Atomic Physics and Electron Transitions

When an electron moves from a lower to a higher energy level, it absorbs energy, gains potential energy, becomes excited, and moves farther away from the nucleus.

The specific wavelengths of light emitted when electrons in an atom transition from excited states back to their ground states.

The energy difference between two energy levels in an atom.

The process of exciting atoms with energy and then observing the light they emit as electrons return to their ground states.

The lowest energy state of an atom, where its electrons occupy their normal, stable orbitals.

A higher energy state of an atom where one or more electrons have absorbed energy and moved to a higher energy level.

The phenomenon where an excited substance emits light, often giving off a specific color.

The set of specific, measurable, wavelengths of light emitted when electrons transition from higher energy levels to lower energy levels in an atom. Each element has a unique emission spectrum.

A series of spectral lines in the hydrogen emission spectrum that correspond to transitions from higher energy levels (n≥2) to the ground state (n=1). These lines fall in the ultraviolet region of the electromagnetic spectrum.

A series of spectral lines in the hydrogen emission spectrum that correspond to transitions from higher energy levels (n≥3) to the second energy level (n=2). Some of these lines fall within the visible region of the electromagnetic spectrum.

A series of spectral lines in the hydrogen emission spectrum that correspond to transitions from higher energy levels (n≥4) to the third energy level (n=3). These lines fall in the infrared region of the electromagnetic spectrum.

The energy difference between two energy levels in an atom is equal to the energy of the photon emitted or absorbed during a transition between those levels.

This equation relates the energy change during an electronic transition to the frequency of light emitted or absorbed (frequency is inversely proportional to wavelength). It is derived from the conservation of energy.

The specific wavelengths of light an atom can absorb when its electrons are excited to higher energy levels. Each element has a unique absorption spectrum.

The wavelengths of light absorbed by an element in its absorption spectrum correspond exactly to the wavelengths of light emitted in its emission spectrum.

The process where an electron in an atom moves to a higher energy level by absorbing a photon with the exact energy difference between the two levels.

A shorthand way to represent electron configurations using the principal quantum number (n) and azimuthal quantum number (l).

The number of electrons a subshell can hold is determined by the azimuthal quantum number (l).

The s subshell (l=0) has only one orbital, and the p subshell (l=1) has three orbitals.

The magnetic quantum number (ml) defines the specific orbital within a subshell where an electron is most likely to be found.

The possible values of ml range from -l to +l, including 0. For example, the p subshell (l=1) has ml values of -1, 0, and +1, representing three orbitals.

Each orbital within an atom can hold a maximum of two electrons.

The shapes of orbitals are determined by the subshell they belong to. s orbitals are spherical, while p orbitals are dumbbell-shaped.

The s orbital is spherical, while the three p orbitals are dumbbell-shaped and aligned along the x, y, and z axes.

Although subshell energies increase with increasing l values, subshells from different principal energy levels (n) can overlap.

The energy of the 4s subshell is lower than the 3d subshell, even though the 4s subshell is at a higher principal energy level.

Materials composed of atoms with unpaired electrons are attracted to a magnetic field due to the alignment of their spins. This property is called paramagnetism.

Materials consisting of atoms with all paired electrons are weakly repelled by a magnetic field. This property is called diamagnetism.

Valence electrons are the outermost electrons of an atom. They determine the atom's chemical behavior and are easily removed for bonding.

The outermost s and p subshells are involved in bonding for elements in Groups IIIA to VIIIA (13-18).

For transition metals, the highest s and d subshells (even if not the same principal quantum number) are involved in bonding.

A rule stating that electrons will individually occupy all available orbitals within a subshell before pairing up in any one orbital. It helps predict the electron configuration of atoms.

The configuration of electrons in an atom's orbitals, indicating the arrangement of electrons in different energy levels and subshells.

When a stable electron configuration is achieved by having a half-filled or completely filled subshell, making the atom more stable.

The tendency of electrons to repel each other, especially when in the same orbital.