Electron Configurations of Atoms

Understanding electron configurations is fundamental to organic chemistry, as the arrangement of electrons dictates an atom's reactivity and bonding behavior. You'll recall these principles from general chemistry.

Determining Ground State Electron Configurations

To determine the ground-state electron configuration for an atom, we follow a systematic procedure:

1. Find the Atomic Number (Z): Locate the element on the periodic table and identify its atomic number, which equals the number of electrons in a neutral atom. For Carbon (C), Z = 6.
2. Apply the Aufbau Principle: Electrons fill atomic orbitals starting from the lowest energy level upwards. (Think of filling stairs from the bottom.) While the energy diagram typically shows 1s at the bottom, the filling order proceeds as 1s, 2s, 2p, 3s, 3p, etc.
   
3. Pauli Exclusion Principle and Hund's Rule:
   • Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons, and these electrons must have opposite spins.
   • Hund's Rule of Maximum Multiplicity: For degenerate orbitals (orbitals of the same energy, like the three p orbitals or five d orbitals), electrons will singly occupy each orbital with parallel spins before any orbital is doubly occupied.
     

Core vs. Valence Electrons

Understanding the distinction between core and valence electrons is crucial in organic chemistry:

• Valence Electrons: These are the electrons in the outermost shell (the highest principal energy level, n). They are the "glue" that binds atoms together in chemical bonds and are directly involved in chemical reactions.
  • For Carbon (1s2 2s2 2p2), the valence shell is n=2, so the valence configuration is 2s2 2p2. Carbon has 4 valence electrons.
• Core Electrons: These are the electrons in the inner shells, closer to the nucleus. They are generally unreactive and do not participate in bonding.
  • For Carbon, the electrons in the 1s2 orbital are the 2 core electrons.

It's vital to determine the number of valence electrons an atom has, as this directly relates to the number of bonds it typically forms and helps in calculating formal charges. You should be proficient at determining valence electrons for any atom, especially those in the second row of the periodic table.

Abbreviated (Noble Gas) Electron Configurations

For larger atoms, electron configurations can be abbreviated using the symbol of the preceding noble gas to represent the core electrons.

Procedure:

1. Find the noble gas that comes before the element in question on the periodic table.
2. Write the symbol for that noble gas in brackets.
3. Continue the electron configuration from that point, including only the valence electrons.

Ground vs. Excited States in Electron Configurations

While we often focus on the most stable arrangement of electrons, atoms can also exist in higher-energy states. Understanding the difference between ground and excited states is essential for predicting an atom's behavior, especially in spectroscopic contexts.

Ground State Electron Configuration

The ground state electron configuration represents the lowest energy, most stable arrangement of an atom's electrons. To determine the ground state:

• Follow the Aufbau Principle: Electrons fill orbitals from the lowest energy level up (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.).
• Adhere to the Pauli Exclusion Principle: A maximum of two electrons with opposite spins can occupy any single orbital.
• Apply Hund's Rule: For degenerate orbitals (orbitals of the same energy, like the three 2p orbitals), electrons will occupy each orbital singly with parallel spins before any orbital is doubly occupied.

Essentially, a ground state configuration means electrons are in their "proper" places, filling up the orbitals systematically and efficiently to achieve the lowest possible energy for the atom.

Excited State Electron Configuration

An excited state electron configuration occurs when one or more electrons in an atom absorb energy and jump to a higher energy orbital than they would occupy in the ground state. These states are less stable and temporary; excited electrons will eventually release the absorbed energy (often as light) and return to a lower energy state.

How to Identify an Excited State:

An electron configuration is in an excited state if it violates either the Aufbau Principle or Hund's Rule, while still maintaining the correct total number of electrons for the neutral atom. The Pauli Exclusion Principle (two electrons per orbital max, with opposite spins) is always followed, even in excited states.

Look for these common indicators of an excited state:

1. Electrons in Higher Orbitals Before Lower Ones are Full: If you see an electron in a higher energy orbital (e.g., a 4s electron) while a lower energy orbital (e.g., a 3d or 3p) that should be filled first is not, it's an excited state.
2. Violations of Hund's Rule: If degenerate orbitals are not singly occupied with parallel spins before pairing, or if electrons are paired up prematurely, it's an excited state (though sometimes this is less obvious in linear notation).

Practice Identifying States

For the following electron configurations, identify the elements (by writing their symbols, e.g., Ni, Ca, Se) and determine if they are in the ground state or excited state. Type either "ground" or "excited" in the appropriate cell.