Of The Following Which Sublevel Is Filled Last

Of the Following, Which Sublevel is Filled Last? Understanding Electron Configuration and the Aufbau Principle

Determining which sublevel is filled last in an electron configuration is crucial for understanding the properties of elements and their behavior in chemical reactions. This seemingly simple question delves into the fascinating world of quantum mechanics and the arrangement of electrons within atoms. This comprehensive guide will explore the Aufbau principle, Hund's rule, and the exceptions to these rules, ultimately enabling you to confidently predict the last filled sublevel for any element.

Understanding Electron Configuration

Before we delve into which sublevel is filled last, let's establish a firm understanding of electron configuration. An electron configuration describes the arrangement of electrons in an atom's orbitals. These orbitals are regions of space around the nucleus where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons, according to the Pauli Exclusion Principle.

Orbitals are grouped into subshells, which are in turn grouped into shells. The principal quantum number (n) designates the shell (n=1, 2, 3, etc.), while the azimuthal quantum number (l) specifies the subshell within a shell. The subshells are labeled s, p, d, and f, corresponding to l values of 0, 1, 2, and 3 respectively. Each subshell holds a specific number of orbitals and therefore a specific number of electrons:

• s subshell: One orbital, holding a maximum of 2 electrons.
• p subshell: Three orbitals, holding a maximum of 6 electrons.
• d subshell: Five orbitals, holding a maximum of 10 electrons.
• f subshell: Seven orbitals, holding a maximum of 14 electrons.

The Aufbau Principle: Filling Orbitals in Order of Increasing Energy

The Aufbau principle, which translates from German to "building-up principle," dictates the order in which electrons fill the atomic orbitals. Electrons fill the orbitals starting with the lowest energy level and progressing to higher energy levels. However, the energy levels aren't always straightforward; the energy of an orbital depends on both the principal quantum number (n) and the azimuthal quantum number (l).

The general order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p...

This order might seem arbitrary, but it reflects the increasing energy of the orbitals. Remember, the energy of an orbital isn't solely determined by its principal quantum number (shell). For instance, the 4s orbital is lower in energy than the 3d orbital, hence the 4s orbital is filled before the 3d orbital.

Visual Aids: Understanding Orbital Energy Levels

Several diagrams can help visualize the order of orbital filling. One popular method is the diagonal rule (also known as the Madelung rule), where you draw diagonal lines across the subshells, following the order of filling. Another helpful visual is an energy level diagram, which explicitly shows the relative energies of different orbitals. These diagrams, readily available online and in chemistry textbooks, provide a clear picture of the Aufbau principle in action.

Hund's Rule: Filling Orbitals Within a Subshell

Once we know the order of filling from the Aufbau principle, we need to consider Hund's rule. Hund's rule states that electrons will individually occupy each orbital within a subshell before doubling up in any one orbital. This is because electrons, being negatively charged, repel each other. They prefer to occupy separate orbitals within a subshell to minimize repulsion, resulting in a more stable configuration.

Exceptions to the Aufbau Principle

While the Aufbau principle and Hund's rule provide a reliable framework for predicting electron configurations, there are exceptions. These exceptions arise because the energy differences between orbitals can be quite small, and interelectronic repulsions can sometimes lead to a more stable configuration by violating the Aufbau principle. Some common examples include:

• Chromium (Cr): Instead of the expected configuration [Ar] 3d⁴ 4s², chromium has the configuration [Ar] 3d⁵ 4s¹. This is because a half-filled d subshell (5 electrons) and a half-filled s subshell (1 electron) provide greater stability due to enhanced electron exchange energy.

• Copper (Cu): Similar to chromium, copper deviates from the expected [Ar] 3d⁹ 4s² to [Ar] 3d¹⁰ 4s¹. The completely filled d subshell (10 electrons) contributes to enhanced stability.

• Other Transition Metals and Lanthanides/Actinides: Other transition metals and elements in the lanthanide and actinide series also show exceptions to the Aufbau principle due to the complex interplay of electron-electron interactions and relativistic effects.

Determining the Last Filled Sublevel: A Step-by-Step Approach

Now, equipped with the knowledge of the Aufbau principle, Hund's rule, and the exceptions, let's outline a step-by-step process to determine the last filled sublevel for any given element:

1. Identify the Atomic Number: Find the atomic number (Z) of the element from the periodic table. The atomic number corresponds to the number of electrons in a neutral atom.

2. Apply the Aufbau Principle: Use the Aufbau principle or a visual aid like the diagonal rule to determine the order of orbital filling.

3. Fill Orbitals According to Hund's Rule: Follow Hund's rule to fill orbitals individually within a subshell before pairing up electrons.

4. Consider Exceptions: Be mindful of the exceptions to the Aufbau principle, particularly for transition metals and lanthanides/actinides. Consult a reliable resource like a periodic table with electron configurations or a chemistry textbook if you encounter an element known for exceptions.

5. Identify the Last Filled Sublevel: The sublevel containing the last electron added is the last filled sublevel.

Examples: Determining the Last Filled Sublevel

Let's illustrate this process with a few examples:

• Oxygen (O, Z=8): Following the Aufbau principle, the electron configuration is 1s² 2s² 2p⁴. The last filled sublevel is 2p.

• Titanium (Ti, Z=22): The electron configuration is [Ar] 3d² 4s². The last filled sublevel is 4s.

• Iron (Fe, Z=26): The electron configuration is [Ar] 3d⁶ 4s². The last filled sublevel is 4s. Note that even though the 3d subshell is almost filled, 4s is filled last.

• Gold (Au, Z=79): This element shows an exception. While the expected configuration might suggest otherwise, the actual configuration is [Xe] 4f¹⁴ 5d¹⁰ 6s¹. The last filled sublevel is 6s.

Conclusion: Mastering Electron Configuration for Deeper Understanding

Determining the last filled sublevel is a fundamental concept in chemistry, crucial for understanding an element's chemical properties and reactivity. While the Aufbau principle provides a solid foundation, remembering Hund's rule and being aware of the exceptions are essential for accurate predictions. By mastering these principles, you can confidently navigate the fascinating world of electron configurations and gain a deeper appreciation for the intricate behavior of atoms. The ability to determine the last filled sublevel is a key skill for success in advanced chemistry studies. Practice with various elements and consult resources as needed to solidify your understanding.