Ionisation Enthalpy, Oxidation States — Definition

Definition

Imagine an atom as a tiny solar system, with electrons orbiting a central nucleus. To pull one of these electrons away from the atom, you need to supply energy, much like launching a rocket against Earth's gravity.

This energy is called Ionisation Enthalpy (IE). Specifically, the *first ionisation enthalpy* is the energy needed to remove the first electron from a neutral gaseous atom. If you then want to remove a second electron from the now positively charged ion, that requires the *second ionisation enthalpy*, and so on.

Each successive ionisation enthalpy is always higher than the previous one because you're trying to remove an electron from an increasingly positive ion, which holds onto its remaining electrons more tightly.

Now, let's talk about Oxidation States. When atoms form chemical bonds, they either gain, lose, or share electrons. The oxidation state (or oxidation number) is a way to keep track of these electron transfers.

It's a hypothetical charge an atom would have if all its bonds were considered purely ionic. For example, in NaCl, Na loses an electron to become $Na^+$ (oxidation state $+1$), and Cl gains an electron to become $Cl^-$ (oxidation state $-1$).

In molecules like water ($H_2O$), oxygen is more electronegative than hydrogen, so we assign oxygen an oxidation state of $-2$ and each hydrogen an oxidation state of $+1$, assuming the electrons in the O-H bonds are completely transferred to oxygen.

Transition elements, found in the d-block of the periodic table, are unique because they typically exhibit *variable oxidation states*. This means a single transition metal element can exist in compounds with different oxidation numbers.

For instance, iron can be found as $Fe^{2+}$ (oxidation state $+2$) or $Fe^{3+}$ (oxidation state $+3$). This variability arises because the energy difference between their outermost $ns$ electrons and the inner $(n-1)d$ electrons is very small.

This allows both sets of electrons to participate in chemical bonding, leading to a diverse range of possible oxidation states. This characteristic is fundamental to their chemical behavior, including their roles as catalysts and in forming colored compounds.