Question

Use the first-row transition metals $(Sc$ to $Cu)$ as an example to illustrate the characteristics of the electron configurations of transition metals.

Short answer

First-row transition metals exhibit unique electron configurations where 3d and 4s subshells are filled, leading to diverse chemical properties due to their d orbitals.

Step-by-step solution

Understand the Properties of Transition Metals

First-row transition metals are characterized by their partially filled d subshells. They exhibit unique properties like variable oxidation states, complex formation, and magnetic properties due to these electrons.

Identify the General Electron Configuration

Transition metals generally follow an electron configuration of $\text{{[noble gas]}}(n-1)d^{1-10}n{s}^{0-2}$ where n is the principal quantum number. For Scandium (Sc) to Copper (Cu), the inner filled shell will generally correspond to Argon (Ar) and the outer electrons will be filling 3d and 4s subshells.

Examine Specific Electron Configurations

Systematically examine each element from Sc to Cu:- Scandium (Sc): [Ar] 3d${}^1$ 4s${}^2$- Titanium (Ti): [Ar] 3d${}^2$ 4s${}^2$- Vanadium (V): [Ar] 3d${}^3$ 4s${}^2$- Chromium (Cr): [Ar] 3d${}^5$ 4s${}^1$, showing 3d stabilization through half-filled stability- Manganese (Mn): [Ar] 3d${}^5$ 4s${}^2$- Iron (Fe): [Ar] 3d${}^6$ 4s${}^2$- Cobalt (Co): [Ar] 3d${}^7$ 4s${}^2$- Nickel (Ni): [Ar] 3d${}^8$ 4s${}^2$- Copper (Cu): [Ar] 3d${}^{1}0$ 4s${}^1$, showing full d subshell stability.

Analyze the Trends in Electron Configurations

Note how unfilled or partially filled d orbitals lead to interesting chemical properties like the ability to form colored compounds and diverse oxidation states. Also, observe how copper (Cu) has a filled 3d shell with one electron in 4s, highlighting stability in full sub-shells and unique electron rearrangement for added stability.

Conclude on Transition Metal Characteristics

First-row transition metals exhibit variable oxidation states, form colored compounds, and display paramagnetism, primarily because of their partially filled d orbitals. They often have similar 4s and 3d electron configurations where changes in these configurations influence their chemical properties.

Key concepts

First-row Transition Metals

The first-row transition metals, spanning from Scandium (Sc) to Copper (Cu), are unique in the periodic table due to their ability to exhibit a range of distinct chemical properties. This row of elements is situated between Groups 2 and 13 and regarded as transition metals because their d subshell is progressively filled across the series. These elements serve as an exceptional example of the intermediate nature of transition metals, bridging between highly electropositive metals and less electronegative elements.
These metals showcase a diversity in their chemical behavior that is closely linked to their electron configurations. Their placement in the periodic table reflects a progressive filling of the 3d orbitals after the 4s orbital is occupied. Consequently, these elements have a rich chemistry characterized by complex formation, variability in oxidation states, and remarkable magnetic properties that set them apart from other elements. As we explore these elements, it is important to note the ever-changing interactions of the 3d and 4s electrons, which play a significant role in their wide-ranging functions and capabilities.

Partially Filled d Subshells

Partially filled d subshells are a hallmark of the transition metals and are at the core of their fascinating properties. For the first-row transition metals, these partially occupied d orbitals are what contribute to characteristics such as their ability to form complex ions and display a wide variety of colors in compounds.
Transition metals typically have the electron configuration $\text{{[noble gas]}}(n-1)d^{1-10}n{s}^{0-2}$, where n represents the principal quantum number. Throughout the series from Scandium to Copper, the increasing number of d electrons means that the electron configuration becomes more complex. For example, in Chromium (Cr), electrons are configured as $\text{{[Ar]}}3d^{5}4s^1$ showing a preference for half-filled stability. Such configurations highlight the energy stabilization achieved through specific arrangements of the d and s electrons.
These partially filled d subshells also allow for unique chemical bond formation. They can overlap with orbitals of other atoms to form covalent bonds, and the diverse arrangements of these electrons contribute to the structural variety in metal complexes.

Variable Oxidation States

One of the most intriguing attributes of transition metals is their ability to exhibit variable oxidation states. This property is primarily due to the presence of partially filled d orbitals, which allow these metals to lose different numbers of electrons. The first-row transition metals provide excellent examples of this phenomenon.
For instance, Vanadium (V) can demonstrate oxidation states ranging from +2 to +5. This variability results from the ability to lose both s and d electrons, leading to different degrees of metal oxidation. Each oxidation state corresponds to a unique electron configuration and thus leads to distinct chemical reactivity.
Variable oxidation states are crucial for many industrial processes and biochemical systems. For example, iron, in its +2 and +3 oxidation states, plays a vital role in biological systems, including oxygen transport and electron transfer._