Problem 76 - Chemistry

Question

The ground-state electron configurations listed here are incorrect. Explain what mistakes have been made in each and write the correct electron configurations. Al: \(1 s^{2} 2 s^{2} 2 p^{4} 3 s^{2} 3 p^{3}\) \(\mathrm{B}: 1 s^{2} 2 s^{2} 2 p^{5}\) \(\mathrm{F}: 1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{6}\)

Step-by-Step Solution

Answer

The correct electron configurations are: Aluminium: \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{1}\), Boron: \(1 s^{2} 2 s^{2} 2 p^{1}\), Fluorine: \(1 s^{2} 2 s^{2} 2 p^{6}\).

Step 1: Correcting the Configuration for Aluminium

For Aluminium (Al), it's given as \(1 s^{2} 2 s^{2} 2 p^{4} 3 s^{2} 3 p^{3}\). However, if you check the periodic table, you'll find that Aluminium has an atomic number of 13, which means it has 13 electrons. The given configuration adds up to \(2+2+4+2+3=13\) electrons. Therefore, the initial count is correct but the placement of the electrons in \(2 p^{4}\) is not accurate. The configuration should be \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{1}\) following the rule that p-orbitals can at maximum hold 6 electrons.

Step 2: Correcting the Configuration for Boron

For Boron (B), the configuration provided is \(1 s^{2} 2 s^{2} 2 p^{5}\). Boron has an atomic number of 5 indicating 5 electrons. The given configuration adds up to 9 electrons which is incorrect. The correct electron configuration for Boron is \(1 s^{2} 2 s^{2} 2 p^{1}\).

Step 3: Correcting the Configuration for Fluorine

Lastly, the presented configuration for Fluorine(F) is \(1 s^{2} 2 s^{2} 2 p^{6}\). Fluorine has an atomic number of 9 and the given configuration satisfies this with a total of \(2+2+5=9\) electrons. Therefore, the configuration for Fluorine is correct and no changes are required.

Key Concepts

Ground-State Electron Configuration

Understanding the ground-state electron configuration is essential in chemistry. It refers to the arrangement of electrons around the nucleus of an atom at its lowest energy state. The electrons are organized in various orbitals depending on their energy levels. The main goal in determining these configurations is to minimize the energy of the atom, placing electrons in the lowest available energy levels first. This concept is guided by a set of rules including the Aufbau principle, which states that electrons fill orbitals starting from lowest to highest energy. Additionally, Hund’s rule ensures electrons fill degenerate orbitals singly first before pairing up. Therefore, a thorough knowledge of these guidelines helps in predicting the chemical behavior of elements effectively.

Atomic Number

The atomic number is a fundamental property of an element, signifying the number of protons in the nucleus of an atom. It is unique to each element and determines the identity of the element. In a neutral atom, the atomic number also equals the number of electrons, which are distributed among various orbitals. This distribution is critical when determining the electron configuration of an atom. For example, Aluminium has an atomic number of 13. Hence, its electron configuration needs to account for 13 electrons in its ground-state. Understanding atomic numbers is crucial for identifying mistakes in incorrect electron configurations and correcting them.

Periodic Table

The periodic table is an invaluable tool in chemistry, organizing elements based on their properties. It's arranged in increasing order of atomic numbers from left to right and top to bottom. Elements are categorized into periods (rows) and groups (columns) with similarities in chemical behavior. By examining its structure, one can predict the electron configuration of an atom. The periodic table helps determine the order of electron arrangement as s, p, d, and f blocks correspond to different types of orbitals. This layout simplifies the process of writing out the electron configuration of an element, ensuring electrons fill available orbitals in a systematic way. Moreover, elements in the same group have similar electron configurations in their outermost shells.

P-Orbitals

P-orbitals are a type of orbital in which electrons can be found. They have a distinctive dumbbell shape and are higher in energy compared to s-orbitals of the same shell. Each p subshell can hold up to six electrons, divided into three different p orbitals: p𝑥, p𝑦, and p𝑧. This plays a vital role when it comes to filling electrons in the configuration of elements, especially for those beyond the second period in the periodic table. For instance, Aluminium's electron configuration includes a 2p orbital which can hold up to six electrons. It's important to correctly fill the p orbitals to follow the rules of electron placement properly, such as the one that states no two electrons can have the same set of quantum numbers.

Electron Placement

Electron placement refers to the process by which electrons are distributed among the atomic orbitals. This placement follows several key principles.

• Aufbau Principle: Electrons fill the lowest energy orbitals first before moving to higher ones.
• Pauli Exclusion Principle: No two electrons can occupy the same quantum state simultaneously.
• Hund's Rule: Electrons fill degenerate orbitals (orbitals at the same energy level) singly before pairing up.

Ensuring correct electron placement is crucial for accurate electron configurations. Mistakes in placement can lead to incorrect electronic structures, impacting an atom's chemical properties. For instance, errors in placement as seen in the original exercise can lead to misconceptions about an element's reactivity and its potential to form bonds. Thus, accurate electron placement ensures a proper understanding of the atom's behavior in different chemical contexts.