What Led to the Bohr Model of Hydrogen Atom?
In 1897, Sir J.J. Thomson discovered electrons as negatively charged particles present in every element's atom, but without any knowledge of the distribution of electrons, the positive charge, and the mass inside the atom. Subsequently, in 1904, Sir Thomson suggested a model for the atom, also known as the 'plum pudding model,' which stated that the electrons are embedded like plums in a distribution (or pudding) of positive charge within the atom. Thomson's model failed to explain emission spectra and alpha particle scattering. Rutherford came up with another model in which the electrons revolve around the nucleus in different orbits. The revolution is driven by the electrostatic force of attraction between the nucleus and the electrons. But Rutherford's model failed to account for the stability of atoms and the origin of line spectra. To address the shortcomings of these previous models, Prof. Neils Bohr, in 1913, applied Planck's quantum theory and proposed three postulates that came to be known as the Bohr Model of Atom. So, let us discuss the Bohr Model of Hydrogen Atom (class 12) in detail.
Prior to Bohr, there were a number of scientists who were working on the structure of an atom. Rutherford was one of them and his model was the closest one to the Bohr model of the atom. In fact, Niels Bohr had helped to overcome the problem in Rutherford’s model of an atom.
Rutherford’s model of an atom had a major drawback, that it could not explain the stability. It showed that electrons in an atom revolve around a positively recharged centre called the nucleus. However, later on, it was found that any particle in a circular motion would undergo acceleration and thus would lose energy. So the electron would take a spiral path and would finally fall into the nucleus and the atom would collapse. But this is not true because this does not take place in reality.
Bohr said that electrons do revolve around the nucleus but their energy remains fixed. He explained that the energy of the electrons remains fixed because they are restricted to some fixed orbits. Each of these orbits is at a fixed distance from the nucleus and is associated with a fixed amount of energy. These energy levels are represented by the letters K, L, M, N, or the numbers 1, 2, 3, 4, starting from the centre.
So the final model of Neils Bohr was similar to Rutherford’s model of an atom which states that an atom consists of a positively charged centre in which the electrons revolve. The only difference was that in Bohr’s model, electrons revolve around the nucleus in fixed orbits with fixed energy. For his work on the structure of the atom, he got a Nobel Prize in 1922.
Bohr's Theory of Hydrogen Atom and Hydrogen-like Atoms
A hydrogen-like atom consists of a tiny positively-charged nucleus and an electron revolving around the nucleus in a stable circular orbit.
Bohr's Radius:
If 'e,' 'm,' and 'v' be the charge, mass, and velocity of the electron respectively, 'r' be the radius of the orbit, and Z be the atomic number, the equation for the radii of the permitted orbits is given by r = n2 ｘr1, where 'n' is the principal quantum number, and r1 is the least allowed radius for a hydrogen atom, known as Bohr's radius having a value of 0.53 Å.
Limitations or Problems of the Bohr Model
The model does not explain the Heisenberg Uncertainty Principle. The Bohr atomic model theory states that the electrons have both radius and orbit. which means to have a position and momentum simultaneously. This does not match with the Heisenberg Uncertainty Principle.
The theory Bohr devised was a mixture of classical and quantum physics. Quantum physics superseded classical physics, meaning quantum physics has everything classical physics has. This makes the approach of understanding the model of the atom invalid in some aspects.
The model could not explain the various intensities of the spectral lines, which are classified under the Stark effect.
The model could not explain the existence of the hyperfine structure of some spectral lines.
The model makes inaccurate spectral line predictions when it is concerned with larger atoms such as helium, lithium and oxygen or any other element. Bohr’s model only works with hydrogen.
The model does not explain the Zeeman effect, which is the splitting of spectral lines when placed in the magnetic field.
FAQs on Bohr Model of the Hydrogen Atom
1. How should I study the Bohr Model?
To study the Bohr Model effectively, students should first attempt to understand the concept and where it should be used. Studying tips such as flashcards and mnemonics (patterns to aid in memory) help students in memorizing the formulas and rules along with principles and finally limitations of the topic. Students are advised to refer to the online course material from the official website of Vedantu, www.vedantu.com. The course material was curated by subject-matter experts for the sole purpose of helping out students in mapping and structuralizing the concepts for better understanding.
2. What are the improvements to the Bohr Model?
There were a few improvements to the Bohr atomic model. They are the Sommerfeld model - which is also known as the Bohr-Sommerfeld model- and Wolfgang Pauli’s model. The Sommerfeld model was better at explaining the spectral effects of atoms but it could not contain the magnetic quantum number. It showed how electrons orbited around the atoms in an elliptical fashion. Wolfgang’s model was improved to understand the quantum mechanics of an atom. It used waves to describe the orbital aspects of the atom.
3. What are the main points to remember about the Bohr Model?
The main features of the Neils Bohr atomic model are:
Orbits have a nucleus around which the electrons revolve, without emitting radiant energy. These orbits have definite energy. They are also known as energy shells or energy levels.
These orbits are denoted by K, L, M, N shells or 1, 2, 3, 4, 5 starting from the centre. The electron is in the ground state when its energy level is the lowest.
When the electron jumps from one orbit to another energy level, it either emits or absorbs energy while doing so.
4. How do I study Bohr’s Model and other topics in physics easily?
There are a few effective tips you can use to study physics easily. Some of the most effective proven ones are listed as follows:
Master the basics - once you familiarize yourself and gain mastery over the basics of physics, the following chapters become easier.
Keep working on improving your math skills - since maths is an important factor in physics, pay attention to it. You can even categorize the topics for the equations, like algebra for basic equations and geometry for area and volume problems.
5. State the important drawbacks of rutherford's model of the atom?
Rutherford's model of the atom suffered from two drawbacks:
Stability of Atoms: The electrons orbiting the nucleus possess centripetal acceleration. According to classical electrodynamics, accelerated charged particles emit electromagnetic energy, and its continuous radiation by the orbiting electrons leads to the electrons losing energy. Thus, the orbits' radii should continuously decrease, and eventually, the electrons will collapse into the nucleus. Therefore, atoms cannot remain stable.
Explanation of Line Spectrum: In Rutherford's model, the continuous change in the radii of orbits must also cause a constant change in the frequency of the revolution of electrons. Thus, the electrons will radiate electromagnetic waves of all frequencies, giving rise to a 'continuous' spectrum. But, experimental observation proves that the atomic spectra are not continuous, but consist of discrete spectral lines, each corresponding to a specific frequency. So, an atom should be radiating waves of particular frequencies.
6. What is the Hydrogen Spectrum?
The Hydrogen spectrum was studied by Balmer. It consists of discrete bright lines (Hα, Hβ, H४...) on a dark background. From one end to another, the brightness and the separation between these lines gradually decrease; these lines are members of a series called the Balmer series, and they are found in the visible part of the spectrum. Other series were found in the invisible parts of the spectrum - Lyman series in the ultraviolet region, and the Paschen, Brackett, and Pfund series in the infrared area. The wavelength of the lines in these series can be represented by the Rydberg equation:
\[\frac{1}{\lambda }\] = R \[\left [ \left ( \frac{1}{n12} \right )-\left ( \frac{1}{n22} \right ) \right ]\] , where R is the Rydberg constant, n1, and n2 are the lower and the higher energy levels, respectively, of the hydrogen atom. The transitions of electrons between the energy levels give rise to the line spectrum of hydrogen. For Balmer series, n1 = 2.
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