Transition elements (also known as transition metals) are a fascinating group in the periodic table. One of their most remarkable characteristics is their ability to exhibit variable oxidation states—that is, they can lose different numbers of electrons in different compounds. But why does this happen? Let’s explore in depth.
🔍 What Are Transition Elements?
Transition elements are elements found in the d-block of the periodic table. These include metals like iron (Fe), copper (Cu), manganese (Mn), and chromium (Cr), among others. A key feature of these elements is the incomplete filling of their d-orbitals in at least one of their oxidation states.
⚡ Understanding Oxidation States
The oxidation state (or oxidation number) of an element refers to the charge it appears to carry when combined with other elements in a compound. For example:
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In FeCl₂, iron has a +2 oxidation state.
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In FeCl₃, iron has a +3 oxidation state.
So, iron shows variable oxidation states: +2 and +3.
The Role of d-Orbitals
The primary reason transition elements show variable oxidation states lies in their electronic configuration. Transition metals have electrons in both s and d orbitals, and the energy difference between these orbitals is relatively small. This allows electrons from both the s and d orbitals to be involved in bonding, leading to multiple oxidation states.
Electronic Configuration
Consider the general electron configuration of a transition metal: [Ar] 3dⁿ 4s², where n is the number of d electrons. When forming ions, transition metals lose electrons to achieve a positive oxidation state. Unlike main group elements, which lose electrons only from their outermost s or p orbitals, transition metals can lose:
Electrons from the 4s orbital.
Electrons from the 3d orbitals, depending on the chemical environment.
For example:
Scandium (Sc): [Ar] 3d¹ 4s²
Loses two 4s electrons to form Sc²⁺ ([Ar] 3d¹).
Can also lose the 3d electron to form Sc³⁺ ([Ar]), a stable noble gas configuration.
Iron (Fe): [Ar] 3d⁶ 4s²
Loses two 4s electrons to form Fe²⁺ ([Ar] 3d⁶).
Loses an additional 3d electron to form Fe³⁺ ([Ar] 3d⁵), which is stable due to a half-filled d subshell.
The ability to lose varying numbers of 4s and 3d electrons allows transition metals to form ions with different charges, corresponding to different oxidation states.
Small Energy Gap Between 4s and 3d Orbitals
The energy difference between the 4s and 3d orbitals is small, especially as the effective nuclear charge increases across the d-block. This small energy gap means that the energy required to remove a 3d electron (after losing 4s electrons) is not significantly higher than removing a 4s electron. As a result, transition metals can form ions with a range of oxidation states without a prohibitive energy cost.
🧪 Why Do Transition Metals Show Variable Oxidation States?
There are several key reasons:
1. Small Energy Difference Between (n-1)d and ns Orbitals
In transition elements, the electrons are filled in both the (n-1)d and ns orbitals. For example:
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In iron (Fe): Electronic configuration = [Ar] 3d⁶ 4s²
When forming ions, electrons are first removed from the 4s orbital, then from 3d orbitals. Since the energy gap between 4s and 3d orbitals is very small, different numbers of electrons can be removed fairly easily, leading to multiple oxidation states.
➡️ Example:
Fe → Fe²⁺ (by removing 2 electrons: 4s²)
Fe → Fe³⁺ (by removing 3 electrons: 4s² + 1 from 3d)
2. Participation of d-Electrons in Bonding
Unlike s- and p-block elements, transition metals have valence d-electrons that also participate in bonding. As the oxidation state increases, more d-electrons are used in bonding. This gives them a wide range of possible oxidation states depending on the chemical environment.
➡️ Example:
Manganese (Mn) shows oxidation states from +2 to +7 in compounds like MnO, MnCl₂, KMnO₄, etc.
3. Stability of Different Oxidation States Varies Across the Series
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In the early transition series (Sc to Mn), higher oxidation states are more stable.
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In the later series (Fe to Zn), lower oxidation states become more stable.
This is due to increased nuclear charge and the relative energy of the d-orbitals.
4. Formation of Stable Ions with Different Ligands
Transition metals form complexes with ligands like water, ammonia, chloride, etc. Different ligands stabilize different oxidation states due to ligand field stabilization energy (LFSE). Hence, the same metal can form compounds with different oxidation states depending on the ligands.
5. Exchange of Electrons Without Large Energy Barriers
Since the ionization energies of successive electrons are not drastically different (especially for the first few), transition metals can lose more than just their outermost electrons without much energy input.
📚 Examples of Variable Oxidation States
| Element | Common Oxidation States |
|---|---|
| Iron (Fe) | +2, +3 |
| Copper (Cu) | +1, +2 |
| Chromium (Cr) | +2, +3, +6 |
| Manganese (Mn) | +2, +4, +7 |
| Vanadium (V) | +2, +3, +4, +5 |
🧠 Conclusion
The ability of transition metals to show variable oxidation states is primarily due to:
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The closeness in energy of the (n-1)d and ns orbitals,
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The involvement of d-electrons in bonding,
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The formation of stable complexes with various ligands, and
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Their ability to lose different numbers of electrons easily.
This property is what makes transition elements so versatile in chemistry—from colorful compounds and catalysts to biological functions and industrial applications.
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