15 Aug, 2024

· Chemistry

What is the electron configuration of palladium (Pd)

Short Answer

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Long Explanation

Explanation

Palladium (Pd), with an atomic number of 46, exhibits an anomalous electron configuration that deviates from the expected pattern. Normally, we would anticipate that palladium would follow the straightforward Aufbau principle, filling up the electron shells in the order of increasing energy levels.

Expected Electron Configuration

Typically, the expected electron configuration of Pd would be:

[Kr] \, 4d^8 \, 5s^2

However, palladium's actual electron configuration is unique.

Actual Electron Configuration

The actual electron configuration of palladium (Pd) is:

[Kr] \, 4d^{10} \, 5s^0

In this configuration, palladium has 10 electrons in the 4d orbital and none in the 5s orbital.

Why This Anomaly Occurs

The deviation arises due to the stability imparted by a fully filled 4d subshell. The energy difference between the 4d and 5s orbitals in palladium is such that having a completely filled 4d orbital (with 10 electrons) and an empty 5s orbital results in a more stable configuration.

Summary of Key Points

Expected Configuration: $[Kr] \, 4d^8 \, 5s^2$
Actual Configuration: $[Kr] \, 4d^{10} \, 5s^0$
Reason for Anomaly: Enhanced stability due to the fully filled 4d subshell.

Understanding these anomalies is crucial in the field of chemistry, as they significantly affect the chemical properties and reactivity of the element.

Verified By

Johnathan Clark

Chemistry Content Writer at Math AI

Johnathan Clark, with a Master's in Chemistry from the University of São Paulo, is a young high school chemistry teacher and part-time contract writer. His engaging classroom experiments translate into compelling written content that makes chemistry exciting and practical.

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Concept

Anomalous Electron Configuration

Understanding Anomalous Electron Configuration

Anomalous electron configuration refers to instances where electrons in atoms do not follow the predicted electron configuration based on the usual order of energy levels. This often occurs due to subtle interactions that make certain arrangements more stable, even if they deviate from the expected pattern.

The Aufbau Principle and Expected Configurations

Generally, the electron configuration is predicted using the Aufbau principle, which states that electrons fill orbitals starting from the lowest available energy levels before filling higher levels. This often follows the order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, and so forth.

The expected configuration of an element like Chromium (Cr) with an atomic number of 24 would be:

24 \text{ e}^-: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^4

Examples of Anomalous Configurations

However, Chromium actually has an anomalous configuration:

24 \text{ e}^-: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5

Copper (Cu) is another element that shows this behavior. Instead of the expected configuration:

29 \text{ e}^-: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^9

It has an actual configuration of:

29 \text{ e}^-: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10}

Reasons for Anomalous Configurations

Increased Stability: Certain configurations are more stable due to exchange energy and symmetry. For instance, a half-filled or fully filled d-subshell offers additional stability.

Energy Sublevels: The energy difference between the 4s and 3d orbitals is relatively small. Electrons may occupy or vacate these orbitals to achieve a more stable state.

Illustrative Equations

For Chromium:

\begin{align*} \text{Expected:} & \quad 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^4 \\ \text{Actual:} & \quad 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5 \end{align*}

For Copper:

\begin{align*} \text{Expected:} & \quad 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^9 \\ \text{Actual:} & \quad 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10} \end{align*}

These configurations highlight the importance of considering stability factors beyond the basic Aufbau principle.

Concept

Stability Of Fully Filled Subshells

Stability of Fully Filled Subshells

The stability of fully filled subshells in atoms is a key concept in understanding atomic structure and electron configurations. The main idea revolves around the arrangement of electrons within an atom's orbitals and how this configuration impacts the atom's overall stability.

Quantum Mechanics and Electron Configuration

In quantum mechanics, electrons occupy specific energy levels known as orbitals. These orbitals are grouped into subshells denoted by the letters $s, p, d,$ and $f$ . The stability of an atom greatly depends on how these orbitals are filled.

Pauli Exclusion Principle and Hund's Rule

Pauli Exclusion Principle: This principle states that no two electrons in an atom can have the same set of quantum numbers. In simple terms, an orbital can hold a maximum of two electrons with opposite spins.
Hund's Rule: According to Hund's rule, electrons will first occupy different orbitals within the same subshell singly and with parallel spins before pairing up.

Increased Stability of Fully Filled Subshells

Atoms achieve enhanced stability when their subshells are either completely filled or half-filled. This increased stability is due to:

Symmetrical Distribution: Fully filled or half-filled subshells result in a more symmetrical and balanced distribution of electrons, leading to lower energy states.
Exchange Energy: Full and half-filled configurations maximize the exchange energy, which is a stabilizing factor associated with parallel electron spins.

For instance, the electron configuration of noble gases, such as neon ( $Ne$ ) and argon ( $Ar$ ), demonstrates full $p$ subshells:

Neon ( $Ne$ ): $1s^2\,2s^2\,2p^6$ Argon ( $Ar$ ): $1s^2\,2s^2\,2p^6\,3s^2\,3p^6$

These configurations show full $s$ and $p$ subshells, making these atoms extremely stable and chemically inert.

Mathematical Representation

The stability can be mathematically understood by considering the exchange integral $K$ , which represents the exchange energy due to symmetric wavefunctions of parallel spins. For instance, for a fully filled $p$ -subshell:

\Delta E_{exchange} = K \sum_{i<j} \langle \phi_i \phi_j | \frac{1}{r_{12}} | \phi_i \phi_j \rangle

Here, $\phi$ represents the wavefunctions, and $r_{12}$ is the distance between two electrons. Fully filled subshells maximize this stabilizing exchange interaction.

Real-World Examples

Chromium ( $Cr$ ): Despite violating the Aufbau principle, chromium has a $3d^5\,4s^1$ configuration instead of $3d^4\,4s^2$ . This is because a half-filled $3d$ -subshell provides extra stability.
Copper ( $Cu$ ): Copper exhibits a $3d^{10}\,4s^1$ configuration rather than $3d^9\,4s^2$ , favoring the fully filled $d$ -subshell for enhanced stability.

In summary, the stability of fully filled subshells is a fundamental concept that explains the preferred electron configurations of atoms and their resulting stabilities. This understanding is crucial for predicting and explaining various chemical properties and reactivities.