Question

Write electron configurations for each transition metal. (a) Zn (b) \(\mathrm{Cu}\) (c) \(\mathrm{Zr}\) (d) \(\mathrm{Fe}\)

Short answer

[Zn] [Ar] 3d10 4s2, [Cu] [Ar] 3d10 4s1, [Zr] [Kr] 4d2 5s2, [Fe] [Ar] 3d6 4s2

Step-by-step solution

Understanding Electron Configuration

Electron configuration refers to the distribution of electrons among the energy levels, sublevels, and orbitals of an atom. For transition metals, they generally follow the pattern of filling the 3d sublevel after filling the 4s sublevel, but there can be exceptions due to electron-electron interactions and stability.

Writing Electron Configuration for Zn

Zinc (Zn) has an atomic number of 30. The electron configuration can be written as follows: The 4s sublevel is filled before the 3d. All sublevels fill in the order of 1s, 2s, 2p, 3s, 3p, 4s, and then the 3d sublevel. For zinc, the electron configuration ends at 3d10, after filling the 4s2 sublevel first.

Writing Electron Configuration for Cu

Copper (Cu) has an atomic number of 29. The electron configuration for Cu is an exception to the general rule because it's more stable with a filled or half-filled d sublevel. Consequently, one electron from the 4s sublevel moves to the 3d sublevel, leading to the configuration [Ar] 3d10 4s1.

Writing Electron Configuration for Zr

Zirconium (Zr) has an atomic number of 40. The electron configuration follows: [Kr] 4d2 5s2. This is because we fill the 5s sublevel before starting to fill the 4d sublevel.

Writing Electron Configuration for Fe

Iron (Fe) has an atomic number of 26. The electron configuration is [Ar] 3d6 4s2. Iron follows the general rule where the 4s sublevel is filled with 2 electrons before continuing to fill the 3d sublevel.

Key concepts

Energy Levels and Orbitals

When exploring the world of atoms, the concepts of energy levels and orbitals are fundamental. Energy levels, also referred to as electron shells, are like concentric circles around the nucleus of an atom, with each successive level further away from the nucleus holding more energy. Within these energy levels are orbitals, which are specific regions where the probability of finding an electron is the highest.

An easy way to visualize orbitals is as different-shaped 'clouds' within an energy level: spherical for 's' orbitals, dumbbell-shaped for 'p' orbitals, and increasingly complex for 'd' and 'f' orbitals. It's important to remember that each orbital can hold a maximum of two electrons. Understanding how electrons fill these orbitals is crucial for grasping concepts such as chemical bonding and atomic properties.

Electron Sublevel Filling Order

The journey of an electron through the complex road map of an atom follows specific rules. Among them, the electron sublevel filling order is a guide for understanding how electrons populate the orbitals. Sublevels are denoted by the letters s, p, d, and f, and they fill up in a particular sequence based on the increasing energy.

The rule of thumb is called the Aufbau principle, which states that electrons occupy the lowest-energy orbital available. Following this, the order generally begins with 1s, then 2s, followed by 2p, 3s, 3p, and so forth. Visual aids like the Aufbau diagram can help students remember this sequence more easily. In identifying how electrons populate these sublevels, using the order as a checklist can ensure that you don't miss any steps when writing electron configurations.

Atomic Number and Electron Distribution

The atomic number of an element—often shown as 'Z'—is a pivotal piece of information, as it not only tells us the number of protons in an atom's nucleus but also reveals the total number of electrons when the atom is neutral. A balanced understanding of the atomic number and electron distribution is vital to write accurate electron configurations.

To populate the orbitals with electrons, simply match the total number of electrons (which equals the atomic number) with the available spots in the energy sublevels based on the filling order. For transition metals, this means looking closely at the 3d and 4s sublevels and knowing when to place electrons in which. This is because transition metals often display behaviors that buck the standard energy order, challenging us to think critically about the electron distribution.

Exceptional Electron Configurations

Among all the transition metals, certain atoms stand out with their exceptional electron configurations. Elements like chromium and copper do not follow the regular filling order, due to their quest for greater stability. Electrons in these cases are rearranged, resulting in configurations that may initially appear perplexing but make sense when considering principles of atomic stability.

For instance, electrons prefer to fill orbitals in a way that minimizes electron-electron repulsion, often leading to half-filled or fully-filled d-sublevels. These anomalies in electron configurations are exceptions to the otherwise systematic filling order, emphasizing the need for students to engage with the content on a case-by-case basis. Knowing these exceptions helps deepen understanding of electron configurations and prepares students for more complex concepts in chemistry.