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Niels Bohr proposed an atomic model in 1913, describing an atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the positively charged nucleus, similar to the orbits of the planets around the sun in our solar system, with attraction provided by electrostatic forces. The model was accepted in 1921.
The Bohr model of an atom is the name given to this model because it was developed by Niels Bohr. Bohr proposed an atomic model of a hydrogen atom, which is still in use today.
The stability of electrons revolving in orbits
The stability of electrons revolving in orbits was properly explained by Bohr’s model, which is still in use today. These orbits were given the moniker “energy shells” by him. L
The following is Bohr’s explanation for energy level:
Stationary states or energy levels are terms used to describe the various orbits in which electrons rotate. The stationary states/energy levels of an electron are denoted by the numbers n = 1, 2, 3,……….. The principal quantum numbers are a group of integers that are also known as the principal quantum numbers.
The energy of the stationary state in which an electron is placed can be calculated as follows:
In this case, RH is referred to as the Rydberg constant, and its value is 2.1810–18 J.
The energy of an electron when it is far removed from the influence of the nucleus is taken to be zero in this calculation. When an electron is in a stationary condition, the principal quantum number of the electron is taken to be equal to 1. An ionised hydrogen atom is a type of hydrogen atom that has been ionised. The above equation is given a negative sign because, as a result of the passage of an electron from one orbit to another, the stationary state energy is expelled and so the total energy is decreased.
When an electron is placed in the lowest stationary state/energy level that is feasible, it is referred to as being in the ground state of matter. During this energy level, the electron spins around the nucleus in an orbit with the shortest feasible radius. This state has an energy of –13.6 eV, which is very low.
Electrons Have Different Energy Levels
As you may recall from high school chemistry, an atom is made up of electrons that orbit around a nucleus. The electrons, on the other hand, are unable to determine their own orbit. They are limited to orbits with only a specific range of energy. Despite the fact that electrons are capable of jumping from one energy level to another, they can never have orbits with energies that exceed the permitted energy levels.
Let’s take a look at the most basic atom, which is a neutral hydrogen atom. The graphic below depicts the various energy levels of the object. The x-axis depicts the electron energy levels that are permitted in a hydrogen atom, which are numbered from 1 to 5. The energy of each level is represented on the y-axis in electron volts (eV). In electrochemistry, one electron volt (eV) is the amount of energy gained by an electron as it travels over a potential difference of one volt (1.6 x 10-19 Joules).
The electrons in a hydrogen atom must be in one of the energy levels that are permitted. It is necessary that an electron has exactly -13.6 eV of energy in order to be in the first energy level. The second energy level requires -3.4 eV of energy, which means it is in the second state. It is not possible for an electron in a hydrogen atom to have values of -9 eV, -8 eV, or any other value in between.
Consider the case in which the electron wishes to move from the first energy level, n = 1, to the second energy level, n =2. Given that the second energy level contains more energy than the first, the electron must obtain energy in order to go from n = 1 to n = 2 in order to gain momentum. Adding (-3.4) – (-13.6) = 10.2 eV of energy to the system is required in order to reach the second energy level.
By absorbing light, the electron can obtain the energy it requires to function. If an electron leaps from the second energy level to the first energy level, it must release some energy in the form of light in order to complete the transition. Photons are discrete packets of light that are absorbed or emitted by the atom, and each photon has a specific amount of energy. In order for the electron to leap between the n = 1 and n = 2 energy levels, it must absorb or emit a photon with an energy of exactly 10.2 eV in order for it to absorb or emit.
The amount of energy that a photon carries is determined by the wavelength of the photon. A particular wavelength must be observed for the photons absorbed or emitted by electrons hopping between the energy levels of n = 1 and n = 2, because these photons must have an exact 10.2 electron volts of energy in order to be absorbed or emitted. The equation can be used to determine this wavelength.
E = hc/ℷ, and
The energy of a photon (in electron volts) is given by E, and the Planck’s constant is given by h (4.14 x 10-15 EV s), and the speed of light is given by c (3.14×108 m/s). When this equation is rearranged to find the wavelength, the result is
ℷ= hc/E.
A photon with an energy of 10.2 eV has a wavelength of 1.21 x 10-7 m, which places it in the ultraviolet portion of the spectrum, and its wavelength is 1.21 x 10-7 m. As a result, in order for an electron to make the transition from n = 1 to n = 2, it must absorb a photon of ultraviolet light. Whenever the electron’s n = 2 value dips below 1, it emits an ultraviolet light photon known as an alpha photon.
Conclusion
The transition from the second to the third energy level is significantly more gradual. This jump requires only 1.89 eV of energy to complete. It takes even less energy to move from the third to the fourth energy level, and even less energy to move from the fourth to the fifth energy level.
