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Electromagnetic Radiation

A form of energy that travels through space at the speed of light, encompassing a range of wavelengths and frequencies.

Wavelength (λ)

The distance between successive peaks of a wave, typically measured in meters. It is inversely related to frequency.

Frequency (ν)

The number of waves that pass a point in one second, measured in hertz (Hz). It is directly related to energy.

Electromagnetic Spectrum

The range of all types of electromagnetic radiation, including radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, and gamma rays, each with distinct properties and applications.

Quantum (Planck)

The smallest discrete quantity of energy proportional to the frequency of the radiation it represents, foundational to quantum theory.

Photon (Einstein)

A particle representing a quantum of light or other electromagnetic radiation, carrying energy proportional to its frequency.

Wave Characteristics

Light exhibits wave-like properties characterized by wavelength, frequency, and speed. The relationship is given by the equation

Particle Characteristics

Light can also be described as particles (photons). The energy of a photon is calculated using the equation E = hν where E is energy, h is Planck's constant (6.626 x 10^-34 J·s), and ν is frequency.

Atomic Emission Spectra

When atoms are excited, they emit light at specific frequencies, resulting in a unique atomic emission spectrum, which differs from a continuous spectrum that contains all wavelengths.

Excited State Electrons

Electrons that have absorbed energy and moved to a higher energy level compared to their ground state.

Ground State Electrons

The lowest energy state of an electron in an atom, where it is closest to the nucleus.

Bohr’s Model

A model of the atom that describes electrons in fixed orbits around the nucleus, with quantized energy levels.

De Broglie

Proposed that particles, such as electrons, exhibit wave-like behavior, leading to the concept of electron waves.

Heisenberg Uncertainty Principle

States that it is impossible to simultaneously know both the position and momentum of an electron with absolute certainty.

Schrodinger

Developed wave functions that describe the probability of finding an electron in a given space, forming the basis of quantum mechanics.

Energy Level

Indicates the size of the orbital and the energy of the electron; higher levels correspond to greater energy.

Sublevel

Defines the shape of the orbital (s, p, d, f) and the distribution of electrons within an atom.

Atomic Orbital

Describes the orientation of the orbital in space, which can be spherical or dumbbell-shaped.

Electron Spin

A quantum property of electrons that describes their intrinsic angular momentum, represented as either +1/2 or -1/2.

Aufbau Order

The principle that electrons occupy the lowest energy orbitals first before filling higher ones.

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of four quantum numbers, ensuring that each electron has a unique state.

Hund’s Rule

Electrons will fill degenerate orbitals (orbitals of the same energy) singly before pairing up, to minimize repulsion.

Orbital Notation

A method of representing electrons in orbitals using arrows to indicate spin direction.

Electron Configuration

A notation that shows the distribution of electrons among the various orbitals, using superscripts to indicate the number of electrons in each orbital.

Noble Gas Shorthand

A simplified way to write electron configurations by using the nearest noble gas to represent inner-shell electrons.

Valence Electrons

The electrons in the outermost shell of an atom, crucial for chemical bonding and reactivity.

Ions

Charged particles formed when atoms gain or lose electrons.

Cations

Positively charged ions formed by losing electrons.

Anions

Negatively charged ions formed by gaining electrons.