Table of Contents
Atomic structure is defined as the structure of an atom containing a nucleus present in the center in which the protons or positively charged particles and neutrons (neutral) are present. The negatively charged particles are termed as electrons and they revolve around the nucleus.
In the 1880s, the first scientific theory of atomic structure was explained by John Dalton. A variety of different models have been evolved over the past decades to understand the functions of an atom. As a result, there are five basic atomic models which helped us to describe and comprehend the structure of the atom. Each of these models mentioned below had its own advantages and drawbacks.
The five atomic models that shaped the modern atomic theory are:
• John Dalton’s atomic model
• J.J. Thomson’s atomic model
• Ernest Rutherford’s atomic model
• Niels Bohr’s atomic model
• Quantum Numbers/model
I. Dalton Atomic Model
The English chemist and scientist named John Dalton stated that all matter is made up of atoms, which are undividable. He also proposed that all the atoms present in an element are the same, but the atoms of different elements generally differ in their size and mass. The following are the postulates of Dalton’s theory;
• All matter is made up of particles called atoms.
• Atoms are indivisible particles.
• Specific elements generally have only one type of atom present in them.
• Each atom has a constant mass respectively that differs from element to element.
• Atoms can neither be formed nor can be destroyed but can be transformed from one form to another.
Drawbacks of Dalton Atomic Model:
• The theory was not able to describe the existence of isotopes.
• Dalton’s atomic theory does not explain the existence of subatomic particles. Dalton’s atomic theory projected that the atoms were indivisible. However, the discovery of subatomic particles (for example, protons, electrons, and neutrons) disproved this postulate.
• Dalton’s atomic theory failed to explain isobars( two different elements having the same mass number. For Instance: 40Ar and 40Ca)
II. Thomson Atomic Model
The English chemist and scientist Sir Joseph John Thomson described the structure of the atom in the early 1900s.
He was awarded the Nobel prize later for the finding of “electrons”. His work is chiefly based on an experiment titled a cathode ray experiment. The working of this experiment is as follows:
Cathode Ray Experiment: It has a tube made of glass which further has two openings, one opening is for the vacuum pump and the other one is for the inlet through which a gas is pumped.
A high voltage electric current is passed through a discharge tube that contains gas at a very low pressure, a green glow is thus seen at the other end of the discharge tube. This green glow or fluorescence observed is the result of the rays which are released from the cathode towards the anode. These rays are termed as cathode rays.
Conclusions: Based on this observation from his cathode ray experiment, Thomson defined the atomic structure as a positively charged sphere that contains negatively charged particles called electrons.
It is usually stated as the “plum pudding model” because it can be pictured as a plum pudding dish where the pudding represents the positively charged atom and the plum pieces in it represent the electrons.
Drawbacks of Thomson’s Atomic Model: The theory did not mention anything about the nucleus present in the center of an atom.
III. Rutherford Atomic Model
Rutherford, a famous scientist revised the structure of an atom with the discovery of another subatomic particle named as a Nucleus. His atomic model is built on the Alpha ray scattering experiment.
Alpha Ray Scattering Experiment Structure:
• Rutherford took a gold foil as he wanted a fragile layer.
• In this experiment, fast-moving alpha particles were bombarded on a thin gold foil.
• Alpha particles are referred to the helium ions with a +2 charge and thus have a significant amount of energy.
• Rutherford predicted that the alpha particles would pass through the gold foil but some of the particles deflected and striked the fluorescent screen.
• Since most of the rays passed straight through the gold foil, Rutherford thus observed that most of the space inside the atom is vacant or empty.
• Few rays which got reflected is because of the repulsion with some other positive charge present inside the atom.
• 1/1000th of rays got forcefully deflected because of the presence of a very strong positive charge confined in the center of the atom. He named this strong positive charge as “nucleus”.
• He stated that most of the charge and the mass of the atom is present in the center (Nucleus).
Rutherford’s Structure of Atom Based on the above comments and assumptions, Rutherford projected his own atomic structure which is;
• The nucleus is present at the center of an atom, generally where most of the charge and mass of the atom is concentrated.
• Electrons revolve around the nucleus (which is present in the center) in circular paths called orbits, just like the planets revolve around the sun.
Limitations of Rutherford Atomic Model:
• If electrons present in an atom revolve around the nucleus, then they have to spend energy, as a result, a lot of energy will be spent by the electrons, and ultimately, electrons will lose all their energy and will fall into the nucleus thus Rutherford was unable to explain the stability of atoms.
IV. Bohr Atomic Model
Neils Bohr model is the most widely used atomic model which helped to define the atomic structure of an element that is built on Planck’s theory of quantization. Bohr’s Postulates:
• The electrons present inside the atoms are positioned in discrete orbits termed as “stationary orbits”.
• The energy levels of these shells can be denoted as quantum numbers.
• Electrons can travel to higher levels by absorbing energy and can also move to lower energy levels by emitting or releasing their energy.
• When an electron stays in its rest or stationary form, then there will be no absorption or release of energy.
Drawbacks of Bohr’s Atomic Model:
• Bohr’s atomic structure is applicable only for single-electron species. For instance; H, He+, etc.
• Bohr’s theory was unable to explain both Stark and Zeeman’s effects.
V. Quantum Numbers
• Principal Quantum number (n): It represents the orbital number or the shell number of the electron.
• Azimuthal Quantum numbers (l): It represents the orbital (sub-orbit) of the electron.
• Magnetic Quantum number: It represents the number of energy states present in each orbit.
• Spin Quantum number(s): It represents the direction of spin, that is when S = -½ the spin of an electron is Anticlockwise and ½ then the spin of an electron is Clockwise.
Electronic Configuration of an Atom:
The electrons are usually filled in the s, p, d, f orbits as per the following rule; 1.
Aufbau’s principle: The filling of electrons must take place by following the ascending order of the energy orbitals that is;
• Initially, the lower energy orbital should be filled and then the higher energy levels.
• Ascending order of energy orbitals is as follows 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on. 2.
Pauli’s exclusion principle: this principle states that no two electrons can have all the four quantum numbers mentioned above to be the same or similar.
NOTE: If two electrons are positioned in the same energy state then they should be positioned with opposite spines.
Atomic Model Citations
- Nagaoka’s atomic model and hyperfine interactions. Proc Jpn Acad Ser B Phys Biol Sci . 2016;92(4):121-34.
- Dalton’s disputed nitric oxide experiments and the origins of his atomic theory. Chemphyschem . 2008 Jan 11;9(1):106-10.
- Atomic Theory and Multiple Combining Proportions: The Search for Whole Number Ratios. Ann Sci . 2015 Apr;72(2):153-69.