Atom Model & Periodic Properties

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1 One Change Physics (OCP) MENU Atom Model & Periodic Properties 5.1 Atomic Mode In the nineteenth century it was clear to the scientists that the chemical elements consist of atoms. Although, no one gas so far seen individual atoms, yet there is no doubt for its existence, Electron, one of the fundamental constituent principle, was first of all invented by J.J. Thomson. Electrons are negatively charged particles and the atom as a whole is electrically neutral. Hence it became clear that the atom must also contain positively charged material to balance the negative charge of the electrons. This was confirmed by the discovery of positive rays which indicated that a normal atom consisted of both negative and positive charges. Thomson also measured the ratio of charge to mass e/m for an electron whereas the charge e on the electron was measured by Millikan by his famous oil drop experiment. Thus the mass of the electron was found by dividing charge e by the ratio e/m. Since the normal atom is electrically neutral. The quantity of positive charge per atom was thus known. Thomson also showed that the mass of the electron is for the order of 1/2000th of the mass of hydrogen atom which clearly indicated that the entire mass of the atom is associated with its positive charge. It was not known at that time that how the positive and negative charges are distributed in an atom. To account for the experimentally observed spectroscopic data at that time, several theories have been proposed from time to time regarding the atomic structures which are known as atomic models. The various models are as follows: (1) Thomson s Plum pudding model, (2) Rutherford s nuclear model, (3) Bohr s model, (4) Summerfield s relativistic model, (5) Vector model and (6) Wave mechanical model. Following we discuss first three models one by one for understand the model. 5.2 Thomson s plum pudding model J.J. Thomson proposed the plum pudding model considering the atom to consist of a heavy sphere about meter in diameter According to this model, the electrons and positive particles are uniformly distributed inside the spherical atom. The uniform distribution is such that repulsive force action between two adjacent electrons is balanced by the equal and opposite attractive forces of action between the electron and nearly positive particle. This model is shown in Fig 5.2 According to Thomson,

2 Fig 5.2 Thomson s Atom model If there is a single electron in the atom (like a hydrogen atom), the electron must be situated at the centre of the positive sphere. For an atom with two electrons (for helium atom), the electrons should be situated symmetrically with respect to the center of the sphere, i.e., opposite sides of the centre at a distance r/2, where r is the radius of the positive sphere. In a three electrons system of the atom, the electrons should be at the corners of a symmetrically placed equilateral triangle, the side of which was equal to the radius of the sphere. In general, the electrons of an atom are located in a symmetrically pattern with respect to the centre of the sphere. It was suggested that spectral radiations are being resulted by simple harmonic motion of these electrons on both sides of their mean positions. Moreover, the stability of the atom was very well explained on the basis of this model. This atomic model was given up after some time due to the following reasons (i) According to electro magnetic theory, the vibrating electron should radiate energy and the frequency of the emitted spectral line should be the same as the electron, In case of hydrogen atom, Thomson observed only one spectral line of about 1300A0 by this assumption while experimental observations reveal that hydrogen spectrum consists of five different series with several lines in each series. (ii) This model could not provide any satisfactory mechanism for explaining the large deflection suffered by α particles in Rutherford s experiment (Rutherford s experiment is discussed in the next article.) 5.3 Rutherford s Experiment on α Particle scattering A fine pencil of particles (helium nuclei each with a charge of +2e) was obtained from a radioactive material like radium or radon by placing it in a narrow tunnel like opening in a lead block as shown in Fig 5.3 The α- particle re-emitted in all directions but only those came out

Fig 5.3 Ruther ford s Experiment on α Particle Scattering which are along with the axis of opening. Thus they are confined to a narrow beam. This beam is now made to strike a thin gold foil G.

3 Fig 5.3 Ruther ford s Experiment on α Particle Scattering which are along with the axis of opening. Thus they are confined to a narrow beam. This beam is now made to strike a thin gold foil G. A florescent screen was placed behind the gold foil. The α particles passing through gold foil produce flashes on the fluorescent screen. Rutherford observed the following points in his experiment: (i) Most of the a particles went straight trough the gold foil and produced flashes on the screen as if there were nothing inside the gold foil. (ii) Few particles collided with the atoms of the foil which have scattered or deflected through considerable large angles. (iii) Few particles even turned back towards source it self. It was a surprising result. The small angle scattering could be expected on the basis of Thomson model but the large angle scattering goes against this model. The reason is that the electrons due to their small mass are not in a position to bring about large deflection of α particles should be subjected to very feeble force of repulsion in different directions by the charges spread throughout the mass. In this way the net deviation of α- particle should be negligible and small. Thus no large scale scattering is possible in accordance with Thomson s model. The detailed experiments by Geiger and Marsden showed that large angle scattering can be expected by assuming that a massive positive point charge exists at the centre of each gold atom as shown in Fig 5.3 (a) Fig 5.3(a) α Particles Scattering 5.4 Rutherford s Nuclear Atomic Model Rutherford, in 1911, proposed a new type of model of the atom. According to this model the positive charge of the atom. Instead of being uniformly distributed throughout a sphere of atomic dimensions, is concentrated in a very small volume (less than cm in diameter) at its centre, This central core, now called nucleus, is surrounded by cloud of electrons which makes the entire atom electrically neutral. On the basis of Rutherford s nuclear atomic model the large angle scattering of positively charged &alpha- particle can be easily explained. The reason of the large angles scattering is the mutual repulsion (as per Coulomb s law) between the &alpha-particles and the concentrated positive charge on the nucleus i.e., the &alpha particle experienced a repulsive force which

4 continues to increase according to inverse square of the distance between α-particle and nucleus. When α- particle approaches the positively charged nucleus, it experiences a great repulsive force and gets deflected to a much greater angle. The &alpha-particles are sufficiently away from the nucleus. They pass through the empty space with no appreciable deviation as most of the volume of the atom is empty. Here the force between electron and &alpha-particle is ignored. Drawbacks of Rutherford s nuclear atomic model. Following are the two drawbacks of Rutherford s model: (i) The stability of the atom as a whole, and (ii) Distribution of electrons outside the nucleus. The stability of the atom as a whole. If the electrons surrounding the nucleus are at rest, the equilibrium conditions of the system cannot be reached by electrostatic forces between positively charged nucleus and negatively charged electrons outside the nucleus. For example, consider the case of a helium atom whose nucleus carries a charge +2e with two electrons outside the nucleus each having a charge e. suppose the two electrons are symmetrically placed at a distance r from the nucleus as shown in Fig 5.4(a) Fig 5.4(a) e x 2e Force of attraction between nucleus and each electron = 4π ε0 r2 e x e Force of repulsion between two electrons = - 4π ε0 (2r) This shows that the attractive force is eight times greater than the repulsive force. As the two forces do not balance each other, the electrons will fall into the nucleus thereby destroying the stability of the atom. 2 To overcome the difficulty of stability of atom, Rutherford proposed that electrons might be assumed to revolve round the nucleus with such a speed that the outward centrifugal force balances the net electrostatic attraction towards nucleus. In this way Rutherford introduced the planetary model for atom just like a solar system in which planets revolved around the central heavy sun. Thus for an electron having a charge e and mass m revolving in an orbit of radius r with velocity v around nucleus with charge +Z, we have,

5 orbits. For this expression, it can be concluded that an electron can rotate in an infinite number of The Rutherford s ideas of the revolving electron lead to a serious difficulty from the point of view of electromagnetic theory. According to electromagnetic theory a revolving electron should radiated energy continuously. This energy may be come from the atomic system itself. Due to continuous loss of energy, the electron will gradually approach the nucleus by spiral path as shown Fig 5.4 (b) and Fig 5.4 (b) finally fall to the nucleus. Hence the revolving idea of the electron destroys the stability of the atom. i.e., the very purpose for which it was postulated. Obviously, either the Rutherford s nuclear atom model with revolving electrons is defective or the classical electromagnetic theory fails to explain the present case. This problem was actually solved by Niel s Bohr in 1913 by admitting the failure of classical theory and by applying the quantum theory to the Rutherford nuclear atom with revolving electrons. 5.5 Bohr s Atomic Model Bohr developed his theory of atomic structure by retaining the essential features of Rutherford s planetary model are listed below. (i) The atom consists of a central positively charged hard core called the nucleus, and (ii) The electrons revolving round their nucleus in circular orbit the out ward centrifugal force is balanced by the net electrostatic attraction towards the nucleus. However, using the Planck s quantum theory, Bohr made the following assumptions : (iii) He assumed that an electron in the fielded of nucleus is not capable of mobbing along each of the paths that were possible on the basis of classical theory but the electron can move along one of the discrete sets of allowed paths. Only those orbits are possible for which the orbital as angular momentum of the electron is equal to an integral multiple of h/2π. hence As the momentum of revolving electron is mv and its moment about then nucleus is m v r,

6 h m v r = n - 2π Where n = 1, 2, 3, for first, second and third orbits respectively and h is Plank s constant, such orbits are known as stationary orbits. (iv) No energy is radiated by the electron as long as it remains in its definite or stationary orbit. Thus the permitted orbits are non-radiating paths of electrons. (v) The radiation of energy takes place only when an electron jumps from one permitted orbit of higher energy to another permitted orbit of lower energy. The difference of energies is radiated and must be a quantum of energy hv, i.e., E2 E1 = hv. 5.6 Bohr s Theory of Hydrogen atom Bohr s theory of hydrogen atom is based on the following assumptions: (i) An electron in an atom moves in a circular orbit about the nucleus under the influence of Coulomb s force of attraction between the electron and nucleus. As the atom as a whole is stable the Coulombian force of attraction is balanced by Newtonian centrifugal force Fig 5.6, i.e., Fig 5.6 (ii) Only those orbits are possible for which the orbital angular momentum of the electron is equal to an integral multiple of h/2π i.e., Where h is Plank s constant. (iii) The electron moving in such allowed orbit does not radiate electromagnetic radiation. Thus the total energy of the electron revolving in any one of the so many stationary orbits remains constant. (iv) Electromagnetic radiations are emitted if an electron jumps from stationary orbit of

7 higher energy E2 to another stationary orbit of lower energy E1. The frequency hv of the emitted radiation is related by the equation E2 E1 = hv. Now we shall calculate some quantities concerning the Bohr s orbit. Therefore, Bohr s Theory of Hydrogen atom confirms nucleus energy level is there. So energy level maintain by nucleus. 6.1 Introduction Periodic Properties Towards the 19th century it was realised as a result of the discovery of the electron and the phenomenon of radio-activity that atom is not invisible but consists of a number of sub-atomic particles. In fact it has a complex structure of its own on which considerable light has been thrown as a result of research work carried out during the last sixty years. In the atoms main particles are electron, protons and neutron. Atoms are unit of the substance and element is nothing but the single atom Size of Atoms and ions According to the wave-mechanical picture of the atom, since the probability of finding electrons in an atom or ion never becomes zero even at very large distances from the nucleus, the idea that atom and ions have definite size is quite untenable. Thus to define the atomic radius is quite arbitrary and due to this vagueness in concept, and number of radii have been defined for an atom. Three Types of radii are commonly used. These are given below: 6.2 Atomic Radius The distance between the nucleus and outer most shell of an atom or half of the distance between two consecutive nucleus is called Atomic Radius. These distances are calculated in three

8 ways as follows. The distance between two nucleus in d. Therefore, Atomic Radius is r = d/2. This is measured in many ways based materiel type. Namely (a) Covalent bond, (b)the crystal Radius (Atomic or Metallic radius). (c) The van der Waals Radius or Collision Radius. Atomic radius -10 measured in A = Ionic Radii Ionic radius may be defined as the distance between the nucleus of an ion and the point up to which the nucleus has influence on its electron cloud is called Ionic radius Determination of Ionic radii. The following two methods are used. (i) Lande s Method, (ii) Pauling s Method measures based on crystal types. 6.4 Ionisation Potentials or Ionisation Energies We know that in an atomic system when electrons on the outer most orbit absorb energy, they are raised to higher energy levels and give rise to the corresponding spectral lines. If this process of absorption of energy is continued, a stage comes when the electron goes completely out of the influence of the nucleus and its removal results in the formation of a positive ion, The amount of energy need to effect such a removal of the most loosely bound or outermost electrons from an isolated gaseous atom of an element in its lowest energy state is called the Ionisation Potential or ionization energy, Thus: M(g) + Ionisation energy M+(g) + e-(g) Gaseous Natural Gaseous Removed infinite away atom cation from the nucleus This energy is usually measured in electron volts (ev) or kilocalories (Kcal). The electrons are removed in stages one by one; the energy corresponding to the removal of the first electron is

9 known as the first ionization potential. Then when one more electron is removed, we have second and similarly stepwise we have third and fourth etc, ionization potential. Thus: M(g) + First ionisation potential, (I2) M+(g) + e-(g) M(g) + Second ionisation potential, (I2) M2+(g) + e-(g) The values of first ionization potentials (in ev) Valenc Ionisation Potential ev No e At Symbo l Config uratio I1 I2 I3 I4 I5 I6 I7 I8 n 1 H 1s He 1 s Li 2s Be 2s B 2s2 2p C 2s2 2p N 2s2 2p O 2s2 2p F 2s2 2p Ne 2s2 2p It would be observed from the Table show that I2 is invariably higher than I1, I3 is higher than I2 and so on i.e., I1< I2 < I3. The reason is that it is more difficult to remove an electron from a positively charged particle than from a neutral particle. 6.5 Electron A nities In ionization potentials, we express the energy supplied to the atom to knock out successively one, two or more electrons, and thereby produce ions. In the concept of electron affinities, we refer to a reverse process. Here we have an energy associated with the addition of an electron (or even more electrons in succession) to a neutral atom. The electron affinity of an element is the amount of energy released when an electron is added

10 to an isolated neutral gaseous atom in its lowest energy sate (i.e., ground state) to form an anion. Electron affinities in ev of elements. Search RECENT POSTS Hello world! RECENT COMMENTS A WordPress Commenter on Hello world! ARCHIVES December 2018 CATEGORIES Uncategorized META Log in Entries RSS Comments RSS WordPress.org Basics Science Science Introduction

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