The Electron Configuration: Videos & Practice Problems

Electron Orbital Diagrams

The periodic law plays a crucial role in determining the electron arrangements of elements, which can be visually represented through electron orbital diagrams. These diagrams illustrate how electrons are distributed among various orbitals, particularly focusing on the concept of degenerate orbitals—sets of orbitals that have the same energy level. According to Hund's rule, these degenerate orbitals are filled in a specific manner: they are first half-filled before any orbital is completely filled.

To understand the filling of these orbitals, we can examine the different subshells:

1. **s Sublevel**: The s subshell contains 1 orbital and can hold a maximum of 2 electrons. In this orbital, one electron spins up while the other spins down, adhering to the Pauli exclusion principle.

2. **p Sublevel**: The p subshell consists of 3 orbitals, allowing for a maximum of 6 electrons. Following Hund's rule, each of the three orbitals is first half-filled with one electron (all spinning up) before pairing occurs, resulting in the configuration of up, up, up, down, down, down.

3. **d Sublevel**: The d subshell has 5 orbitals, accommodating a total of 10 electrons. Similar to the p subshell, the d orbitals are half-filled first, with one electron in each orbital before pairing begins, leading to a filling sequence of up, up, up, up, up, down, down, down, down, down.

4. **f Sublevel**: The f subshell contains 7 orbitals and can hold a maximum of 14 electrons. Again, following Hund's rule, the orbitals are half-filled first before any pairing occurs, resulting in a filling pattern that allows for maximum stability.

In summary, the maximum electron capacities for each subshell are as follows: the s subshell can hold 2 electrons, the p subshell can hold 6, the d subshell can hold 10, and the f subshell can hold 14. Understanding these principles is essential for grasping the electron configuration of elements and their chemical behavior.

Electron Orbital Diagrams Example

When filling the d set of orbitals for an atom with 8 electrons, it's essential to understand the structure and energy of these orbitals. The d set consists of 5 orbitals, which are degenerate, meaning they have the same energy level. According to Hund's rule, electrons will occupy each orbital singly before pairing up. This means that for the first 5 electrons, you will place one electron in each of the 5 d orbitals, represented as:

↑ ↑ ↑ ↑ ↑

After filling the first 5 orbitals, you have 3 electrons remaining to place. You will then return to the first orbital and begin pairing the electrons. The next 3 electrons will be placed as follows:

↓ ↓ ↓

Thus, the final electron configuration for the d orbitals with 8 electrons will have 5 orbitals filled with one electron each and 3 orbitals with a pair of electrons, resulting in:

↑↓ ↑↓ ↑↓ ↑ ↑

This configuration illustrates that the first 5 d orbitals are half-filled, and the remaining 3 electrons are paired in the first three orbitals. This method of filling ensures that the electrons are distributed in a way that minimizes repulsion and maximizes stability within the atom.

Which electron configuration represents a violation of Hund's Rule? 

• A.
• B.
• C.
• D.

Answer: (D)

Periodic Table Blocks

The periodic table can be visualized in a way that highlights its structure through distinct color-coded blocks: the s block (blue), p block (yellow), d block (purple), and f block (red). Each block corresponds to specific electron sublevels, which are crucial for understanding electron configurations.

The s sublevel contains 1 orbital, allowing it to hold a maximum of 2 electrons. This is reflected in the two columns of the s block. In contrast, the p sublevel has 3 orbitals, accommodating up to 6 electrons, which is why the p block consists of 6 columns. The d sublevel, with 5 orbitals, can hold up to 10 electrons, corresponding to the 10 positions in the d block. Lastly, the f sublevel can contain 14 electrons due to its 7 orbitals, represented by the 14 slots in the f block.

Starting with hydrogen, the electron configuration begins as \(1s^1\). As we progress through the periodic table, we add electrons sequentially. For example, helium is \(1s^2\), and in the second row, we continue with \(2s^1\) and \(2s^2\), followed by \(2p^1\) through \(2p^6\). This pattern continues with the s and p blocks in subsequent rows, leading to configurations like \(3s^2\), \(4s^2\), and \(5s^2\), and so forth.

When examining the d block, it is important to note that the principal quantum number decreases by one. For instance, the configuration for the d block starts at \(4s\) but is labeled \(3d\), continuing with \(4d\) and \(5d\). The f block follows a similar pattern, with configurations like \(4f\) and \(5f\) corresponding to the elements between atomic numbers 57 to 72 and 89 to 104, respectively.

By understanding this reimagined periodic table, one can effectively determine the electron configuration for any element or ion. This foundational knowledge is essential for further studies in chemistry, particularly in areas such as bonding, reactivity, and the behavior of elements in various chemical contexts.

Electron Configuration Example

To determine the ground state electron configuration of fluorine, which has an atomic number of 9, we start by recognizing that this means fluorine has 9 electrons. The electron configuration describes how these electrons are distributed among the various atomic orbitals.

We begin with the lowest energy level, which is the 1s orbital. The filling order of the orbitals follows the Aufbau principle, where electrons occupy the lowest energy orbitals first. For fluorine, the electron configuration proceeds as follows:

1. The 1s orbital can hold a maximum of 2 electrons, so we write 1s₂.

2. Next, we fill the 2s orbital, which also holds 2 electrons: 2s₂.

3. Finally, we move to the 2p orbital. Fluorine is located in the 2p block and has 5 electrons in this orbital, which we denote as 2p₅.

Putting this all together, the complete ground state electron configuration for fluorine is:

1s₂ 2s₂ 2p₅.

When we add the electrons from each orbital (2 from 1s, 2 from 2s, and 5 from 2p), we confirm that the total is 9 electrons, consistent with the atomic number of fluorine. This relationship highlights that in a neutral atom, the number of electrons equals the number of protons, which is defined by the atomic number.

Utilizing the periodic table effectively can guide you in determining the electron configurations for any element or ion, ensuring a clear understanding of their electronic structure.

Which electron configuration represents a violation of the Auf Bau Principle? 

• A.
• B.
• C.
• D.

Answer: (B)

Identify the element with the given electron orbital diagram. 

• A. Silicon
• B. Fluorine
• C. Sulfur
• D. Chlorine
• E. Phosphorus

Answer: (C)

Write the electron configuration and electron orbital diagram for the following element: 
Mn (Z = 25) 

• A. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵
• B. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 4d⁵
• C. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴
• D. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵

Answer: (D)

Write the ground-state electron configuration for the following element: 
Lead (Z = 82) 

• A. $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{10}5d^{10}6p^6$
• B. $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}6p^2$
• C. $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}5d^66p^6$
• D. $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}5p^66s^24f^{14}5d^{10}6p^2$

Answer: (D)

Here's what students ask on this topic:

What is the periodic law and how does it relate to electron configurations?

The periodic law states that when elements are arranged by increasing atomic weight, elements in the same group share similar chemical properties. This law is fundamental because these chemical properties influence the electron arrangements or configurations of the elements. Electron configurations describe how electrons are distributed in an atom's orbitals, which directly affects the element's reactivity and atomic structure. Understanding the periodic law helps us predict and explain trends in electron configurations across the periodic table, as elements in the same group tend to have similar valence electron arrangements, leading to similar chemical behavior.

How do Hund's rule and degenerate orbitals affect electron configuration?

Hund's rule states that electrons fill degenerate orbitals—orbitals of the same energy—by first occupying each orbital singly with parallel spins before pairing up. This minimizes electron-electron repulsion and stabilizes the atom. For example, in the p sublevel, which has three orbitals, electrons will half-fill each orbital (one electron per orbital) before any orbital gets a second electron. This rule is crucial for correctly writing electron configurations and understanding the magnetic properties of atoms. It ensures that electrons are distributed to maximize total spin, affecting the atom's energy and chemical properties.

What is the Aufbau principle and how is it used to write electron configurations?

The Aufbau principle guides the order in which electrons fill atomic orbitals. It states that electrons occupy the lowest energy orbitals first before filling higher energy ones. Starting from the 1s orbital, electrons fill orbitals in a sequence determined by their energy levels, such as 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on. This principle helps us write the ground state electron configuration of elements by following the energy order, often visualized using the Aufbau diagram. It ensures that the electron configuration reflects the most stable, lowest-energy arrangement of electrons in an atom.

How are the s, p, d, and f blocks arranged in the periodic table and what do they represent?

The periodic table is divided into blocks based on the electron sublevels being filled: s, p, d, and f blocks. The s block (blue) includes two columns and corresponds to the filling of s orbitals, which hold a maximum of 2 electrons. The p block (yellow) has six columns, representing the filling of p orbitals with up to 6 electrons. The d block (purple) contains ten columns for d orbitals, which can hold up to 10 electrons. Lastly, the f block (red) includes 14 columns for f orbitals, holding up to 14 electrons. These blocks help visualize electron configurations and predict element properties by showing which sublevel is being filled as you move across the table.

What are the maximum number of electrons that can be held in s, p, d, and f sublevels?

The maximum number of electrons in each sublevel is determined by the number of orbitals and the Pauli exclusion principle, which allows two electrons per orbital with opposite spins. The s sublevel has 1 orbital, so it can hold up to $2_{}$ electrons. The p sublevel has 3 orbitals, holding up to $6$ electrons. The d sublevel has 5 orbitals, accommodating up to $10$ electrons. The f sublevel has 7 orbitals, allowing up to $14$ electrons. These limits are essential for writing accurate electron configurations and understanding the structure of atoms.