Atomic Structure and Spectra Quiz

A model of the atom where a dense, positively charged nucleus is surrounded by negatively charged electrons. It was revolutionary for its time, but fell short of explaining the spectral emission of light by atoms.

An experiment conducted by Geiger and Marsden, directed by Rutherford, where alpha particles were shot at a thin gold foil. The results provided evidence for the nuclear model of the atom.

Positively charged particles emitted by radioactive substances, such as bismuth-214. These particles are used in the Rutherford scattering experiment.

The thin sheet of gold used in Rutherford's experiment. Alpha particles were shot at this foil to study how they interact with the atoms.

The deflection of alpha particles when they hit the gold foil. The amount of scattering depends on the interaction with the nucleus.

The detector used in Rutherford's experiment to observe the scattered alpha particles. The screen produces flashes of light when struck by an alpha particle.

Specific colors or wavelengths of light emitted by atoms. These are unique for each element and cannot be explained by Rutherford's model.

Rutherford's model couldn't explain why atoms emit specific colors of light, as the electron's orbit would produce a continuous spectrum instead. This led to the development of new models like Bohr's model.

The distance between the nucleus and the electron in an atom.

The formula for calculating orbital radius is: r = (e^2)/(2 * E) where e is the charge of the electron, and E is the energy of the electron.

Unique patterns of light emitted or absorbed by an element, serving as a fingerprint for identification.

A spectrum with bright lines on a dark background, created by excited atoms releasing specific wavelengths of light.

Dark lines in a spectrum, corresponding to wavelengths absorbed by the material as white light passes through it.

Regular patterns observed in the frequencies of light emitted by a particular element.

Hydrogen is the simplest atom because it has only one proton and one electron.

The key to understanding atomic spectra is the excitation of atoms, which causes them to emit or absorb specific wavelengths of light.

The series of visible spectral lines of hydrogen, discovered by Johann Jakob Balmer in 1885.

The line with the longest wavelength in the Balmer series, at 656.3 nm, appearing red.

An empirical formula that predicts the wavelengths of the lines in the Balmer series.

A fundamental constant in atomic physics, representing the energy difference between the ground and first excited state of hydrogen.

Other spectral series of hydrogen, discovered after the Balmer series, named after their discoverers.

A spectrum containing all wavelengths of light, without any distinct lines or gaps.

The shortest wavelength in a spectral series, beyond which no distinct lines appear, only a continuous spectrum.

Regular patterns observed in the frequencies of light emitted by a particular element.

A spectral series in hydrogen where electrons transition from higher energy levels to the ground state (n=1).

Electrons can only exist in specific, stable orbits around the nucleus, called stationary orbits, without radiating energy.

The angular momentum of an electron in a stationary orbit is quantized, meaning it can only take on certain values.

An electron can jump between stationary orbits by absorbing or emitting a photon, the energy of which is equal to the energy difference between the orbits.

The relationship between the energy difference between two energy levels in an atom and the frequency of the photon emitted or absorbed during an electron transition.

An integer that describes the quantized values of an electron's energy, angular momentum, and spin, determining its specific state within an atom.

A model of the atom where electrons orbit the nucleus in specific energy levels, quantized according to Planck's constant. This model successfully explains the spectral lines of hydrogen.

A specific, quantized energy state that an electron can occupy in an atom, described by the principal quantum number (n). Higher energy levels correspond to orbits farther from the nucleus.

Energy is not continuous, but exists in discrete, fixed values. This applies to electron energy levels in an atom, explaining why atoms emit or absorb specific wavelengths of light.

A unit of energy commonly used in atomic physics, equal to the energy gained by an electron when it moves through a potential difference of one volt.

The lowest possible energy state an electron can occupy within an atom. Electrons are naturally found in this state unless excited.

An energy level higher than the ground state. Electrons can be excited by absorbing energy, moving to a higher energy level. They eventually decay back down to the ground state, releasing the absorbed energy.

Bright or dark lines on a spectroscope, indicating specific wavelengths of light emitted or absorbed by an atom. Each element has a unique set of lines, like an atomic fingerprint.

The spectrum of light emitted by excited atoms as they transition to lower energy levels. This spectrum consists of bright lines, each corresponding to a specific wavelength of light.

In multi-electron atoms, electron-electron interactions are significant, making the simple Bohr model inaccurate. This is because the forces between electrons are comparable to the electron-nucleus force, unlike in the solar system where planetary interactions are negligible compared to the Sun's gravitational pull.

Bohr introduced quantum numbers to describe electron orbits. These numbers quantize energy levels, meaning electrons can only exist at specific energy states. This is different from classical physics, where electrons could have any energy.

Quantum mechanics, the more refined theory, supports Bohr's idea of quantized energy levels. However, it allows for multiple quantum states at a given energy level, unlike Bohr's single-state model.

Bohr's model incorrectly assumed the electron's revolution frequency is the same as the frequency of the emitted spectral line. In reality, the emitted frequency is determined by the difference in energy levels divided by Planck's constant.

Bohr's model is a semi-classical approximation, combining elements of classical physics (like orbits) with quantum ideas (quantified energy levels). It's not a complete picture of atomic structure.

Despite its limitations, Bohr's model remains valuable for understanding atomic spectra and basic concepts. Its simplicity and ability to explain key features of hydrogen spectra make it a useful teaching tool.

Bohr's model only works for hydrogen. Its limitations include the incorrect assumption about the frequency of electron revolution and the lack of explanation for multiple electron atoms.

Bohr's model was a revolutionary step in understanding atomic structure. It introduced the concept of quantized energy levels and provided a foundation for later quantum theories.