KONFIGURASI ELEKTRON BERDASARKAN TEORI ATOM MEKANIKA KUANTUM

Summary

This video delves into electron configuration based on quantum mechanical theory. It explains how electrons occupy shells and subshells within an atom, detailing the different types of subshells (s, p, d, f) and their electron capacities. The video also introduces important rules like the Aufbau principle, Hund’s rule, and Pauli's exclusion principle, which guide electron distribution across orbitals. Several examples, including phosphorus, chromium, and copper atoms, illustrate how electron configurations are determined and why certain configurations are more stable. The video further discusses simplifying electron configurations using noble gas notation.

Takeaways

• 🔬 The video explains electron configuration based on quantum mechanics.
• ⚛️ Electrons in an atom occupy shells, which consist of subshells, and subshells contain orbitals.
• 🔢 For phosphorus (atomic number 15), its electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁵, with the last 3p indicating it has five electrons in the third shell's p subshell.
• 🌀 The Aufbau principle states that electrons fill lower energy subshells first before moving to higher energy subshells.
• 📉 Subshells fill in a specific energy order: 1s, 2s, 2p, 3s, 3p, 3d, 4s, and so on.
• 🎯 Hund's rule: electrons occupy orbitals singly first, then pair up, with opposite spins, when all orbitals in a subshell have one electron.
• 🚫 Pauli's exclusion principle: no two electrons can have the same set of four quantum numbers in the same orbital.
• 🔄 Half-filled and fully filled subshells are more stable, leading to electron shifts for stability, like in chromium or copper.
• ⚙️ Electron configurations can be abbreviated using noble gases, such as writing [Ne] for neon or [Ar] for argon.
• 🧪 Example configurations were provided for elements like carbon, neon, chlorine, calcium, and titanium, illustrating the rules.

Q & A

What is the significance of the electron configuration for phosphorus with an atomic number of 15?

The electron configuration of phosphorus (atomic number 15) is written as 1s2 2s2 2p6 3s2 3p5. This configuration shows that phosphorus has electrons in three shells, with the last shell containing five electrons in the 3p sublevel. The 3 indicates the third shell, p represents the sublevel, and 5 shows the number of electrons in this sublevel.

How are sublevels and orbitals structured within an electron configuration?

In electron configurations, sublevels (s, p, d, f) are part of a shell. Each sublevel contains orbitals, and each orbital can hold a maximum of two electrons. For example, the s sublevel has one orbital, the p sublevel has three orbitals, the d sublevel has five orbitals, and the f sublevel has seven orbitals.

What is the Aufbau principle in quantum mechanics, and how does it relate to electron configuration?

The Aufbau principle states that electrons occupy the lowest energy sublevels first before moving to higher ones. This helps determine the order in which electrons are filled in an atom's electron configuration, ensuring that the atom achieves the lowest possible energy state.

How can we determine the energy levels of sublevels using the Aufbau principle?

To determine the energy levels of sublevels, one can use the diagonal rule or energy diagram. This starts with the 1s sublevel, followed by 2s, 2p, 3s, 3p, and so on. Electrons fill the sublevels in increasing order of energy, ensuring the atom remains at its lowest energy configuration.

What does Hund's rule state about filling orbitals within a sublevel?

Hund's rule states that electrons will fill orbitals in a sublevel singly before pairing up. This means that if there are multiple orbitals within a sublevel (like the p sublevel), each orbital will receive one electron before any orbital gets a second electron, and these unpaired electrons will have parallel spins.

What is the Pauli Exclusion Principle, and how does it limit the arrangement of electrons in orbitals?

The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers. As a result, each orbital can hold a maximum of two electrons, and these electrons must have opposite spins.

What is meant by the 'half-filled and fully filled sublevel stability' rule?

According to the 'half-filled and fully filled sublevel stability' rule, atoms with half-filled or fully filled sublevels are more stable than those that are not. This leads to some exceptions in electron configurations, where electrons shift from one sublevel to another to achieve this more stable arrangement, as seen in elements like chromium and copper.

Why is the electron configuration for chromium unusual?

Chromium (atomic number 24) has an unusual electron configuration due to the half-filled and fully filled sublevel stability rule. Instead of following the expected configuration of 1s2 2s2 2p6 3s2 3p6 4s2 3d4, one electron from the 4s sublevel shifts to the 3d sublevel, resulting in a configuration of 1s2 2s2 2p6 3s2 3p6 4s1 3d5, making both the 4s and 3d sublevels more stable.

What is a noble gas configuration, and how is it used to simplify electron configurations?

A noble gas configuration is a shorthand way of writing electron configurations by using the electron configuration of the nearest noble gas as a starting point. For example, magnesium (atomic number 12) has the electron configuration [Ne] 3s2, where [Ne] represents the electron configuration of neon (1s2 2s2 2p6), simplifying the notation.

How does the electron configuration of copper illustrate the 'full and half-full sublevel stability' rule?

Copper (atomic number 29) normally would be expected to have the configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d9, but to achieve a more stable configuration, one electron from the 4s sublevel shifts to the 3d sublevel, resulting in a configuration of 1s2 2s2 2p6 3s2 3p6 4s1 3d10. This fully fills the 3d sublevel, increasing the atom's stability.