Key Trends in Oxidation States Across Transition Series
What are Transition Elements?
The d-block elements are called transition elements. These are the elements that have incompletely filled (partly filled) d-subshells in their ground state or in any one of their oxidation states. Here, we will discuss the trends in the stability of oxidation states of transition metals.
The transition metals exhibit a variable number of oxidation states in their compounds. This is one of the notable features of the transition elements. Few elements show exceptions for this case, most of these show variable oxidation states. These different oxidation states are relatable to the electronic configuration of their atoms. The stability of oxidation states for transition metals depends on the electronic configuration of the elements. For example, oxidation states of transition metals exhibited by elements of the first series are listed in the table given below.
Oxidation states overview
The oxidation state of an element can be defined as the degree of capacity e of an element in a chemical compound to lose electrons from its valence shell full stop is also referred to as the degree of oxidation. Transition elements exhibit a wide range of oxidation states in their compounds. For example, magnesium shows a wide range of oxidation states that starts from +2 and goes up to +7 in its various compounds. However, some of the elements exhibit very few oxidation States. The elements that are included in very few oxidation states are zinc and scandium. The very few oxidation states of any of the elements are because of the fact that they have very few electrons to lose from their valence shells. For example, scandium has too many d electrons and hence very few orbitals to share electrons with others for very high valence. The variable oxidation states in valence electrons are due to the very few amount of electrons that are filled in d-orbital in such a way that their oxidation states differ from each other by a unitary.
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Stability of oxidation states
Chromium manganese and Cobalt shows a very high oxidation State. Manganese does not exhibit a +7 oxidation state in the case of halides but in compound MnO3F the +7 oxidation state is known. ΔhydH of Cu+2 is more than Cu+, which compensates for the second ionisation enthalpy of Cu, therefore the Cu+2 (aq) is known to be more stable than Cu+ (aq). In Mn2O7 the oxidation state of Mn is +7 which proves that the oxidation state of manganese is higher in case of its oxides. The Mn oxide, Mn2O7 shows a higher oxidation states in comparison to Mn fluorides, MnF4 as oxygen has the tendency to form multiple bonds with metals. Mn-O-Mn bridge is included in Mn2O7 where each Mn is tetrahedrally surrounded by O’s. The other elements in the compound such as such as Fe2O3 and V2O4 exhibits +3 and +4 oxidation states respectively. Due to the inert pair effect in the p-block elements, the heavier elements prefer a low oxidation state which is just the opposite in case of the d-block elements. As in group 6, Mo (VI) is found to have higher stability in comparison to Cr (VI).
Oxidation State of Transition Elements
The existence of the transition elements in different oxidation states means that their atoms can lose a different number of electrons. This is due to the participation/contribution of inner (n-1) d-electrons in addition to outer ns-electrons because the energies of the ns and (n-1)d-subshells are almost equal. For example, scandium has the electronic configuration 3d14s2. The oxidation state of sc is +2 when it uses both of its 4s- electrons for bonding. It can also show the oxidation state of +3 when it uses its two s- electrons and one d- electron.
The other atom similarly shows oxidation states equal to ns and (n-1) d- electrons. The oxidation state of zn is +2. The chromium shows variable oxidation states; +2, +3, +4, +5, +6. The highest oxidation state of chromium is +6.
As we know the oxidation state of Zn is +2. It does not show a variable oxidation state. Therefore, does not consider a transition element.
Variable Oxidation State of Metals
The elements of the second and third transition metals series also exhibit variable oxidation states are given below:
Second Transition Series
Third Transition Series
It can be noted from the above tables that the stability of a given oxidation state is dependent upon the nature of the elements with which the metal is combined. The highest oxidation states are shown in the compounds of fluorides and oxides because fluorine and oxygen are the most electronegative elements.
The elements that show the greatest (high) number of oxidation states occur in or near the middle of the series. For example, in the first series of the d-block elements, manganese exhibits all the oxidation states from +2 to +7. The small number of oxidation states at the extreme left-hand side end is due to the lesser number of electrons to lose or share. On the other hand, at the extreme right-hand side end, it is due to a large number of d- electrons so that only fewer orbitals are available in which the electron can share with other compounds for higher valence.
In the +2 oxidation state (O.S) and +3 oxidation state, the bonds formed are mostly ionic. In the compounds or molecules of higher oxidation states (generally formed with oxygen and fluorine), the bonds are essentially covalent. Therefore, the bonds in +2 O.S and +3 O.S are generally formed by the loss of two or three electrons, while the bonds in higher oxidation states are formed by sharing of d- electrons.
Did You Know?
The variable oxidation state of transition elements is due to the participation or contribution of inner (n-1) d and outer ns- electrons.
The lowest oxidation state corresponds to the number of ns orbital electrons.
Except for scandium, the most common oxidation state of the first-row transition elements is +2 which arises due to the loss of two 4s- electrons. This means that after scandium 3d block orbitals become more stable and, therefore, are lower in energy than the 4s-orbitals. As a result, electrons are first removed from 4s- orbitals.
FAQs on Transition Elements: Understanding Oxidation States
1. Why do transition elements exhibit variable oxidation states?
Transition elements show variable oxidation states because the energy difference between the outer (n)s orbital and the inner (n-1)d orbital is very small. This allows electrons from both subshells to participate in chemical bonding. Consequently, they can lose a varying number of electrons, leading to different oxidation states in their compounds.
2. What is the most common oxidation state for the first-row transition elements and why?
Except for scandium, the most common oxidation state for the first-row transition elements is +2. This arises from the loss of the two electrons from the outer 4s orbital. After scandium, the 3d orbitals become progressively more stable, so the 4s electrons are typically lost first during ionisation.
3. How is the stability of a particular oxidation state in transition elements determined?
The stability of an oxidation state is determined by the overall thermodynamic stability of the ion. Key factors include:
- Ionisation Enthalpy: The energy required to remove electrons.
- Hydration or Lattice Enthalpy: The energy released when the ion is in solution or forms a crystal lattice.
- Electronic Configuration: States corresponding to half-filled (d⁵) or completely-filled (d¹⁰) d-orbitals, such as in Mn²⁺ or Zn²⁺, are particularly stable.
4. Why do elements at the beginning and end of the transition series show fewer oxidation states?
Elements at the beginning of a series (like Scandium) have very few electrons to lose from their d-orbitals, limiting their options. Elements at the end of a series (like Zinc) have nearly or completely full d-orbitals, making it energetically unfavourable to unpair and involve them in bonding. This results in less variability in oxidation states compared to elements in the middle.
5. Which transition elements show the highest oxidation state, and in what type of compounds?
The highest oxidation states are typically found in elements near the middle of the series, such as Manganese (Mn) which can show up to +7, and Osmium (Os) which can show up to +8. These highest oxidation states are stabilised in compounds with highly electronegative elements like oxygen and fluorine (e.g., in Mn₂O₇ and OsO₄), as these elements are strong enough to draw out multiple electrons from the metal atom.
6. Why is Cu²⁺(aq) more stable than Cu⁺(aq) in aqueous solutions?
Although removing the second electron from copper requires significant energy (high second ionisation enthalpy), the Cu²⁺ ion has a much higher charge density than the Cu⁺ ion. This allows the Cu²⁺ ion to have a significantly more negative (favourable) hydration enthalpy. The large amount of energy released during the hydration of Cu²⁺ ions more than compensates for the energy cost, making Cu²⁺(aq) more stable than Cu⁺(aq) in aqueous solutions.
7. Do transition elements exhibit a zero oxidation state? If so, provide an example.
Yes, transition elements can exhibit a zero (0) oxidation state, especially in organometallic compounds called metal carbonyls. This occurs when they form bonds with ligands like carbon monoxide (CO). A classic example is Nickel Tetracarbonyl, [Ni(CO)₄], where the oxidation state of the nickel atom is zero.
8. Why is zinc (Zn) not considered a transition element based on its oxidation state?
Zinc is not considered a transition element because it does not have an incompletely filled d-subshell in its ground state (3d¹⁰4s²) or in its only common oxidation state, Zn²⁺ (3d¹⁰). A key characteristic of transition elements is a partially filled d-orbital, which is responsible for their variable oxidation states and other properties. Since zinc lacks this feature, it is classified as a d-block element but not a typical transition element.
9. How does the stability of higher oxidation states change down a group in the d-block?
In the d-block, heavier elements tend to have more stable higher oxidation states. For example, in Group 6, Molybdenum (Mo) and Tungsten (W) in the +6 state are much more stable than Chromium in the +6 state, which is a strong oxidising agent. This increased stability in heavier elements is because their valence electrons are further from the nucleus and more shielded, making them easier to involve in bonding.
