The reason for the existence of distinct energy levels or states in an atom can be explained by the quantization of energy in quantum mechanics. According to quantum theory, electrons in an atom can only occupy specific energy levels, and the energy of an electron is quantized, meaning it can only have certain discrete values.
The energy levels in an atom are determined by the electron's arrangement around the nucleus and are characterized by the principal quantum number (n). The ground state corresponds to the lowest energy level, where electrons are closest to the nucleus. The ground state is typically denoted as n = 1.
When energy is supplied to an atom, such as through absorption of photons or thermal excitation, electrons can absorb this energy and move to higher energy levels. These higher energy levels are known as excited states. The first excited state corresponds to the energy level immediately above the ground state (n = 2), the second excited state corresponds to the next energy level (n = 3), and so on.
The energy difference between each energy level is not equal. As the energy levels get farther from the nucleus, the spacing between them decreases, resulting in a converging series of energy levels. This is why the energy difference between the ground state and the first excited state is typically larger than the energy difference between higher excited states.
When an electron transitions from a higher energy level to a lower one, it emits energy in the form of photons. This emission corresponds to the emission spectra observed in atomic spectroscopy. The emitted photons have energies corresponding to the energy difference between the initial and final energy levels of the electron.
In summary, the existence of distinct energy levels in an atom arises from the quantization of energy in quantum mechanics. These energy levels are characterized by the principal quantum number, and transitions between these levels result in the absorption or emission of discrete amounts of energy.