Chapter 2.
Structure of Atom
Some Important Points and Terms of the Chapter
1. The word 'atom' has been derived from the Greek word a-tomio' which means 'uncutable 'or non-divisible'
2. J.J.Thomson, in 1898, proposed that an atom possesses a spherical shape (radius approximately 10-10 m) in which the positive charge is uniformly distributed. The electrons are embedded into it in such a manner as to give the most stable electrostatic arrangement (Fig. 2.4, NCERT Page 5) Many different names are given to this model, for example, plum pudding, raisin pudding or watermelon.
3. Rutherford's Nuclear Model of Atom:
a) Most of the space in the atom is empty as most of the alpha particles passed through the foil undeflected
b) A few positively charged alpha particles were deflected. The deflection must be due to enormous repulsive force showing that the positive charge of the atom is not spread throughout the atom as Thomson had presumed. The positive charge has to be concentrated in a very small volume that repelled and deflected the positively charged a-particles.
e)Calculations by Rutherford showed that the volume occupied by the nucleus is negligibly small as compared to the total volume of the atom. The radius of the atom is about 10-¹⁰ m, while that of nucleus is 10-¹⁵m.
d) On the basis of above observations and conclusions, Rutherford proposed the nuclear model of atom (after the discovery of protons). According to this model :) The positive charge and most of the mass of the atom was densely concentrated in extremely small region. This very small portion of the atom was called nucleus by Rutherford.
(i) The nucleus is surrounded by electrons that move around the nucleus with a very high speed in circular paths called orbits. Thus, Rutherford's model of atom resembles the solar system in which the nucleus plays the role of sun and the electrons that of revolving planets.
(ii) Electrons and the nucleus are held together by electrostatic forces of attraction.
4. The number of protons present in the nucleus is equal to atomic number (Z).the nucleus is equal to atomic number (Z ).i.e. Atomic number (Z) = number of protons in the
nucleus of an atom = number of electrons in a nuetral atom.
5. Protons and neutrons present in the nucleus are collectively known as nucleons. The total
number of nucleons is termed as mass number (A) of the atom.mass number (A) = number of protons (Z)+ number of neutrons (n)
6. Isobars are the atoms with same mass number but different atomic number for example, 6C¹⁴ and
7N¹⁴. On the other hand, atoms with identical atomic number but different atomic mass number are known as Isotopes. e.g. c cH C& CCP
7. Drawbacks of Rutherford Model –
According to the electromagnetic theory of Maxwell, charged particles when accelerated should emit electromagnetic radiation (This feature does not exist for planets since they are uncharged),. Therefore, an electron in an orbit will emit radiation, the
energy carried by radiation comes from electronic motion. The orbit will thus continue to shrink.
Calculations show that it should take an electron only 10 s to spiral into the nucleus. But this
does not happen. Thus, the Rutherford model cannot explain the stability of an atom.
8. The frequency (v ), wavelength () and velocity of light (c) are related by the equation (2.5). c =
v.The other commonly used quantity specially in spectroscopy, is the wavenumber (v). It is
defined as the number of wavelengths per unit length. Its units are reciprocal of wavelength unit,
i.e., m
9. H. Hertz performed a very interesting experiment in which electrons (or electric current) were
ejected when certain metals (for example potassium, rubidium, caesium etc.) were exposed to a
beam of light. The phenomenon is called Photoelectric effect. For photoelectric effect: hv =
h+1/2 mv
(i) The energy 1s
10. Planck's quantum theory.
radiated or absorbed by a body not continuously but discontinuously in form of small packets.
(ii) Each packet is called quantum. In case of light, the quantum is called "photon'. The energy of
quantum is directly proportional to the frequency (V) of the radiation. E vE = hv, Where h
is Planck's constant. Its value is 6.625 x 10% Joule second.
11. The spectrum of radiation emitted by a substance that has absorbed energy is called an emission
spectrum. Atoms, molecules or ions that have absorbed radiation are said to be "excited". To
produce an emission spectrum, energy is supplied to a sample by heating it or irradiating it and
the wavelength (or frequency) of the radiation emited, as the sample gives up the absorbed
energy, is recorded.
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12. An absorption spectrum is like the photographic negative of an emission spectrum. A
continuum of radiation is passed through a sample which absorbs radiation of certain
wavelengths.
13. Line Spectrum of Hydrogen: When an electric discharge is passed through gaseous hydrogen,
the H molecules dissociate and the energetically excited hydrogen atoms produced emit
electromagnetic radiation of discrete frequencies. The hydrogen spectrum consists of several
series of lines named after their discoverers.
Spectral region
Series
|Ultraviolet
Lyman| 2,3..
Balmer2 3,4. Visible
Paschen 3 4,5. nfrared
Brackett4 5,6.. Infrared
Pfund 5 |6,7. Infrared
The Swedish spectroscopist, Johannes Rydberg. noted that all series of lines in the hydrogen spectrum
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could be described by the following expression
14. Bohr's Model For Hydrogen Atom
a) The electron in the hydrogen atom can move around the nucleus in a circular path of fixed
radius and energy. These paths are called orbits, stationary states or allowed energy states.
These orbits are arranged concentrically around the nucleus.
b) An electron can move only in those orbits for which its angular momentum is integral
multiple of h/2[] that is why only certain fixed orbits are allowed. The angular
momentum of an electron in a given stationary state can be expressed as in equation
m, vr =n n-1,2,3.
c) The energy of an electron in the orbit does not change with time. However, the electron
will move from a lower stationary state to a higher stationary state when required amount
of energy is absorbed by the electron or energy is emitted when electron moves from
higher stationary state to lower stationary state. The energy change does not take place
in a continuous manner.
d) The frequency of radiation absorbed or emitted when transition occurs between two
stationary states that differ in energy by AE, is given by :
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15. Bohr's theory can also be applied to the ions containing only one electron, similar to that present
in hydrogen atom. For example, He" Li*", Be" and so on. The energies of the stationary states
associated with these kinds of ions (also known as hydrogen like species) are given by the
E-2.181o"
and radii by the expression
pm
expression
16. Limitations of Bohr's Model: It fails to account for the finer details (doublet, that is two closely
spaced lines) of the hydrogen atom spectrum observed by using sophisticated spectroscopic
techniques. This model is also unable to explain the spectrum of atoms other than hydrogen, for
example, helium atom which possesses only two electrons. Further, Bohr's 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).
17. Dual Behaviour of Matter: The French physicist, de Broglie in 1924 proposed that matter, like
radiation, should also exhibit dual behaviour i.e., both particle and wavelike properties.
18. The de Broglie relation. :de Broglie relation state that the wavelength associated with a moving
object or an electron is inversely proportional to the momentum of the particle.
where p is the momentum of particle = mv.
19. Heisenberg's Uncertainty Principle. It is not possible to determine the position and velocity
simultaneously for a sub-atomic particle like electron at any given instant to an arbitary degree of
precision. Consequently, it is not possible to talk of path of the electron in which it moves. If
Ax' is uncertainty in position and AP' is uncertainty in momentum then
A.AP2
20. Orbital. It is a region or space where there is maximum probability of getting electron.
21. Quantum numbers. They are used to get complete information about electron, i.e., location,
energy. spin, etc. These quantum numbers also help to designate the electron present in an
orbital.
22. Principal quantum number. It specifies the location and energy of an electron. It is measure
of the efiective volume of the electron cloud. It is denoted by 'n'. Its possible values are 1, 2,
3,4
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23. Angular momentum quantum number. It is also called 'azimuthal quantum number". It
determines the shape of the orbital. It is denoted by T". The permitted values of are 0, 1, 2,
etc., upto n-1. For a given value of n, l=0 to n-1. eg, ifvalue of n is 4,I can have values
mr = i0+D
0, 1,2, 3. It determines angular momentumn.
24. Magnetic quantum number. It is denoted by "m' and its value depends on value of l' since
magnetism is due to angular momentum. It determines the magnetic orientation of an orbital,
i.e., the direction of orbital relative to magnetic field in which it is placed. Its permitted values
are I to + l including zero, e.g., when I = 1, then m = -1, 0, +1. It has total number of values
equal to 21 + 1.
25. Spin quantum number. It indicates, the direction in which electron revolves. Spin is
magnetic property and is also quantized. It has two permitted values + á or -. The spin
angular momentum of an electron is constant and cannot be changed.
26. (n+1) rule: The relative order of energies of various sub-shells in a multi-electron atom can be
predicted with the help of (n+l) rule (also called Bohr-Bury rule)According to this rule a sub-
shell with lower values of (n+1) has lower energy.In case two sub-shell has equal value of (n+1).
the sub-shell with lower value of n has lower energy
27. Pauli's Exclusion Principle. No two electrons in an atom can have all the four quantum
numbers same. It can also be stated as An orbital can have maximum two electrons and theymust be of opposite spin quantum numbers.
28. Aufbau principle. Electrons are filled in the various orbitals in the increasing order of thei
energies, ie., orbital having lowest energy will be filled first and the orbital having highes
energy will be filled last.Increasing energy ofatomic orbitals for multi-electron atom
ls <2s< 2p< 3s <3p <4s 3d<4p 5s <4d <5p< 6s < 4/< Sd < 6p < 7
29. Hund's rule of masimum multiplicity. No electron pairing takes place in p. d and forbital
until cach orbital in the given sub-shell coatains onc electron. e.g.. N (7) has clectroni
configuration I 2 2p, 2p, 2p, according to Hund's rule and not Ls 2 2p 2p
30. The valence electronic configurations of Cr and Cu, therefore, are 3d' 4s' and 3d"as
respectively and not 3d 4s and 3d'4s. It has been found that there is extra stability(Stability
Completely Filled and Half Filled Subshells) associated with these electronic configuration
31. Three orbitals of 2p subsbell (2p 2p, and 2p, orbital.
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