Rutherford performed an alpha particles scattering experiment which revolutionized the understanding of scientists about the structure of atom. In this article, you will know, how Rutherford's experiment works, what conclusions it led to, and its applications.
Alpha particles scattering experiment
Setup: Rutherford fired a beam of alpha particles (positively charged helium nuclei) at a very thin sheet of gold foil. Alpha particles are energetic and were thought to be indivisible at the time.
Expectation: The prevailing atomic model (Thomson's model) pictured the atom as a uniform sphere of positive charge with negatively charged electrons scattered throughout. Based on this model, Rutherford expected the alpha particles to pass through the gold foil with minimal deflection, like shooting bullets through pudding.
Observations from the Experiment
(i) Most of the α-particles (nearly 99.9%) went straight without suffering any deflection.
(ii) A few of them got deflected through small angles.
(iii) A very few α-particles (about one in 20,000) did not pass through the foil at all but suffered large deflections (more than 90°) or even come back in the direction from which they have come i.e. a deflection of 180°.
Conclusions made by Rutherford
(1) Since most of the α-particle went straight through the metal foil undeflected, it means that most of the space inside atom is empty.
(2) Since few of the α-particles were deflected from their original path through moderate angles; it was concluded that whole of the positive charge is concentrated in a very tiny space (later known ass nucleus).
(3) The positively charged heavy mass which occupies only a small volume in an atom is called nucleus and it is present at the center of atom.
(4) Electrons must be outside the nucleus, but the experiment didn't tell him how they were arranged.
(5) A very few of the α-particles suffered strong deflections on even returned on their path indicating that the nucleus is rigid and α-particles recoil due to direct collision with the heavy positively charged mass.
(6) There is a mathematical relation between the angle of deviation and number of particles undergone that deviation.
where, μ is the number of particles; and θ is the deviation angle.
This experiment laid the foundation for the nuclear model of the atom, which is the basis for our modern understanding of atomic structure.
Rutherford model of gold atom
Atom is a spherical substance with positively charged heavy mass nucleus placed at the center in a very tiny space and electrons spread in the outer space of the nucleus. This model correctly describe the nucleus but not the electrons. Hence, it is also known as Rutherford's nuclear model.
Applications of Rutherford's model of atom
The nuclear model gave by Rutherford explained various theories. Such as,
(1) Protons are present in the nucleus: An atom consists of a heavy positively charged nucleus where all the protons are present.
(2) Radius of Atom and radius of Nucleus: The volume of the nucleus is very small and is only a minute fraction of the total volume of the atom. Nucleus has a radius of the order of 10-13 cm and the atom has a radius of the order of 10-8 cm
Thus radius (size) of the atom is 105 times the radius of the nucleus.
(3) Volume of atom and volume of nucleus: There is an empty space around the nucleus called extra nuclear part. In this part electrons are present. The no. of electrons in an atom is always equal to no. of protons present in the nucleus. As the nuclear part of atom is responsible for the mass of the atom, the extra nuclear part is responsible for its volume. The volume of the atom is about 10^15 times the volume of the nucleus.
(4) Place of electrons in the atom: Electrons revolve round the nucleus in closed orbits with high speeds. This model was similar to the solar system, the nucleus representing the sun and revolving electrons as planets.
Limitations of the Rutherford's atomic model
(1) Stability of the atom: It could not explain the stability of an atom. According to Maxwell, electron loses it's energy continuously in the form of electromagnetic radiations. As a result of this, the e- should loss energy at every turn and move closer and closer to the nucleus following a spiral path. The ultimate result will be that it will fall into the nucleus, thereby making the atom unstable.
(2) If the electrons loss energy continuously, the observed spectrum should be continuous but the actual observed spectrum consists of well defined lines of definite frequencies (discontinuous). Hence, the loss of energy by electron is not continuous in an atom. (The spectrum made by electrons will be explained later in this chapter.)
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