Rutherford's model

O

Fig. 3. Setup of alpha particles scattering experiment

ther ground-breaking experiment which provided us with our understanding of the structure of the atom is Ernest Rutherford's experiments with alpha particle bombardment of thin metal foils. These experiments clearly provided the information which led to the formulation of the atom as a nuclear atom. Most of the atom is empty space. The vast majority of the mass of an atom, and all of the positive charge, is concentrated in a very tiny volume called the nucleus. A typical atom has a diameter of 10^-8 cm, while the diameter of a nucleus is about 100,000 times smaller: on the order of 10^-13 cm. ¹

We know that the mass of an alpha particle is almost 7500 times greater than that of an electron. Most of the alpha particles were found to pass through the foil without changing their direction, although the thickness of the foil was hundreds of thousands of atomic diameters.

It followed from the results of these experiments that the predominating part of the space occupied by an atom of a metal contains no heavy particles—only electrons can be there.

On the basis of these considerations, Rutherford in 1911 the planetary model of the atom, in which the electrons revolve in orbits around the nuclear “sun”. The main features of this model — the presence in an atom of a positively charged heavy nucleus surrounded by negatively charged electrons. The atom as a whole is electrically neutral. The centrifugal force produced owing to rotation of the electrons around the nucleus is balanced by the force of electrostatic attraction of the electrons to the oppositely charged nucleus. The size of the nucleus is very small in comparison with that of an atom as a whole:

Rutherford's model, however, in some respects contradicted firmly established facts. We shall note two of these contradictions.

First, Rutherford's theory could not explain the stability of the atom. An electron orbiting about a positively charged nucleus ought to emit electromagnetic energy in the form of light waves like an oscillating electric the nucleus. After exhausting all of its energy, it should "fall" onto the nucleus, and the atom will stop existing. This conclusion contradicts the real properties of atoms, which are stable formations and can exist without decay for an exceedingly long time.

Second, Rutherford's model led to improper conclusions on the nature of atomic spectra. As indicated above, an electron revolving about a nucleus should approach the latter, continuously changing its speed. This signifies that the emission spectrum of an atom should be continuous, but this does not correspond to actual facts. Thus, Rutherford's theory could explain neither the existence of stable atoms nor the presence of line spectra of atoms.

Rutherford postulated that the nucleus carries a positive charge, but it was Henry Moseley who actually determined the charges of nuclei of most of the atoms. Moseley observed a specific wavelength of the X rays produced in special tube depend on the kind of metal used as a target. X rays are electromagnetic radiation of high frequency and short wavelength. X rays are produced when high-speed electrons (cathode rays) strike the metal target in a evacuated tube. On the basis of this experimental data in 1913 he formulated the law:

The square root of frequency (n=1\l) of definite lines of identical series of the characteristic X-ray spectrum directly proportional to atomic number of an element a (Z — b) a, b are constant, Z is an atomic number, l is a wavelength of irradiation of an element with atomic number Z

So, the atomic number (Z) in the Periodic Table is equal to the number of protons in the nucleus of the atom.

Rutherford had to postulate the existence of a neutral particles in atom due to the failure to account for known mass of the atom in terms of the number of the protons in nucleus. For example, a helium nucleus has a charge of +2 furnished by the 2 protons with total mass 2 amu but helium atoms have a mass 4 amu. Only in 1932 James Chadwick experimentally proved the existence of such particles that have a mass approximately 1 amu and were electrically neutral. Accordingly, the particle was called neutron.

Here we consider a brief summary as to fundamental subatomic particles:

Table 1. Fundamental Subatomic Particles

Particle	Symbol	Charge	Mass
Electron	e-	-1	0.0005486 amu
Proton	p+	+1	1.007276 amu
Neutron	no	0	1.008665 amu

The number of protons in the nucleus of the atom is equal to the atomic number (Z).

The number of electrons in a neutral atom is equal to the number of protons.

The mass number of the atom (M) is equal to the sum of the number of protons and neutrons in the nucleus.

The number of neutrons is equal to the difference between the mass number of the atom (M) and the atomic number (Z).

Isotopes are atoms of the same element with different numbers of neutrons, such as the ²⁰Ne and ²²Ne isotopes of neon or the ³⁵Cl and ³⁷Cl isotopes of chlorine.