Atomic Structure

Anode Rays
The Rutherford Model
Drawback of Rutherford Model
Bohr's Model
Quantum Model

Cathode Rays (Discovery of electrons)

At ordinary pressure, gases do not conduct electricity. Though gases conduct electricity at very low pressures and when a high potential difference is applied across them. Conduction of electricity through gases is also called gaseous discharge.
The rays produced by the discharge of electricity through gases are called cathode rays .
The presence of electrons in an atom was proved by Sir J.J. Thomson .

Characteristics of Electron

(i) An electron is a negatively charged particle found in the atoms of all elements.
(ii) An electron has mass      times that of a proton (which is the positively charged particle in the atom of

an element).
(iii) An electron has the absolute charge of -1.6 x 10–19 Coulombs .

Anode Rays (Discovery of protons)

Anode rays, consisting of positively charged particles known as protons have been experimentally produced. The existence of protons in the atoms was shown by Goldstein.

Characteristics of Proton

(i) A proton is a positively charged particle .
(ii) A proton has relative mass of 1 a.m.u .
(iii) A proton has the absolute charge of 1.6 x 10–19 Coulombs but the charge of proton is opposite to that of electron.


The third fundamental particles (inside the nucleus) was proved by Chadwick in 1932 . It was called neutron .

Characteristics of Neutron

(i) A neutron is a neutral particle .
(ii) A neutron has relative mass of 1 a.m.u .


Atomic Models

The Plum Pudding Model

J.J. Thomson proposed the plum pudding model. This model proposed that atoms are blobs of a positively charged jellylike material, with electrons suspended in it like raisins in a pudding.

The Rutherford Model

In 1911, Ernst Rutherford and his co-workers carried out a series of experiments using alpha-particles . A beam of alpha-particles was directed against a thin foil of about 0.0004 cm thickness of gold, platinum, silver or copper respectively. The foil was surrounded by circular fluorescent zinc.

The following observations were made:

1. Most of the alpha-particles (nearly 99%) went straight without suffering any deflection.

2. A few of them got deflected through small angles .

3. A very few (about one in 20,000) did not pass through the foil at all but suffered large deflections (more than 90°) or even came back in more or less the direction from which they have come, i.e., a deflection of 180° .

Atomic Spectra

1. Emission spectra - When the radiation emitted from some source e.g. by passing electric discharge through a gas at low pressure or by heating some substance to high temperature etc. is passed directly through the prism and then received on the photographic plate, the spectrum obtained is called ‘Emission spectrum’ .

2. Continuous spectra - A continuous spectra results when the gas pressures are higher. Generally, solids, liquids, or dense gases emit light at all wavelengths when heated.

Hydrogen Emission and Absorption Series

Emission or absorption processes in hydrogen give rise to series, which are sequences of lines corresponding to atomic transitions, each ending or beginning with the same atomic state in hydrogen. Thus, for example, the Balmer Series involves transitions starting (for absorption) or ending (for emission) with the first excited state of hydrogen, while the Lyman Series involves transitions that start or end with the ground state of hydrogen.


Drawbacks of Rutherford Model

According to the classical electromagnetic theory , there is emission of electromagnetic radiation and loss of energy, when any charge moves around any other charge. Thus, when a negatively charged electron revolves around a positively charged nucleus it must emit radiations. As a result of this, the electron should lose energy at every turn and move closer and closer to the nucleus. The ultimate result will be that it will fall into the nucleus, thereby making the atom unstable.
Bohr made calculations and find out that an atom would collapse in 10-8 seconds . Since the atom is quite stable, it means the electrons do not fall into the nucleus, thereby this model does not explain the stability of the atom .

Bohr's Theory of Hydrogen Atom

Niels Bohr suggested that the problem about hydrogen spectrum can be solved if we can make some assumptions. According to classical theory, the frequency of the electromagnetic waves emitted by a revolving electron is equal to the frequency of revolution.


(i) Every atom consists of nucleus and suitable number of electrons revolved around the nucleus in circular orbits.

(ii) Electrons revolved only in certain non-radiating orbits called stationery orbits for which the total angular momentum is an integral multiple of h/2p where h is plank's constant.

L is the Angular momentum of the revolving electrons.

(iii) Radiation occurs when an electron jumps from one orbit to another. It is emitted when electron jumps from higher orbit to a lower orbit i.e., E2 - E1 = hf, where f is frequency of radiation.

Bohr’s model helps to calculate the energies of various stationary states in hydrogen atom. The energy (En) associated with each stationary state is called its energy level . It is given by the expression,

Where n is the quantum number and has integral values 1, 2, 3 ......
For hydrogen

De Broglie Wavelength

In 1924, the French physicist de Broglie suggested that the electron has a dual nature . In other words, electron have a dual character; they behave as particles as well as waves .

The wavelength associated with a particle of mass 'm' , moving with velocity 'v' is given by de Broglie's relation as

Heisenberg's Uncertainity Principle

According to Heisenberg’s uncertainty principle, it is not possible to measure, simultaneously, the position and the momentum of a particle with unlimited precision. Mathematically this may be expressed 

The Quantum Model

Improving on the Bohr Model, Sommerfeld, in order to account for the additional lines present in the spectra of atoms, assumed that each principal energy level contains a number of sub-levels , each of which possesses slightly different energy. The subsidiary orbits are designated s, p, d and f.

The increasing energy levels of these sub-levels are: s < p < d < f

There are four quantum numbers. These are principal quantum number , azimuthal quantum number , magnetic quantum number and spin quantum number .
Permissible values of the quantum numbers for various orbitals are mentioned in the table given below:

Shapes of Atomic Orbitals

(i) s-orbital :The shape of the cloud is the shape of the s-orbital. The cloud is not uniform but denser in the region where the probability of finding the electron is maximum .
The size of an orbital increases as the principal quantum number increases, thus a 2s-orbital is larger than 1s-orbital. A 2s-orbital differs from 1s-orbital in having a nodal surface. A nodal surface is the region in space where the probability of finding electron is zero.

(ii) p-orbital :There are three p-orbitals px, py and pz . They are dumb-bell shaped , the two levels being seperated by a nodal plane , i.e., a plane where there is no likely hood of finding the electron. The p-oritals have a marked directional character, depending on whether the px, py and pz orbital is considered.
The p-orbitals consist of two lobes with the atomic nucleus lying between them. The axis of each p-orbital is perpendicular to other two.

(ii) d-orbital : There are five d-orbitals which are directional and of same energy. The shapes of four d-orbitals resemble four leaf cloves . The fifth d-orbital loops differently.
These d-orbitals are shown below.

The Aufbau Principle

Aufbau principle states that the orbitals get filled up in an increasing order of their energies . It means that the last added electron will occupy the available orbital with the least energy.

There is a simpler method to get the order.

(i) Write a column of 's' orbitals from 1 to 8.

(ii) Write a second column for the 'p' orbitals starting at n=2. (1p is not an orbital combination allowed by quantum mechanics)

(iii) Write a column for the 'd' orbitals starting at n=3.

(iv) Write a final column for 4f and 5f. There are no elements that will need a 6f or 7f  shell to fill.

(v) Read the chart by running the diagonals starting from 1s.

Now the order of orbitals are known to fill, one thing remaining is memorizing how large each orbital is:

s orbitals have 1 possible value of m to hold 2 electrons
p orbitals have 3 possible value of m to hold 6 electrons
d orbitals have 5 possible value of m to hold 10 electrons
f orbitals have 7 possible value of m to hold 14 electrons

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