6.5 Electron Configuration | General Chemistry

Summary

This chemistry lesson explores electron configurations in atoms, detailing the Aufbau principle, Hund's rule, and the Pauli exclusion principle. It covers standard configurations, noble gas notations, and exceptions like those for copper, silver, and gold. The instructor, Chad, also discusses ions, valence electrons, and the concept of excited states versus ground states, providing a comprehensive guide to electron configurations for students preparing for exams.

Takeaways

• 📚 The lesson covers the rules for electron configuration in atoms, including the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
• 🔋 The Aufbau principle states that electrons fill the lowest energy orbitals first, working their way up in energy levels.
• 👥 The Pauli exclusion principle dictates that no two electrons in an atom can have the same set of four quantum numbers, meaning they can't occupy the same orbital with the same spin.
• 🧲 Hund's rule explains the filling of degenerate orbitals, where each orbital gets one electron before any pairing occurs, and all unpaired electrons have the same spin.
• 📝 Standard electron configurations are discussed, with examples and notable exceptions, emphasizing the importance of understanding the order of filling orbitals.
• 🏫 The instructor introduces the concept of abbreviated noble gas configurations as a shorthand for electron configurations, simplifying the notation for atoms after a certain point in the periodic table.
• 🌟 The lesson touches on ions, explaining how to determine electron configurations for cations and anions, and the special considerations for transition metal cations.
• 💡 The number of valence electrons is related to an atom or ion's electron configuration, with the periodic table providing a guide to determine the number of valence electrons for main group elements.
• 🚫 The script mentions five notable exceptions to the standard electron configuration rules: copper, silver, gold, chromium, and molybdenum, which do not follow the typical filling order due to increased stability with half-filled or fully-filled subshells.
• 🔄 The concept of an excited state is introduced, where an electron is promoted to a higher energy orbital, contrasting with the ground state which follows the Aufbau principle.
• 🔗 The instructor emphasizes the importance of understanding electron configurations for various chemical applications and how the periodic table can be used as a tool to determine the order of orbital filling.

Q & A

What are the three main principles discussed in the lesson for filling electrons in an atom?

The three main principles discussed are the Aufbau Principle, Hund's Rule, and the Pauli Exclusion Principle. The Aufbau Principle states that electrons fill in the lowest energy orbitals first. Hund's Rule says that equal energy orbitals are each filled with one electron before any pairing occurs, and all unpaired electrons in these orbitals must have the same spin. The Pauli Exclusion Principle states that no two electrons in an atom can have the same four quantum numbers, meaning no two electrons can be in the same orbital with the same spin.

What is the significance of the term 'degenerate' in the context of orbitals?

In the context of orbitals, 'degenerate' refers to orbitals that have the same energy level. For example, all three 2p orbitals are degenerate and can each hold up to two electrons with opposite spins. The order in which these degenerate orbitals are filled does not matter, but Hund's Rule dictates that they should be filled with one electron each before any pairing occurs.

How does the energy level of subshells change in a multi-electron system?

In a multi-electron system, the energy levels of subshells within the same shell can change due to electron-electron repulsion. This causes the subshells to split into different energy levels, with the 2p subshell being higher in energy than the 2s subshell, for instance. The order of energy levels for subshells in larger shells can be different, such as 4s being filled before 3d.

What is the purpose of the periodic table in determining electron configurations?

The periodic table provides a guide to the order in which electrons are filled in orbitals. Elements in the periodic table are arranged in increasing atomic number, which corresponds to the order in which electrons are added to an atom. The periodic table also helps in identifying the subshells that are filled last, which can be useful in writing electron configurations, especially for ions or excited states.

How is the electron configuration of an ion different from that of its neutral atom?

The electron configuration of an ion differs from its neutral atom because ions have gained or lost electrons. Cations, which are positively charged ions, have lost electrons and thus have fewer electrons in their outermost shell compared to the neutral atom. Anions, which are negatively charged ions, have gained electrons and have additional electrons in their outermost shell. The number of electrons lost or gained determines the charge of the ion.

What is the noble gas configuration and why is it used?

The noble gas configuration is an abbreviated way of writing electron configurations by using the electron configuration of the noble gas that precedes the element in the periodic table. It is used to simplify the notation, especially for elements with higher atomic numbers, by indicating the electron configuration up to the last noble gas core and then adding the remaining valence electrons.

Why does the electron configuration of copper (Cu) deviate from the expected configuration based on the Aufbau Principle and Hund's Rule?

The expected electron configuration for copper based on the Aufbau Principle and Hund's Rule would be [Ar] 4s2 3d9. However, copper's actual ground state configuration is [Ar] 4s1 3d10. This deviation occurs because there is extra stability when a subshell is either completely filled or half-filled, which is the case for copper's 3d subshell.

What are the five exceptions to the standard electron configurations that are commonly taught?

The five exceptions to the standard electron configurations that are commonly taught are copper (Cu), silver (Ag), gold (Au), chromium (Cr), and molybdenum (Mo). These elements have special stability when their d or f subshells are half-filled or completely filled, leading to electron configurations that deviate from the expected patterns.

How do you determine the number of valence electrons in an atom?

The number of valence electrons in an atom can be determined by looking at the group in the periodic table that the element belongs to for main group elements. For transition metals, the number of valence electrons is typically the number of electrons in the outermost s and d subshells. However, there are exceptions, and the actual number can depend on the chemical behavior of the element.

What is an excited state in the context of electron configurations?

An excited state refers to a condition where an electron in an atom has absorbed energy and moved to a higher energy orbital, resulting in a configuration that violates the Aufbau Principle. This is a temporary state, and the atom will eventually return to its ground state by releasing the energy.