Bohr-Rutherford Model and Atomic Spectra - Vocabulary Flashcards

Vocabulary flashcards covering key terms and definitions related to the Bohr-Rutherford atom model, energy levels, transitions, and atomic spectra.

Bohr-Rutherford Model

Atomic model with a dense nucleus containing protons and neutrons, surrounded by electrons on discrete energy levels.

Nucleus

The tiny, very dense central region of an atom containing protons and neutrons; protons determine the element’s identity.

Proton

Positively charged particle inside the nucleus.

Neutron

Electrically neutral particle inside the nucleus.

Electron

Negatively charged particle that orbits the nucleus in defined energy levels.

Energy level (shell)

Allowed energy states for electrons around the nucleus; levels become closer together at higher energies.

Ground state

The lowest energy state of an atom where electrons occupy the lowest available levels.

First energy level (n = 1)

Closest energy level to the nucleus; lowest energy and can hold up to 2 electrons.

Potential energy (PE)

Stored energy due to position; in Bohr’s model, electrons nearer the nucleus have lower (more negative) PE.

Energy gap between levels

The energy difference between two adjacent levels; the gap is largest between levels 1 and 2.

Unpopulated energy level

An energy level that is empty in the ground state but can be reached by absorbing energy.

Excitation

Process of absorbing energy to move an electron to a higher energy level.

Relaxation

Return of an excited electron to a lower energy level, releasing energy as light.

Ionization energy

Energy required to remove an electron from an atom (from the valence shell).

Quantum

The smallest discrete unit of energy exchange.

Planck's constant (h)

6.626 × 10^-34 J·s; constant relating energy to frequency (E = hf).

Frequency (f)

Number of cycles per second; unit is s^-1 (Hz); higher f corresponds to higher photon energy.

Wavelength (λ)

Distance between successive crests of a wave; inversely related to frequency.

Speed of light (c)

Constant 3.00 × 10^8 m/s; relates frequency and wavelength via c = λf.

E = hf

Planck’s equation: energy of a photon equals Planck’s constant times its frequency.

c = λf

Equation relating the speed of light to wavelength and frequency.

Photon

A quantum of light energy; particle-like packet that carries E = hf.

Excited state

A higher energy state reached after absorbing energy; electrons are not in ground state.

Absorption spectrum

A spectrum showing dark lines where specific wavelengths are absorbed by atoms.

Emission spectrum

A spectrum of bright lines at specific wavelengths emitted as electrons drop to lower levels.

Lyman series

Hydrogen spectral series for transitions to n = 1 (ultraviolet region).

Balmer series

Hydrogen spectral series for transitions to n = 2 (visible region).

Paschen series

Hydrogen spectral series for transitions to n = 3 (infrared region).

Hydrogen emission spectrum

Bright-line spectrum of hydrogen seen when it emits photons; contains characteristic lines.

Hydrogen absorption spectrum

Dark lines in the continuum spectrum where hydrogen absorbs photons.

Rydberg constant (R)

2.18 × 10^-18 J per atom; used in En = -R/n^2 for hydrogen energy levels.

En = -R/n^2

Hydrogen energy at level n in Bohr’s model; negative values indicate bound states.

Principal quantum number (n)

Integer labeling energy levels (n = 1, 2, 3, …).

Zero of energy (infinity)

Bohr’s convention: energy is defined as zero when the electron is infinitely far away from the nucleus.

Quantum leap

An electron’s transition between allowed energy levels that involves absorbing or emitting a quantum of energy.

Spectroscopy

Study of how matter absorbs or emits electromagnetic radiation; used to identify elements and analyze bonding and concentration.

Visible light

Portion of the electromagnetic spectrum visible to the human eye; Balmer lines lie in this region.

Bohr model successes

Accurately predicted hydrogen emission lines and introduced fixed energy levels; simple visualization.

Bohr model failures

Fails to predict wavelengths for multi-electron atoms and does not fully account for wave-particle duality.