An Introduction
Transition elements are elements found on the periodic table in Groups 3-12 (old groups IIA-IIB) The term refers to the fact that the d sublevel that is being filled is at a lower principal energy level than the s sublevel that came before it. Scandium, the first transition element, has an electron configuration of [Ar]3d14s2. Remember that the configuration is the opposite of the fill order, with the 4 s filling before the 3 d begins. The transition elements are frequently referred to as transition metals because they are all metals. They exhibit typical metallic properties as a group and are less reactive than the metals in Groups 1 and 2.
Some of the more well-known ones are so unreactive that they can be found in nature in their free, or uncombined, form. Platinum, gold, and silver are examples. The transition elements are often referred to as "d -block" elements due to their unique filling order. Compounds containing a variety of transition elements are distinguished by their ability to be widely and vividly coloured. The d-orbitals absorb light of various energies as visible light passes through a transition metal compound dissolved in water. Visible light of a given energy level that is not absorbed results in a clear coloured solution.
Properties of Transition Elements
The transition elements' general properties are as follows:
They are typically metals with a high melting point.
They have a variety of oxidation states.
They usually combine to form coloured compounds.
They are frequently paramagnetic.
They have a high charge/radius ratio.
High density and hardness.
The boiling and melting points are both very high.
Construct paramagnetic compounds.
Variable oxidation states are displayed.
Coloured compounds and ions are common.
Create catalytically active compounds.
Create stable complexes
Oxidation States
The number of electrons that an atom loses, gains, or appears to use when joining with another atom in a compound is related to its oxidation state. It also determines an atom's ability to oxidise (lose electrons) or reduce (gain electrons) other atoms or species. Almost all transition metals have multiple oxidation states that have been experimentally observed. Ions are formed by adding or subtracting negative charges from an atom. Keeping the atomic orbitals in mind when assigning oxidation numbers aids in understanding that transition metals are a special case, but not an exception to this convenient method.
An atom with an oxidation number of -1 accepts an electron to achieve a more stable configuration. The electron donation is then +1. When a transition metal loses electrons, it usually loses s orbital electrons first, followed by d orbital electrons. See Formation of coordination complexes for a more detailed discussion of how these compounds form. Most transition metals have multiple oxidation states because transition metals lose electron(s) more easily than alkali metals and alkaline earth metals. The valence s-orbital of alkali metals contains one electron, and their ions almost always have oxidation states of +1. (from losing a single electron). Similarly, alkaline earth metals have two electrons in their valence s-orbitals, resulting in +2 oxidation state ions (from losing both). Transition metals, on the other hand, are more complex and exhibit a variety of observable oxidation states, owing primarily to the removal of d-orbital electrons.
