In 1859, Julius Plucker started the study of the conduction of electricity through gases at low pressure in the discharge tube. William Crooke, J. Perrin, J. J. Thomson did a further investigation in this field. In 1885, Sir William Crookes carried out several experiments to study the behaviour of heated metal in a vacuum. He found that cathode produces a stream of radiation, which could cause a glow in gases at low pressure. He called these radiations coming out of the cathode as cathode rays. Then he studied the behaviour in the discharge tube. At atmospheric pressure or at a higher pressure no electric current flows through the tube because gases are a poor conductor of electricity.

Discharge Tube:

The Discharge tube is a glass tube used to study the flow of electricity through gases at low pressures. The discharge tube consists of a thick-walled glass tube about 40 to 50 cm in length and about 3 to 4 cm in diameter. The tube is closed at both ends and has a side tube connected to a vacuum pump and a pressure gauge. Electrodes are disc-shaped made up of aluminium are sealed in the tube at two ends. A high D.C. voltage of 10,000 to 15,000 volts can be applied between the electrodes. The tube is filled with air or any other gas in which the electric discharge at different low pressures is to be studied.

As the vacuum pump is operated the pressure of the gas in the discharge tube is gradually reduced and the discharge goes through different stages. The pressure can be read from the pressure gauge and can be made constant. This study of the discharge can be done at different pressures. If discharge tube is highly evacuated so that the pressure in the tube is of the order of 0.01 mm of mercury, certain rays are emitted by the cathode in the form of bluish streamers.  These rays are called cathode rays.

Various Stages of the Discharge:

When the pressure of the gas is equal to the atmospheric pressure the resistance of the gas between the electrodes very high. Therefore, a very high voltage of the order of 50,000 V to 1,00,000 V is necessary to pass the discharge through the tube. This discharge is in the form of a series of white sparks travelling along an irregular path through the tube. The sparks are accompanied by a cracking noise.

When the pressure is reduced to about 10 mm of mercury and a potential difference of about 10,000 V is applied between the electrodes, the discharge occurs in the form of coloured streaks or streamers travelling from the cathode to the anode. If the discharge tube contains air the discharge is a pink colour.

When the pressure is reduced to about 5mm of mercury the discharge widens until it fills the entire tube. It is called the positive column. This type of discharge is called Geissler discharge.

When the pressure decreases to about 3 mm of mercury the positive column gets detached from the cathode where a bluish glow called the negative glow is seen. The space between the positive column and the negative glow is dark and is called Faraday’s dark space.

When the pressure drops to 1 mm of mercury the negative glow gets separated from the cathode glow. A dark space appears between the cathode glow and negative glow. It is called Crooke’s dark space. The positive column gets reduced in size and dark space increases in size.

As the pressure is reduced to about 0.05mm of mercury the negative glow shift towards the anode both Crooke’s dark space and Faraday’s dark space increases in size and the positive column is divided into a number of luminous discs separated from each other by dark spaces. These discs are called striation.

As the pressure is reduced further the positive column goes on decreasing in size and it finally disappears the negative column moves toward the right and grows with fainter and Crooke’s dark space increases in size. When the pressure is 0.01mm of mercury, Crooke’s dark space fills the entire tube. At this stage, the walls glow with a greenish fluorescence. This is due to the impact of some invisible rays emitted from the cathode. These rays are called cathode rays.

If the pressure is reduced stills further the tube becomes non-conducting.

Characteristics of Cathode Rays:

• The cathode rays are emitted normally from the surface of the cathode irrespective of the position of the anode.
• They travel at a speed of about 10⁷ to 10⁹ m/s.
• They travel in straight lines and cast sharp shadows of objects in their paths.

• They are able to produce fluorescence and phosphorescence in many substances when they fall on them. The colour of fluorescence depends upon the nature of substance e.g. Willemite emits green colour. Alumina emits red colours.
• They produce a blackening on a photographic plate when they are incident on it.
• On placing a light paddle wheel in the path of cathode rays in a discharge tube, the blades of paddle wheel rotate. It shows that cathode rays constitute particles.

• They consist of negatively charged particles.
• They generate heat when they strike a target.
• They are deflected by electric and magnetic fields and the direction of deflection shows that they are negatively charged particles.
• They can ionise the gases through which they pass
• They produce heat energy when they collide with the matter. It is due to the kinetic energy possessed by the cathode rays.
• They produce X – rays when stopped by a target.
• The nature of cathode rays does not depend on the nature of the gas and the material of the cathode used in the discharge tube.

Effect of Electric Field on Cathode Rays:

When an electric field is applied to the path of cathode rays, they are deflected towards the positive plate of the electric field. It shows that cathode rays are made up of negatively charged particles.

Effect of Magnetic Field on Cathode Rays:

When a magnetic field is applied to the path of cathode rays, they are deflected. It shows that cathode rays are made up of negatively charged particles.

Uses of Cathode Rays:

• Cathode rays tubes have been developed into C. R. O. (Cathode Ray Oscillators) which has wide application in electronics. They are also used as TV tubes.
• They are used to find the ratio of charge to mass (e/m) of the electrons.
• They are used to produce X – rays.
• They are used in electron microscopes which are used for a magnifying minute object to the extent that detail of the object can be studied.

Importance of Cathode rays in the Study of Atomic Structure:

They produce heat energy when they collide with the matter. It is due to the kinetic energy possessed by the cathode rays. Cathode rays also rotate the paddlewheel in their path. They move from the negative electrode to the positive electrode. Which clearly indicate that the cathode ray consists of a negatively charged particle. Thomson called these particles as negatrons. Stoney changed this name to electrons.

Science > Physics > Cathode Rays and X- Rays > Cathode Rays

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The origin of spectral lines in the hydrogen atom (Hydrogen Spectrum) can be explained on the basis of Bohr’s theory. The hydrogen atom is said to be stable when the electron present in it revolves around the nucleus in the first orbit having the principal quantum number n = 1. This orbit is called the ground state.

The electron gains energy from the surrounding and jumps into a higher orbit with principal quantum number n = 2, 3, 4, 5, ….. These higher orbits are called excited states.  When electrons start revolving in the excited state the atom becomes unstable. To acquire stability the electron jumps from the higher orbit to lower orbit by the emission of the energy of value hν. Where ν is the frequency of radiation energy or radiation photon. This radiation is emitted in the form of spectral lines.

The Expression for the Wavelength of a line in the Hydrogen Spectrum:

Let E_n and E_p be the energies of an electron in the n^th and p^th orbits respectively (n > p) So when an electron takes a jump from the n^th orbit to the p^th orbit energy will be radiated in the form of a photon or quantum such that

E_n –  E_p = hν  ………… (1)

where ν is the frequency of radiation, h = Planck’s constant

This formula is called Bohr’s formula of spectral lines.

The wavelength λ obtained is characteristic wavelength due to jumping of the electron from n^th orbit to p^th orbit. We get different series of spectral lines due to the transition of the electron from different outer orbits to fixed inner orbit.

Energy Level Diagram for Hydrogen Atom:

Energy level diagrams indicate us the different series of lines observed in a spectrum of the hydrogen atom. The horizontal lines of the diagram indicate different energy levels. The vertical lines indicate the transition of an electron from a higher energy level to a lower energy level.

It is very important that as indicated in the diagram each transition corresponds to a definite characteristic wavelength. Thus different transitions give different series of lines. Different Series obtained are a) Lyman series, b)  Balmer series, c)  Paschen series, d)  Brackett series,  e)  Pfund series and f) Henry series

Different Series in Hydrogen Spectrum:

Lyman Series:

If the transition of electron takes place from any higher orbit (principal quantum number =  2, 3, 4,…….) to the first orbit (principal quantum number  = 1). We get a Lyman series of the hydrogen atom. It is obtained in the ultraviolet region.

This formula gives a wavelength of lines in the Lyman series of the hydrogen spectrum. Different lines of Lyman series are

• α line of Lyman series  p = 1 and n = 2
• α line of Lyman series  p = 1 and n = 3
• γ line of Lyman series  p = 1 and n = 4
• the longest line of Lyman series  p = 1 and n = 2
• the shortest line of Lyman series p = 1 and n = ∞

Balmer Series:

If the transition of electron takes place from any higher orbit (principal quantum number = 3, 4, 5, …) to the second orbit (principal quantum number = 2). We get Balmer series of the hydrogen atom. It is obtained in the visible region.

This formula gives a wavelength of lines in the Balmer series of the hydrogen spectrum. Different lines of Balmer series area l

• α line of Balmer series  p = 2 and n = 3
• β line of Balmer series  p = 2 and n = 4
• γ line of Balmer series  p = 2 and n = 5
• the longest line of Balmer series  p = 2 and n = 3
• the shortest line of Balmer series p = 2 and n = ∞

Paschen Series:

If the transition of electron takes place from any higher orbit (principal quantum number = 4, 5, 6, …) to the third orbit (principal quantum number = 3). We get Paschen series of the hydrogen atom. It is obtained in the infrared region.

This formula gives a wavelength of lines in the Paschen series of the hydrogen spectrum.

Brackett Series:

If the transition of electron takes place from any higher orbit (principal quantum number = 5, 6, 7, …) to the fourth orbit (principal quantum number = 4). We get the Brackett series of the hydrogen atom. It is obtained in the far-infrared region.

This formula gives a wavelength of lines in Brackett series of the hydrogen spectrum

Pfund Series:

If the transition of electron takes place from any higher orbit (principal quantum number = 6,7, 8, …….) to the fifth orbit (principal quantum number = 5). We get Pfund series of the hydrogen atom. It is obtained in the far-infrared region.

This formula gives a wavelength of lines in the Pfund series of the hydrogen spectrum

Notes:  Shortest wavelength is called series limit

Continuous or Characteristic X-rays:

Characteristic x-rays are emitted from heavy elements when their electrons make transitions between the lower atomic energy levels. The characteristic x-ray emission which is shown as two sharp peaks in the illustration at left occurs when vacancies are produced in the n=1 or K-shell of the atom and electrons drop down from above to fill the gap.

The x-rays produced by transitions from the n=2 to n=1 levels are called K-alpha x-rays, and those for the n=3 to n = 1 transition are called K-beta x-rays. For a particular material, the wavelength has definite value. Hence these x rays are called continuous or characteristic X-rays. The values of energy are different for different materials.

The frequencies of the characteristic x-rays can be predicted from the Bohr model. Moseley measured the frequencies of the characteristic x-rays from a large fraction of the elements of the periodic table and produced a plot of them which is now called a “Moseley plot”.

Characteristic x-rays are used for the investigation of crystal structure by x-ray diffraction. Crystal lattice dimensions may be determined with the use of Bragg’s law in a Bragg spectrometer.

Science > Physics > Atoms, Molecule, and Nuclei > Hydrogen Spectrum

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In the last article, we have studied Rutherford’s model of an atom, its merits, and demerits. In this article, we shall study Bohr’s Model of an atom, its merits, and demerits.

To overcome the limitations of Rutherford’s model of an atom, Neil Bohr put forward his theory of atom using Planck’s quantum theory. Bohr’s theory is applicable to the hydrogen atom.

Bohr’s Atomic Theory of Hydrogen Atom:

Postulate I (Postulate of Circular Orbit):

In a hydrogen atom, the electron revolves around a circular orbit around the nucleus.  The electrostatic force of attraction between the positively charged nucleus and the negativity charged electron provide necessary centripetal force for circular motion.

The Expression For the First Postulate:

Let m be the mass of an electron revolving around the nucleus in a circular orbit of radius r with a constant speed v round the nucleus.  Let – e and + e be the charges on the electron and the nucleus respectively.

By the first postulate,

Centripetal force =  Electrostatic force

Where ε_o is the electrical permittivity of free space

Postulate – II (Postulate of Selected Orbit):

The electron can revolve only in a certain selected orbit in which the angular momentum of the electron is equal to an integral multiple of nh/2π, where h is the Planck’s constant.  These orbits are called stationary or permissible orbits.  The electron does not radiate energy while revolving in these orbits.

The Expression for the Second Postulate:

Let m be the mass of an electron revolving around the nucleus in a circular orbit of radius r with a constant speed v round the nucleus.

Where, n   =    1, 2, 3………..

n = Principal quantum number

h = Planck’s constant

The integer n is called the principal quantum number and it denotes the number of the orbit.

Postulate –III (Postulate of The Origin of Spectral Lines):

When an electron takes a jump from a higher energy orbit to a lower energy orbit, energy is radiated in the form of a quantum or photon of energy hν, which is equal to the difference of energies of the electron in the two orbits.

Expression for the Third Postulate:

Let E_n and E_p be the energies of an electron in the n^th and p^th orbits respectively (n > p) So when an electron takes a jump from the n^th orbit to the p^th orbit energy will be radiated in the form of a photon or quantum such that

E_n –  E_p = hν

Where ν is the frequency of radiation.

Expression for Radius of Bohr’s Orbit:

Let m be the mass of an electron revolving in a circular orbit of radius r with a constant speed  v around the nucleus.  Let – e and + e be the charges on the electron and the nucleus, respectively.

By the first postulate,

Centripetal force =  Electrostatic force

Where ε_o is the electrical permittivity of free space

From the second postulate of Bohr’s theory

From equation (1) and (2)

This is the required expression for the radius of Bohr’s orbit. Since ε_o, h, π, m, e are constant

∴  r    ∝   n²

Thus the radius of the Bohr’s orbit of an atom is directly proportional to the square of the principal quantum number.

The Expression for Velocity of Electron in Bohr’s orbit:

From the second postulate of Bohr’s theory

Where, n = 1, 2, 3………..

n = Principal quantum number

h = Planck’s constant

This is the required expression for the velocity of the electron in Bohr’s orbit of an atom. Since ε_o, h, π, e are constant

∴   v  ∝ 1 / n

Thus the velocity of the electron in Bohr’s orbit of an atom is inversely proportional to the principal quantum number.

The Expression for Angular Velocity of Electron in Bohr’s Orbit:

Now, ε_o, m, h, π, e are constant

∴ ω  ∝ 1 / n³

Thus the angular velocity of the electron in Bohr’s orbit of an atom is inversely proportional to the cube of the principal quantum number.

Notes:

• The frequency of electron in Bohr’s orbit of an atom is inversely proportional to the cube of the principal quantum number.
• The time period of the electron in Bohr’s orbit of an atom is directly proportional to the cube of the principal quantum number.
• The centripetal acceleration of electron in Bohr’s orbit of an atom is inversely proportional to the fourth power of the principal quantum number

The Expression for Energy of Electron in Bohr’s Orbit:

Let m be the mass of an electron revolving in a circular orbit of radius r with a constant speed v around the nucleus.  Let – e  and + e be the charges on the electron and the nucleus, respectively.

The  potential energy of electron having charge,  – e is given by

The total energy of the electron is given by

The total energy of electron  =  Kinetic energy of electron + Potential energy of the electron

This is the required expression for the energy of the electron in Bohr’s orbit of an atom. Since ε_o, m, h, π, e are constant

∴   E ∝  1 / n²

Thus the energy of an electron in Bohr’s orbit of an atom is inversely proportional to the square of the principal quantum number.

The negative sign indicates that the electron is bound to the nucleus by attractive force and to remove the electron from the atom energy must be supplied to the electron to overcome the attractive force. This energy is called the binding energy of the electron.

Merits of Bohr’s Model of Hydrogen Atom:

• This theory explains the spectrum of hydrogen atom completely.
• The concept of electronic configuration i.e. the distribution of electrons in different orbits was introduced.
• This theory is capable of explaining the line spectra of elements in general.
• This theory can be used to find the ionization potential of an electron in an atom.
• The value of Rydberg can be calculated using this theory.

Demerits of Bohr’s Model of Hydrogen Atom

• Though spectra of a simple atom like hydrogen is explained by Bohr’s Theory, it fails to account for elements containing
  more than one electron.
• A line in an emission spectrum splits up into a number of closely spaced lines when the atomic source of radiation is placed in
  the magnetic field. This is known as the Zeeman Effect. Bohr Theory could not explain this.
• A line in an emission spectrum splits up into a number of closely spaced lines when the atomic source of radiation is placed in an electric field, which is known as the Stark effect. Bohr Theory has no explanation for it.

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In this article, we shall study Dalton’s atomic theory, Rutherford’s model of an atom, its merits, and demerits.

Dalton’s Atomic Theory:

This theory was proposed by English chemist John Dalton in 1808. The main propositions of the theory are as follows:

• Every element is made up of extremely small particles called an atom.
• The atoms are indivisible and they can neither be created nor be destroyed.
• Atoms of the same element resemble each other in all respects but differ from the atoms of other elements.
• Atom is the smallest unit of matter which takes part in a chemical reaction.
• When chemical compounds are formed they do so by the combination of atoms of different elements in the simple proportion of whole numbers.
• Atoms of different elements may combine in more than one proportion to form different compounds.

Fundamental Particles of an Atom:

Electrons:

• Electrons were discovered by British scientist Sir J.J. Thomson.in 1897 (Cathode ray tube experiment).
• The term electron was coined by Stoney.
• The charge on the electron was found by Milikan
• Specific charge ratio or charge to mass ratio or e/m ratio is obtained by J.J. Thomson.

Protons:

• Proton was discovered by E.Goldstein.
• They are located in the nucleus.
• They are positively charged particles.

Neutrons:

• Neutron was discovered by Chadwick in 1932.
• Neutrons have no charge i.e. they are electrically neutral.
• They are located in the nucleus of an atom.

Thomson’s Model of an Atom:

Thomson proposed this model in 1903. According to him, an atom consists of protons and electrons. The total number of protons is equal to the total number of electrons. Thus the net charge of the atom is zero. An atom consists of a positively charged sphere with negatively charged electrons embedded in it. This model of an atom is called the watermelon model or plum (dot) pudding model of an atom.

Geiger Marsden Experiment:

A narrow beam of alpha particles from the radioactive source was incident on a thin gold foil. The scattering of alpha particles takes place. The scattered alpha particles were detected by a detector fixed on a stand. The deviation of alpha particles from their original path is called the scattering angle. They observed that most of the alpha particles just passed through without any deviation as if there is empty space. A few alpha particles were deflected through smaller angles. A few alpha particles deviated through larger angles. This larger deflection is possible only if alpha-particles collide with heavy and positively charged particles inside the atom because like charges only repel each other.  This massive +ve charge is at the centre of the atom and called the nucleus. Very few alpha particles were rebounded i.e. they deviated through 180°. This concludes that the nucleus is very small as compared to the volume of the atom.

Rutherford’s Model of an Aom:

From the observations of the above experiment, Rutherford put forward the concept of his atomic model. The atom consists of a centrally located positively charged nucleus. The whole mass of an atom is concentrated in the nucleus. Around the nucleus, there is empty space in which the negatively charged electrons revolve in different orbits. The total positive charge of the nucleus is equal to the total negative charge on orbiting electrons. Hence atom is electrically neutral. Rutherford’s model of an atom is also called as a planetary model of an atom.

Limitations of Rutherford’s Model of an Atom:

• The negatively charged electrons revolve around the nucleus in a circular orbit, hence they possess centripetal acceleration. According to the classical theory of electromagnetism, accelerated charge radiates energy continuously. Therefore, the electron should radiate energy while going around the nucleus losing its energy continuously. therefore, it should approach nearer the nucleus while going round emitting radiations of increasing frequency and finally falling in the nucleus. Thus it should move in a spiral path and should emit a continuous spectrum and thus structure atom is not stable. Actually, the spectrum observed is the line spectrum of definite frequency, and hence a modification to Rutherford’s atom model was necessary.
• This model of an atom fails to explain the distribution of electrons in different orbit around the nucleus. According to Rutherford’s model of an atom, the atomic spectrum should be continuous. But the atomic spectrum is found to be discontinuous. Rutherford’s model fails to explain the discontinuity of the atomic spectrum.
• This model also fails to explain the line spectra of atoms, which show discrete lines, each line corresponds to a fixed frequency.

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