How does the Bohr model account for the line spectrum of an atom?

Answer 1

An electron in an 'excited' state (orbital) loses a specific amount of energy thus an exact wavelength of energy, when it moves to a lower state. There are only exact orbitals that the electrons can occupy, thus all wavelengths are not emitted.

Answer 2

The Bohr model predicated that electrons surrounded or revolved around the nucleus of an atom in precise "orbits". These orbits could be imagined as an exact 'distance' from the nucleus and it is impossible for an electron to orbit at another distance than one of these orbits. A photon of precise wavelength is emitted when an electron jumps from one outer orbit to an inner (or closer) orbit. The opposite happens if a photon of the exact right amount of energy bumps an electron further out into a more distant orbit. The pushing out is called an 'absorption line'.

Answer 3

Before Bohr's theory, no one understood why the electron could maintain a stable orbit without it losing energy and spiraling into the centre. Bohr's work closely relates to the photoelectric effect, whereby as an electron is given more energy (by a photon) it goes into an excited state in a higher energy level, and when the electron goes back to the ground state it emits a photon. The spectra lines that we can observe are simply these photons emitted as a result of the electrons returning to a ground state. There are multiple spectrum lines for each atom because they can be excited to different degrees depending on the frequency of the photon.

Answer 4

If one adopts Bohr's basic postulate -- that electrons orbiting a nucleus can exist in only discrete distances from the nucleus -- then atomic emission line spectra (at least for hydrogen) begin to make perfect sense. When an electron moves from one such radius to another, its change of energy is exactly to the energy of the photon* emitted by the atom. Since the energy of this photon would be an exact number, so would its frequency. Bohr was eventually able to show that the predicted difference in energy levels -- ie the energy change that would result when an electron would move to a lower energy level -- matched the energy (ie, the frequency) of hydrogen emission lines.

* Bizarrely, however, Bohr didn't accept photons as real particles. He thought that eventually scientists would be able to show that an EM wave of a discrete frequency would result from this energy change, but would not involve discrete energy packets of light.

Answer 5

How the electrons have distinct energy states (orbitals).

Answer 6

Single atoms with one electron.

Answer 7

single atoms with one electron

Answer 8

line spectrum