Explain the production of an emission spectrum in terms of the energy states of an electron?

The production of an emission spectrum occurs due to atomic transition between energy states. This would be explained further using Bohr's energy quantization.

Explanation:Niels Bohr's great contribution to our understanding of atomic stricture was to add energy quantization to the possible orbits of electrons.

According to Bohr, there was only discrete allowed energy levels that the electron could occupy. If energy was added to an atom (through heat, electricity or light) the atom could absorb specific amounts of this energy. This would cause the electron (he dealt with hydrogen in his initial calculations) to jump into a higher energy orbit, known as a stationary state. Each state was characterized by a number,  n .

Suppose a given electron jumped from  

n =1  (the lowest possible energy) into  n = 4 . Such a high-energy electron was said to be in an excited state .

After a very short time, the electron would return to its lower energy level, either in one jump downward, or in a series of smaller jumps through  

n = 3 ,  n = 2  etc.

To make each jump, the atom would have to release its surplus energy, which it did in the form of a photon. Since there were only certain energy transitions possible, only certain colours of light would be seen. (The colour of the light was based on the energy and frequency of the emitted photon.)

These downward transitions of the excited electrons back to the ground state (the lowest energy) produced the line spectrum.

If the transition involved an electron dropping from a higher level into the  

n = 2  state, the photon was visible. Otherwise, ultraviolet and infrared photons could also be produced.

the answer is c i believe