and disproportionates into $Cu^{2+}$ and $Cu$:

$$2Cu^+(aq) \rightarrow Cu^{2+}(aq) + Cu(s)$$

This occurs when the stability of the lower and higher oxidation states is significantly greater than the intermediate one.

NEET-Specific Angle:

For NEET, understanding the *reasons* behind these trends and exceptions is paramount. Questions often test the stability of specific oxidation states, the maximum oxidation state exhibited by an element, or the relative ionisation enthalpies of elements within a series or group.

Pay close attention to the role of $d^5$ and $d^{10}$ configurations in conferring stability, and the impact of lanthanoid contraction on 5d series properties. Remember that for transition metals, electrons are first removed from the $ns$ orbital, then from the $(n-1)d$ orbital, when forming ions.