Electronic Energy Levels and Atomic Configurations Notes on Atomic Orbitals Summary

Electronic Energy Levels in Atoms

Atoms consist of electrons residing in specific energy levels, which are quantified by principal quantum numbers ( $n$ ).
Higher principal quantum numbers correspond to higher energy levels and greater distances from the nucleus. - For example, electrons in the 1st energy level (n=1) are significantly closer to the nucleus than those in the 2nd energy level (n=2) or higher levels.

Shape of Atomic Orbitals

Atomic orbitals denote the regions in space where there is a high probability of finding an electron, defined by their quantum numbers.   - s Orbitals:
    - Spherical in shape.
    - Increase in size with higher energy levels (n).   - p Orbitals:
    - Dumbbell-shaped and oriented along three axes: $p_x$ , $p_y$ , and $p_z$ .   - d Orbitals:
    - Characterized by more complex shapes with five distinct orientations, which can be visualized using diagrams.

Electronic Configuration of Atoms

Refers to the distribution of electrons among the orbitals of an atom.
Follows three key principles:
  - Aufbau Principle: Electrons fill orbitals starting from the lowest energy level.
  - Pauli Exclusion Principle: No two electrons in an atom can share the same set of quantum numbers (only two electrons can occupy a single orbital with opposite spins).
  - Hund’s Rule: Electrons will fill degenerate orbitals (orbitals of the same energy) singly before pairing to minimize repulsions.

Atomic and Ionic Radii

Atomic Radius: The distance from the nucleus to the outermost electron shell, typically measured in picometers (pm).
Ionic Radius: The size of an atom after it has gained or lost electrons to form an ion:
- Cations (positively charged ions) are generally smaller than their parent atoms.
- Anions (negatively charged ions) are usually larger than their parent atoms.

Ionization Energies

Definition: The energy required to remove an electron from an atom in the gas phase, measured in kilojoules per mole (kJ/mol).
Trends:
- Generally increases across a period (left to right) due to increased nuclear charge, which holds the electrons more tightly.
- Decreases down a group (top to bottom) as atomic size increases, making it easier to remove outer electrons.

Electron Affinity

Definition: The energy change when an electron is added to a neutral atom, forming a negative ion.
Generally, nonmetals exhibit higher electron affinities, leading to a greater release of energy upon gaining an electron.

Hund’s Rule

States that within a set of degenerate orbitals, one electron will fill each orbital singly with parallel spins before sharing occurs.
This arrangement maximizes the distance between electrons, reducing repulsion and enhancing the stability of the atom.

Pauli’s Exclusion Principle

Formulated by Wolfgang Pauli, this principle states that no two electrons in an atom can have identical sets of four quantum numbers.
Result: Each orbital can contain a maximum of two electrons, which must possess opposite spins.

Additional Information

Incorporating diagrams and images of atomic orbitals and electron configurations can enhance understanding. Consider including:
- Diagrams showing the shape of s, p, and d orbitals.
- Graphs illustrating trends in atomic and ionic radii and ionization energies across periods and groups.

Describe the electronic energy levels in atoms
Describe the shape of atomic orbitals
Understand the electronic configuration of atoms
Know what is meant by the terms: atomic and ionic radii, ionization energies, and electron affinity.
Understand Hund’s Rule and Pauli’s Exclusion Principle

Electrons reside in specific energy levels defined by principal quantum numbers ( $n$ ).
Higher values of $n$ indicate higher energy levels and greater distances from the nucleus (e.g., $n=1$ is closest, $n=2$ farther).

Each principal energy level can hold a maximum of $2n^2$ electrons, where $n$ is the principal quantum number.
For example,   - $n = 1$ can hold up to 2 electrons
  - $n = 2$ can hold up to 8 electrons
  - $n = 3$ can hold up to 18 electrons

Atomic orbitals represent regions with a high probability of finding an electron, characterized by quantum numbers.
Types of orbitals include: - s Orbitals: - Spherical
- Size increases with increasing $n$ - p Orbitals: - Dumbbell-shaped, aligned along $p_x$ , $p_y$ , $p_z$ axes. - d Orbitals: - More complex shapes with five orientations.

Each orbital type can hold different numbers of electrons: - Each s orbital holds a maximum of 2 electrons. - Each p orbital holds a maximum of 6 electrons (3 orbitals × 2 electrons each). - Each d orbital holds a maximum of 10 electrons (5 orbitals × 2 electrons each).

Distribution of electrons in an atom’s orbitals.
Principles: - Aufbau Principle: Fill lowest energy orbitals first. - Pauli Exclusion Principle: No two electrons can have the same quantum numbers (maximum of 2 in each orbital with opposite spins). - Hund’s Rule: Electrons fill degenerate orbitals singly (parallel spins) before pairing.

Notation for electron configuration includes the number of electrons in each orbital type (e.g., $1s^2 2s^2 2p^6$ ).

Atomic Radius: Distance from nucleus to outermost electron shell, typically measured in picometers (pm).
Ionic Radius: Size after gaining/loss of electrons: - Cations (positive ions) are smaller than neutral atoms. - Anions (negative ions) are larger than neutral atoms.

Energy needed to remove an electron from an atom in the gas phase, measured in kJ/mol.
Trends: - Increases across a period (left to right) due to higher nuclear charge. - Decreases down a group (top to bottom) due to increased atomic size.

The first ionization energy refers to removing the outermost electron. Higher ionization energies are observed for subsequent removals.

Energy change when adding an electron to a neutral atom, forming a negative ion.
Nonmetals typically release more energy (higher affinity) upon gaining electrons than metals.

In degenerate orbitals, distribute one electron per orbital with parallel spins before pairing.
Enhances stability by maximizing distance and minimizing repulsions.

No two electrons in one atom can have identical quantum numbers, ensuring each orbital can hold 2 electrons with opposite spins.

Quantum numbers include: - Principal quantum number (n): Indicates the energy level. - Azimuthal quantum number (l): Indicates orbital shape (s=0, p=1, d=2, f=3). - Magnetic quantum number (m_l): Indicates the orientation of the orbital. - Spin quantum number (m_s): Indicates the spin of the electron (+1/2 or -1/2).
Remember the relationship: $l$ can take values from 0 to $n-1$ .
For each $n$ , the value of $l$ determines the shape of the orbital, and for each $l$ , the value of $m_l$ determines the orientation of that shape.