Structure of atoms

‘Atom’ is derived from the Greek word ‘atomos’, which means ‘indivisible’.The English chemist John Dalton introduced this whole concept of atoms being the constituent particles of matter in the early 19th century. He put forth the ‘Atomic Theory’ whose main postulates were as follows:

• Elements consist of indivisible small particles known as atoms

• All atoms of the same element are identical; different elements have different types of atoms. 

• Atoms can neither be created nor destroyed.

And as years passed by, various discoveries were made and various theories were suggested by scientists regarding the atom’s structure and its constituents. The major ones are as follows – 

Discovery of Electron 

Around the mid 19th century, a group of scientists – J.J Thomson, W.Crookes & J.Perrin upon performing experiments based on the conduction of electricity through gases in a discharge tube under low pressures, had made some significant observations and conclusions like – 

• They observed that upon impressing a high voltage of almost 10,000 V across the electrodes, some invisible rays moved from the cathode to the anode. 

• This was detected by the scintillations observed on the Zinc Sulphide screen placed beyond the anode in the experimental setup.

• They observed that they travelled in straight lines with a high velocity of about 109 – 1010 cm/s

• It was observed that they were deflected in the presence of electric and magnetic fields, proving that they were charged and also that they discharged a positively charged gold leaf spectroscope proving that they carried a negative charge.

• They also possessed kinetic energy by placing a tiny wheel in between the vacuum tube and watching its blades set to motion. Hence, it was proved that cathode rays consisted of mass and velocity particles.

These particles were thus first termed as ‘negatrons’ by J.J Thomson and then renamed to ‘electrons’ by the Irish physicist George Stoney. Further on, the charge by mass ratio was successfully calculated by J.J Thomson and then the exact value of the charge of an electron by Robert A. Milikan through his famous Oil Drop Experiment. 

Discovery of Proton

Around the late 19th century, a German physicist Eugen Goldstein had performed the same cathode experiment with a perforated cathode and observed that upon applying the high potential difference in the tube, along with cathode rays passing from cathode to anode, another set of rays too existed simultaneously which were travelling from anode to cathode. He termed them as anode rays or canal rays as they originated from the anode. His significant observations regarding their characteristics are as follows-

• Like cathode rays, it was observed that they travelled in straight lines with a high velocity of about 109 – 1010 cm/s, as observed by scintillations on the Zinc Sulphide screen placed beyond the cathode.

• Like cathode rays it was also observed that they possessed kinetic energy by placing a tiny wheel in between the vacuum tube and watching its blades set to motion.

• It was also observed that they were deflected in the presence of electric and magnetic fields and also that they were attracted to a negatively charged plate, proving that they carried a positive charge.

J.J Thomson accurately measured the charge and mass of these particles in 1906, and later in 1911, these particles were named ‘Protons’ by Ernest Rutherford.

Thomson’s Plum Pudding Model

Soon after discovering electrons and protons, in 1904, the British physicist J.J Thomson had proposed a theory to summarise the structure of the atom known as the ‘Plum Pudding Model’.

According to this theory, the atom was supposed to be a sphere of positive charge embedded with negatively charged particles – electrons to make the whole entity neutral. But this theory was soon discarded after the discovery of neutrons by Ernest Rutherford.

Discovery of Nucleus

In the early 20th century, a New Zealand physicist and his co-workers had performed the landmark experiment known as the ‘Alpha Ray Scattering Experiment’ or ‘The Gold Foil Experiment’ with an experimental setup as follows –

• A radioactive source ( then a He ++ nucleus) that could emit α-particles was placed in front of a very thin (0.0004 cm thick) gold foil surrounded by a circular fluorescent Zinc Sulphide screen. 

The following observations were made during the experiment – 

• Whenever the α-particles struck the screen, they produced flashes

• 99% of the α-particles had followed a straight path, i.e., with no deflection

• Very few of the α-particles got deflected through small angles

• About one in 20,000 α-particles reverted after the incidence, i.e., they suffered a deflection of 180°

According to these observations, the following observations were made subsequently –

• Since most of the α-particles had followed a straight path, it was concluded that the atom is mostly hollow with a lot of space.

• Since very few α-particles get deflected through small angles, it was concluded that the entire positive charge is concentrated in a minimal space in the atom.

• Since a very few α-particles had suffered a considerable deflection, it was concluded that the atom supposedly consisted of a rigid positively charged portion at the centre, calling it the ‘nucleus.’

• The results were expressed in terms of these relations –

• N = K1/[(½)mv2]2 i

• N = K2/[sin4(𝚹/2)]

• N = K3/(Ze)2

• Where N ⇒ Number of α-particles striking the screen

            K1, K2, K3 ⇒ constants

            (½)mv2 ⇒ Kinetic energy of α-particles

            𝚹 ⇒ Scattering angle 

            Ze ⇒ Nuclear charge

Discovery of Neutron

As scientists couldn’t justify the atomic masses of atoms if we consider only protons and electrons, Rutherford in 1920 suggested the presence of another particle being neutral and having mass nearly equal to that of the proton. Then, in 1932, British physicist James Chadwick experimented by bombarding Beryllium with a stream of α-particles. 

It was observed that the produced penetrating radiations weren’t affected by electric or magnetic fields. Thus, it was concluded that these consisted of neutrally charged particles called neutrons, with a mass nearly equal to that of a proton.

Rutherford Model of Structure of an Atom

After the discovery of the fundamental particles of an atom – proton, neutron, electron, Rutherford proposed an atomic model, and its main postulates are as follows – 

1. An atom consists of a positively charged heavy nucleus at its centre consisting of protons and neutrons collectively called nucleons.

2. The nucleus is of minimal volume, almost with a diameter 10-5 times that of the entire atom. The radius of a nucleus is proportional to the cube root of its mass number as – R = R0A1/3 cm where R0 = 1.33×10-13 cm, A is the mass number

3. In the space around the nucleus, electrons are present whose number is equal to the number of protons present in an atom.

4. These electrons revolve with high speed around the nucleus in closed orbits whose centrifugal force experienced is balanced by the force of attraction between nucleus and electrons. 

Drawbacks:

• According to classical electromagnetic theory, when a charged particle moves under the influence of an attractive force, it is supposed to lose energy in the form of electromagnetic radiation. And subsequently, if this occurs, the electrons will spiral in and collapse into the nucleus, which is against the observed stability of an atom.

• If electrons lose energy continuously, it should lead to a continuous spectrum, but the spectrum is discontinuous according to what is observed. 

Bohr’s Atomic Model

To overcome the drawbacks of Rutherford’s model, a quantum mechanical model of the atom was proposed by a Danish physicist, Niels Bohr, in 1913. The main postulates of this theory are as follows-

• An atom consists of a dense nucleus consisting of neutrons and protons.

• The negatively charged electrons revolve around the nucleus in stationary circular paths called orbits, and as an electron remains in the stationary orbits, it doesn’t lose energy by itself.

• Each stationary orbit is allocated with a specific amount of energy and it increases with distance from the nucleus to the orbit.

• When an electron jumps from one stationary orbit to another, emission or absorption of energy occurs in the form of radiation.

Drawbacks:

• This theory doesn’t explain the spectra of multi-electron atoms

• It does not explain the Zeeman and Stark Effects

• It goes against Heisenberg’s principle