Write ground-state electron configurations for these transition metal ions: (a) \(\mathrm{Sc}^{3+}\) (b) \(\mathrm{Ti}^{4+}\) (c) \(\mathrm{V}^{5+}\) (d) \(\mathrm{Cr}^{3+}\) , (e) \(\mathrm{Mn}^{2+},\) (f) \(\mathrm{Fe}^{2+}\) (h) \(\mathrm{Co}^{2+}\) (g) \(\mathrm{Fe}^{3+}\) (i) \(\mathrm{Ni}^{2+}\) (i) \(\mathrm{Cu}^{+}\) (k) \(\mathrm{Cu}^{2+}\) (1) \(\mathrm{Ag}^{+}\) \((\mathrm{m}) \mathrm{Au}^{+}\) (n) \(\mathrm{Au}^{3+},\) (o) \(\mathrm{Pt}^{2+}\)

Understanding the electron configuration of an atom or ion is essential to grasp its chemical behavior. Electron configuration refers to the arrangement of electrons in an atom's orbitals, which are defined by quantum mechanics. Electrons fill orbitals starting from the lowest energy level, moving towards higher levels, following the principles of the Aufbau rule, Hund's rule, and Pauli exclusion principle.

For instance, the ground-state electron configuration for a neutral scandium (Sc) atom, with an atomic number of 21, is expressed as \[\mathrm{[Ar]} 3d^1 4s^2\]. When scandium forms a \(\mathrm{Sc}^{3+}\) ion, it loses three electrons, typically the two 4s electrons and one 3d electron, leading to a configuration of \[\mathrm{[Ar]}\]. Similarly, this approach is used for other transition metals, considering their unique electron-distribution and the common practice of electrons being removed from the s-orbital before the d-orbital when forming ions.

In our exercise, by applying these principles, we confirmed the electron configurations for a series of transition metal ions, which serves as the crucial starting point for predicting their chemical properties.

Transition metal ions are formed when transition metals lose electrons and consequently achieve a positive charge. The electron configurations of these ions are unique because they involve the (n-1)d and ns orbitals. A notable feature of transition metals is the comparable energy levels of the 4s and 3d orbitals, which can lead to various electron configurations.

As transition metals ionize, electrons are typically lost first from the outermost s orbital followed by the d orbitals, which is counterintuitive since the d orbitals are filled earlier. Take titanium (Ti), for example, which loses four electrons to become \(\mathrm{Ti}^{4+}\) and ends with an electron configuration of \[\mathrm{[Ar]}\], despite having electrons in the 3d orbital in its neutral state. This pattern occurs because the energy of 4s electrons is higher when ions form, making them easier to remove. This is why understanding the sequence of orbital filling and the energy relationships between them is pivotal.

The periodic table orders elements by their atomic number and groups them into blocks according to their valence electron configurations. Transition elements belong to the d-block, which includes groups 3 to 12. These metals are characterized by having partially filled d subshells or are capable of forming stable cations with an incomplete d-subshell.

Transition metals offer a rich variety of chemistry due to their multiple oxidation states and ability to form complex ions. In the exercise provided, we examined diverse transition metals, including scandium (Sc), vanadium (V), and others, situated in various groups within the d-block. Each has distinctive electron configurations that play a crucial role in their chemistry, color, magnetic properties, and catalytic activities. A solid grasp of where these elements are located on the periodic table is vital for predicting their possible electron configurations and their subsequent chemical behavior.

The oxidation state is an indicator of the degree of oxidation (loss of electrons) of an atom in a chemical compound. Transition metals are known to exhibit a wide range of oxidation states due to the relatively low energy gap between their ns and (n-1)d electrons. This flexibility allows them to lose different numbers of electrons, leading to multiple stable ionic forms.

For example, in the case of chromium (Cr), it can have an oxidation state of +3 when it forms the \(\mathrm{Cr}^{3+}\) ion, as demonstrated in our exercise. This occurs as chromium loses 3 electrons — two from the 4s orbital and one from the 3d orbital — resulting in an electron configuration of \[\mathrm{[Ar]} 3d^3\].

Determining Oxidation States

The oxidation state can often be determined by the electron configuration and the known stability of the ion. It is a critical concept for understanding redox reactions, the formation of compounds, and the mechanisms of reaction pathways involving transition metals.