Electronic Configurations

The content begins with an overview of the electronic configurations and discusses various topics related to it including the electromagnetic spectrum, energy levels, sublevels, and orbitals, and how to write electron configurations. This forms the foundational knowledge needed for further understanding in chemistry.

The electromagnetic spectrum encompasses all electromagnetic radiation, distinguished by their varying frequencies, wavelengths, and corresponding energy levels.
Frequency (f) is defined as the number of waves that pass a given point per second, while wavelength (λ) is the distance between consecutive peaks of a wave.
The spectrum is segmented into different bands, with gamma rays, X-rays, and ultraviolet (UV) radiation being on the high energy and high frequency end. These can be harmful to health.
All light waves travel at the same speed (the speed of light), which is crucial to understanding their properties.

The speed of light (symbol c) is constant at 3.00 x 10^8 m/s. The relationship among frequency, wavelength, and speed is captured in the equation: c = fλ; where an increase in frequency results in a decrease in wavelength.
The distinction between continuous spectra and line spectra is explained. A continuous spectrum contains all colors of the visible spectrum while a line spectrum shows only specific frequencies of light, indicating quantized energies emitted by atoms.
The concept of quanta is introduced where energy is released or absorbed in specified amounts rather than continuously.

Electrons can only have discrete energy values and any transition of electrons between energy levels results in the absorption or emission of energy.
It is noted that formulae relating frequency and wavelength are included in the IB Chemistry Data Booklet, easing the memorization burden for students.

Electrons orbiting the nucleus can absorb energy and transition to higher energy levels, after which they can revert to their original state, releasing energy in the process. This energy release corresponds to various specific frequencies of light, termed an emission spectrum.
The emission spectra can be qualitatively analyzed using diffraction gratings to yield line emission spectra.

Each line in the emission spectrum represents specific energy values, with electron transitions illustrating limited permissible energy states.
The concept of convergence is discussed, referring to the closeness of lines towards the higher energy end of the spectrum correlating to the ionization energy of the electron.
Historical context is provided with reference to Johannes Balmer, who first observed spectral lines.

The detailed depiction of electron jumps from various energy levels—namely infrared, visible, and ultraviolet—tied to specific energy transitions is provided for clarity.
Summarized data is presented in the form of a table to illustrate how electrons transition between quantized energy states.

A specific example detailing the electron transition emitting visible light is introduced, clarifying the concept further by providing multiple choice format with the correct transition from higher energy level down to n=2 yielding visible light.

Definitions of electron shells and how electrons are distributed around the nucleus is presented, emphasizing the role of principal quantum numbers in defining energy levels.
A formula for the number of electrons per principal energy level is also discussed, showing a pattern following the equation 2n².

The existence and significance of subshells is explained, including their energy hierarchy: s < p < d < f.
Subshells are further divided based on their specifics, showcasing patterns of energy distribution and the arrangement of electrons within them.

Orbitals, defined within subshells, are noted for having specific shapes and distinct energy levels, with capacity for two electrons at maximum per orbital.
The shapes of s and p orbitals are visually introduced.

Diagrams aid in understanding the 3D spatial orientation of orbitals, including s (spherical) and p (dumbbell-shaped) orbitals.
p orbitals extend along x, y, and z axes while s orbitals are uniformly spherical.

The concept of ground state, where the atom remains stable at its lowest energy configuration and fills subshells from lowest to highest energy levels (Aufbau Principle), is outlined.

The filling order of electrons across subshells according to the Aufbau Principle is depicted through diagrams elucidating how energies rank.

Discussion continues on how subshells (s, p, d, f) increase in energy with principal quantum number, clarifying exceptions like the 3d orbital.
Tables succinctly summarize orbitals per subshell and their respective capacities.

Maximum electrons per subshell detailed as 2 for s, 6 for p, 10 for d, and 14 for f. Highlights the degeneracy of orbitals within identical subshell.

A comprehensive table encapsulates the relationship between principal quantum numbers, possible subshells, and electron capacity in various shells presenting a holistic view of atomic structure.

An expansion on the shape of s and p orbitals emphasizing their respective geometries and how they vary with principal quantum numbers.

The significance of writing electron configurations is emphasized, detailing how it communicates the distinct arrangement of electrons in atoms.

The concept of electron spin is examined, detailing how this property leads to pairing in orbitals, alongside an introduction to Hund's Rule and Pauli Exclusion Principle.

Practical applications showing how to determine electron configurations in a systematic manner including full and shorthand versions, and also covering ions formed through electron loss or gain.

Examples like Chromium and Copper elucidate exceptions to common filling orders explaining why certain atoms promote electrons for stability in their d-subshells.

Helpful tips are offered for constructing effective orbital spin diagrams, emphasizing clarity in labeling and significance of arrow directions for denoting electron spins.