Atomic Structure

Atom was considered as the smallest particle up to the 19th century. A series of experiments were performed to reveal the structure of the atom as well as to explain its important properties. These experiments indicated the divisibility of tom into sub-atomic particles and showed that atoms possess a definite internal configuration and composition.

Dalton’s Atomic Theory

In 1808, John Dalton published ‘A new system of chemical philosophy’ in which he proposed the following theory

  1. Matter consists of indivisible atoms
  2. All the atoms of a given element have identical properties including identical mass. Atoms of different elements differ in mass.
  3. Compounds are formed when atoms of different elements combine in a fixed ratio.
  4. Chemical reactions only involve reorganization of atoms. The atoms are neither created nor destroyed in a chemical reaction.
  5. Dalton’s atomic theory could explain the law of chemical combination.

Earlier Atomic Models

Different atomic models were proposed to explain the distribution of charged particles i.e., electron, proton, a neutron in an atom.

Thomson Model of an Atom

Thomson proposed the model of an atom to be similar to that of a Christmus pudding. Thomson proposed that

  1. an atom consists of a positively charged sphere and the electrons are embedded into it.
  2. The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral.

Thomson models

Drawback Although Thomson model explained that atoms are electrically neutral but the results of experiments carried out by other scientists like α-particle scattering experiment could not be explained by this model.

 Rutherford Model of an Atom

Rutherford and his students (Hans Geiger and Ernest Marsden) in 1911 performed α-particles scattering experiments in which they bombarded very thin gold foil with α-particles. On the basis of the observations and conclusions, Rutherford proposed the nuclear model of the atom. According to this model,

  1. There is a positively charged spherical center in an atom, called the nucleus. Nearly all the mass of an atom resides in the nucleus. i.e., protons and neutrons are packed into it.
  2. The electrons revolve around the nucleus in well-defined orbits. Thus, most of the part of an atom is empty.
  3. The size of the nucleus is very small as compared to the size of the atom.

Rutherfold model

Drawbacks According to the classical theory of electrodynamics, any charged particle in a circular orbit would undergo acceleration. During centripetal acceleration, the charged particles would radiate energy. Thus, the revolving electron would lose energy and come closer and closer to the nucleus and finally fall into the nucleus. If this were so, the atom should be highly unstable. But we know that atoms are quite stable, so this model was discarded.

Bohr’s Model of an Atom

According to Bohr, the old classical laws cannot hold good in case of subatomic particles. In order to overcome the objections raised against Rutherford’s model of an atoms, Neils Bohr (1913) utilized the concept of quantisation (Max Planck) and put forward the following postulates on the basis of Plank’s quantum theory about the model of an atom.

  1. The electrons continue revolving in their respective orbits without losing energy. Thus, each orbit (shel) is associated with a definite energy hence, it is also called energy level.
  2. The electrostatic coulombic force of attraction between the nucleus and the electron counterbalanced the centripetal force required for revolving the electron
  3. The electrons can move in only those circular orbits where, the angular momentum (mvr) is a whole number multiple h/2π i.e., it is quantized
  4. Energy is emitted or absorbed by an atom only when an electron moves from one level to another

Bohr's Model of Atom

Drawbacks of Bohr’s Model of an atom

  1. This model is unable to explain the spectrum of atoms other than hydrogen e.g., helium atom which possesses only two electrons.
  2. This theory was also unable to explain the splitting of spectral lines in the presence of magnetic field (Zeeman effect) or an electric field (Stark effect).
  3. It could not explain the ability of atoms to form molecules by chemical bonds.

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