Question

Why is the \(+2\) oxidation state common among the transition metals? Why do so many transition metals exhibit a variety of oxidation states?

Short answer

The +2 oxidation state is common among transition metals because removing two electrons from the outer s orbital leads to a relatively stable ion. They exhibit a variety of oxidation states due to partially filled d orbitals and relatively similar energy levels of the s and d orbitals, which allows for the removal or addition of electrons with varying degrees of energy, resulting in different oxidation states.

Step-by-step solution

Understanding electronic configuration of transition metals

Transition metals are the elements found in groups 3-12 of the periodic table. They have partially filled d orbitals, which give rise to their unique properties and behaviors, such as variable oxidation states. The electronic configuration of transition metals can be written in the form [noble gas] (n-1)d^x ns^y, where n represents the outermost energy level, x is the number of electrons in the d orbitals, and y is the number of electrons in the s orbitals.

Reason for common +2 oxidation state

The +2 oxidation state is commonly seen in transition metals because removing two electrons from the outermost s orbitals results in a relatively stable electronic configuration. Additionally, the removal of the two outermost s electrons doesn't require much energy, as they are relatively far from the nucleus and do not experience strong attractive forces. When these two electrons are removed, the remaining d-electrons are stabilized and contribute to the overall stability of the ion.

Transition metals and variable oxidation states

Transition metals exhibit a variety of oxidation states due to the presence of partially filled d orbitals and relatively similar energy levels of the s and d orbitals. This allows for the removal or addition of electrons with varying degrees of energy, resulting in different oxidation states. In addition to the +2 oxidation state, many transition metals can also lose electrons from their d orbitals with a relatively small increase in ionization energy. The difference in ionization energy between different oxidation states is often small, which makes it possible for transition metals to adopt multiple oxidation states. The incoming d-electrons can have different spins, creating further opportunities for varying oxidation states.

Example: Iron

For example, iron (Fe) has the electronic configuration [Ar] 3d^6 4s^2. In its +2 oxidation state, the electronic configuration becomes [Ar] 3d^6, which is achieved by removing the two 4s electrons. Iron exhibits another common oxidation state of +3, where one more electron is removed from the 3d orbitals, resulting in the configuration [Ar] 3d^5. The stability of both the +2 and +3 oxidation states for iron is due to the resulting electronic configurations, which are relatively stable. In conclusion, the +2 oxidation state is common among transition metals because removing two electrons from the outer s orbital results in a relatively stable ion. Transition metals exhibit a variety of oxidation states due to the presence of partially filled d orbitals and the relative closeness in energy levels of the s and d orbitals, which allows for the removal or addition of electrons with varying degrees of energy.