The Electron Configurations: Exceptions: Videos & Practice Problems

Understanding electron configurations is crucial in chemistry, and certain elements exhibit exceptions to the expected patterns due to their drive for stability. Electrons in p and d subshells are most stable when these orbitals are either half-filled or completely filled, which is related to the symmetry and stability of these configurations. Notably, chromium and copper are among the six elements that commonly demonstrate exceptions to their electron configurations. Chromium, for example, would ideally have a 4s² 3d⁴ configuration, but it actually promotes an electron from the 4s² orbital to the 3d⁵ to achieve a half-filled 3d⁵ configuration, resulting in a more stable 4s¹ 3d⁵ state. Similarly, copper adjusts its electron arrangement from an expected 4s² 3d⁹ to a more stable 4s¹ 3d¹⁰ configuration, completely filling the d subshell. This tendency to achieve half-filled or fully-filled d orbitals is a key factor in the exceptions observed in electron configurations of certain transition elements.

Electron Orbital Stability

Understanding electron configurations is crucial in chemistry, particularly when discussing exceptions that arise in certain elements. The stability of p and d subshells is a key factor in these exceptions. Specifically, p and d orbitals achieve maximum stability when they are either half-filled or fully filled with electrons. This stability is attributed to the symmetry of the electron arrangement.

According to Hund's rule, when electrons occupy degenerate orbitals (orbitals of the same energy), they will fill each orbital singly before pairing up. This means that for p subshells, which can hold a maximum of six electrons, the most stable configuration occurs when three electrons occupy the three available orbitals singly, resulting in a half-filled state. Similarly, for d subshells, which can hold up to ten electrons, a half-filled state occurs when five electrons are distributed across the five orbitals.

When these subshells are fully filled, they also exhibit stability. For instance, in a fully filled p subshell, all six electrons are paired, while in a fully filled d subshell, all ten electrons are paired. This drive for stability—whether through half-filling or full-filling of the p and d subshells—explains the deviations from expected electron configurations in certain elements.

Electron Configuration Exceptions

When studying electron configurations, it's important to recognize that exceptions begin to appear starting with chromium (Cr), which has an atomic number of 24. These exceptions can be observed in specific elements as the atomic number increases. Notably, the elements that exhibit these exceptions include chromium itself, as well as copper (Cu), which is found after skipping four elements: manganese (Mn), iron (Fe), cobalt (Co), and nickel (Ni).

In total, there are six key elements where these exceptions to the standard electron configuration occur. Understanding these exceptions is crucial for accurately determining the electron configurations of transition metals. The electron configuration for chromium, for instance, is typically represented as [Ar] 3d5 4s1 instead of the expected [Ar] 3d4 4s2. Similarly, copper's configuration is [Ar] 3d10 4s1 rather than [Ar] 3d9 4s2.

These deviations occur due to the stability associated with half-filled and fully filled subshells, which lower the energy of the atom. Recognizing these exceptions is essential for a deeper understanding of the behavior of transition metals in various chemical contexts.

d^4 Elements

Chromium is a notable example of an exception in electron configuration due to its unique tendency to achieve a more stable arrangement of electrons. Initially, the electron configuration of chromium can be represented as argon 4s₂3d₄. This configuration indicates that there are four electrons in the 3d subshell. However, both s and d subshells exhibit a preference for being either half-filled or fully filled, which leads to a modification in chromium's electron configuration.

To achieve a more stable configuration, one of the electrons from the 4s subshell can be promoted to the 3d subshell. This results in a new configuration of argon 4s₁3d₅. By promoting an electron, the 3d subshell becomes half-filled, which is energetically favorable. Thus, the correct electron configuration for chromium is argon 4s₁3d₅, highlighting the importance of achieving a half-filled set of d orbitals.

This phenomenon illustrates how electron configurations can deviate from expected patterns to enhance stability, particularly in transition metals like chromium. Understanding these exceptions is crucial for predicting the behavior and properties of elements in the periodic table.

d^9 Elements

Copper, when examined in the context of the periodic table, is initially represented as having the electron configuration of argon followed by 4s²3d⁹. However, a closer look at the 3d⁹ orbitals reveals that by adding just one more electron, the configuration can achieve a more stable state. This stability arises from the tendency of d and p sublevels to be either half-filled or fully filled.

In the case of copper, which is classified as a d⁹ element, the promotion of an electron from the 4s orbital to the 3d orbital occurs. This process transforms the electron configuration to 4s¹3d¹⁰, resulting in a completely filled d orbital. The driving force behind this electron promotion is the quest for stability through fully filled or half-filled d orbitals.

It is essential to recognize that this phenomenon is not unique to copper; there are several elements that exhibit similar behavior. These exceptions in electron configurations highlight the importance of understanding the underlying principles that govern electron arrangements in neutral elements, particularly the stability associated with filled and half-filled d orbitals.

Electron Configuration Exceptions Example

To determine the condensed electron configuration for a silver atom, we start by noting that silver has an atomic number of 47, indicating it has 47 electrons in its neutral state. The initial electron configuration can be derived from the periodic table, where silver is located in the transition metals section. The expected configuration would be based on the noble gas preceding silver, which is krypton (Kr), leading to an initial configuration of [Kr] 5s^2 4d^9.

However, silver is known for its electron configuration exception. In this case, to achieve a more stable configuration, one electron from the 5s orbital is promoted to the 4d orbital. This results in a fully filled 4d subshell. Therefore, the correct condensed electron configuration for silver is [Kr] 5s^1 4d^{10}.

This phenomenon is not unique to silver; it applies to several transition metals where the stability of half-filled or fully filled d orbitals is favored. Understanding these exceptions is crucial for accurately representing the electron configurations of certain elements.

Illustrate the exception to the electron configuration of molybdenum.

Molybdenum (Mo), with an atomic number of 42, has a unique electron configuration that deviates from the expected pattern. Typically, one would anticipate its electron configuration to follow the order of filling orbitals, leading to the configuration of Krypton (Kr) 5s² 4d⁴. However, molybdenum is an exception, similar to chromium, due to the stability associated with half-filled d orbitals.

To achieve a more stable electron configuration, one electron from the 5s orbital is promoted to the 4d orbital. This results in the actual electron configuration of molybdenum being Krypton 5s¹ 4d⁵. The promotion of the electron allows for a half-filled d subshell, which is energetically favorable.

In terms of the electron orbital diagram, the configuration can be illustrated as follows: the 5s subshell contains one electron, while the 4d subshell contains five electrons. According to Hund's rule, the five electrons in the 4d subshell will occupy separate orbitals before pairing up, resulting in the configuration:

5s: ↑
4d: ↑ ↑ ↑ ↑ ↑

This configuration highlights the importance of electron stability and the exceptions that can arise in the filling order of orbitals, particularly in transition metals like molybdenum.

Which of the following is the correct electron configuration of gold? 

• A.
• B.
• C.
• D.
• E.

Answer: (D)

A comparison of the electron configurations of palladium (Pd) and silver (Ag) indicates that: 

• A. Ag has 2 more d electrons and the same number of s electrons as Pd.
• B. Ag has 1 more d electron and the same number of s electrons as Pd.
• C. Ag has 2 more d electrons and 1 less s electron than Pd.
• D. Ag has 1 more d electron and 1 less s electron than Pd.
• E. Ag has 1 more d electron and 1 more s electron than Pd.

Answer: (C)

Here’s what students ask on this topic:

What are the exceptions to electron configuration?

Electron configuration refers to the distribution of electrons in an atom's orbitals. However, there are exceptions to the predicted order, which usually follow the Aufbau principle, Hund's rule, and the Pauli exclusion principle. These exceptions often occur due to the closeness in energy between the s and d orbitals or the p and d orbitals in the higher energy levels.

One common exception is seen in elements such as chromium and copper. In the case of chromium, instead of the expected electron configuration of [Ar] 4s² 3d⁴, it actually has a configuration of [Ar] 4s¹ 3d⁵. This is because half-filled d subshells are more stable due to electron repulsion and exchange energy. Similarly, copper is expected to have a configuration of [Ar] 4s² 3d⁹ but instead has [Ar] 4s¹ 3d¹⁰, as a completely filled d subshell is also more stable.

These exceptions generally occur because atoms seek the most stable electron configuration, which sometimes means deviating from the expected order to achieve a lower energy state. It's important to remember that the stability of electron configurations can be influenced by several factors, including electron-electron interactions and the specific energies of different orbitals.

Why are there exceptions to electron configuration?

Electron configurations typically follow a standard order based on the Aufbau principle, which states that electrons fill orbitals starting with the lowest energy levels before moving to higher ones. However, there are exceptions to these configurations primarily due to electron-electron interactions and the stability associated with half-filled and fully filled subshells.

In certain atoms, electrons will rearrange themselves to achieve a more stable electron configuration. This is often seen in the d and f blocks of the periodic table. For example, chromium (Cr) and copper (Cu) are two classic cases where we see exceptions. Chromium has the electron configuration [Ar] 3d⁵ 4s¹ instead of the expected [Ar] 3d⁴ 4s². This is because a half-filled d subshell (with five electrons) offers extra stability. Similarly, copper has the configuration [Ar] 3d¹⁰ 4s¹ as a fully filled d subshell (with ten electrons) is more stable than the expected [Ar] 3d⁹ 4s².

These exceptions occur due to the subtle balance between the energies of the orbitals and the specific electron configurations that confer extra stability to the atom.

What elements are exceptions in electron configuration?

In electron configuration, exceptions often occur due to the stability provided by half-filled or fully filled subshells. The most common exceptions are found among the transition metals, particularly in the chromium (Cr) and copper (Cu) families.

For chromium, which has an atomic number of 24, one would expect the electron configuration to be [Ar] 4s² 3d⁴. However, it is actually [Ar] 4s¹ 3d⁵. This is because having a half-filled d subshell (with five electrons) provides additional stability.

Similarly, copper, with an atomic number of 29, is expected to have the electron configuration [Ar] 4s² 3d⁹. Instead, its actual configuration is [Ar] 4s¹ 3d¹⁰. In this case, a completely filled d subshell (with ten electrons) is more stable than having a filled 4s and a partially filled 3d subshell.

These exceptions are due to subtle energy differences between the 4s and 3d subshells, and the stability that comes from symmetrical or fully filled electron arrangements. Other elements that show exceptions typically reside in the d-block and f-block of the periodic table, where similar stability considerations come into play.