Atomic Structure and Models

Atoms consist of a uniformly distributed positive charge within a sphere with negatively charged electrons embedded within it.

Helium nucleus containing two protons and two neutrons.

The tiny, dense core of an atom containing protons and neutrons, where most of the atom's mass is concentrated.

The distance an alpha particle comes close to a nucleus before being repelled by its positive charge.

Dalton proposed that all matter is made up of tiny, indivisible particles called atoms.

Rutherford's experiment involved firing alpha particles at a thin gold foil. Most particles passed through, but some were deflected at large angles.

The Thomson model couldn't explain the spectral lines of hydrogen and other atoms and didn't account for the scattering of alpha particles at large angles.

Rutherford's model proposed that atoms have a small positively charged nucleus surrounded by negatively charged electrons orbiting it.

The perpendicular distance between the initial path of an alpha particle and the center of the nucleus it's approaching. It essentially determines how close the alpha particle gets to the nucleus.

The idea that electrons can only exist in specific orbits around the nucleus, not anywhere in between.

Bohr's idea that an electron only emits or absorbs energy when it jumps from one allowed orbit to another.

The force that keeps an electron in its orbit around the nucleus, provided by the attraction between the positive nucleus and negative electron.

The equation that relates the electron's speed, mass, charge, and the distance from the nucleus, used to calculate the electron's speed in a particular orbit.

A key concept in Bohr's model, it states that the angular momentum of an electron in a stable orbit is a whole number multiple of h/2π (Planck's constant).

The radii of the allowed orbits for an electron around a hydrogen nucleus, calculated using Bohr's postulates.

The model of the hydrogen atom proposed by Niels Bohr, incorporating both classical physics and quantum concepts to explain the atom's behavior.

The radius of a stationary orbit in the Bohr model is quantized, meaning it can only take on discrete values. This is determined by the principal quantum number 'n'. Larger 'n' corresponds to larger orbit radii.

The velocity of an electron in a Bohr orbit is inversely proportional to the principal quantum number 'n'. This means that electrons in higher energy levels move slower.

The frequency of revolution of an electron in a Bohr orbit is the number of times it completes a circular path around the nucleus per second. It is related to the electron's velocity and the orbit's radius.

The total energy of an electron in a Bohr orbit is negative, indicating that it is bound to the nucleus. It consists of both kinetic and potential energies, with the potential energy being negative and greater in magnitude than the kinetic energy.

The minimum energy required to remove an electron from its ground state and completely free it from the atom. For hydrogen, this is 13.6 eV.

The discrete wavelengths of light emitted or absorbed by an atom when electrons transition between energy levels. This is explained by the Bohr model and the Rydberg formula.

Describes the wavelengths of spectral lines emitted by hydrogen, and it is based on the quantized nature of electron energy levels and transitions between them.

The lowest energy state of an atom, corresponding to the electron occupying the innermost orbit (n=1).

The energy difference between two energy levels in an atom, where 'E2' is the higher energy level and 'E1' is the lower energy level, is equal to the energy of a photon emitted or absorbed during the transition between these levels. It is a fundamental equation used in Atomic Physics.

The Rydberg formula relates the wavelengths of spectral lines emitted by a hydrogen atom to the energy levels involved in the transitions. It predicts the wavelengths of the different spectral series observed in the hydrogen spectrum.

The Rydberg Constant is a physical constant used in the Rydberg Formula to calculate the wavelengths of spectral lines emitted by a hydrogen atom. It is a fundamental constant in atomic physics.

A series of spectral lines observed in the ultraviolet region of the hydrogen spectrum, emitted when electrons transition from higher energy levels to the ground state (n1=1) of the atom.

A series of spectral lines observed in the visible region of the hydrogen spectrum, occurs when electrons transition from higher energy levels to the second energy level of the atom (n1=2).

A series of spectral lines observed in the infrared region of the hydrogen spectrum, created by electronic transitions from higher energy levels to the third energy level of the atom (n1=3).

A series of spectral lines in the infrared region of the hydrogen spectrum, produced by electronic transitions from higher energy levels to the fourth energy level of the atom (n1=4).

A graphic representation of the permitted energy levels of electrons in an atom (the stationary states, as described by Bohr's model). The different energy levels correspond to different electron orbits around the nucleus. The energy levels are shown on a vertical axis, with higher energies at the top. Electrons can jump between these energy levels, absorbing or emitting photons of light. This process gives rise to the characteristic spectral lines observed in atomic emission spectra.

The total energy of an electron in the nth orbit of a hydrogen atom. It's expressed in electron volts (eV).

The energy level diagram of a hydrogen atom shows different energy levels available for electrons. Each level is labeled by its principal quantum number, 'n'.

The energy level diagram of a hydrogen atom shows different energy levels available for electrons. Each energy level is labeled by its principal quantum number, 'n'.

The minimum potential required to excite an electron in an atom from its ground state to a higher energy level.

The potential required to completely remove an electron from an atom, leaving it positively charged.

de Broglie proposed that particles exhibit wave-like properties. He explained Bohr's quantization postulate by suggesting that electron orbits contain an integral number of electron wavelengths.

Bohr's model, while successful for hydrogen, had limitations. It couldn't explain the fine structure of spectral lines, the intensities of spectral lines, and it only worked for hydrogen.