chp#02

• Properties of Cathode Rays

  • Effect of Electric Field
  • Charge
  • Effect of Magnetic Field
  • Fluorescence
  • Shadow Formation
  • Path
  • Momentum
  • Energy
  • Penetration
  • Ionization
  • Chemical Change
  • X-ray Production
  • Charge to mass ratio (e/m)

  
• Measurement of Charge to Mass Ratio (e/m) of Electron

  1. Introduction

  2. Experiment

  3. Calculation of e/m of electron
• Millikan's Oil Drop Experiment

  1. Introduction

  2. Construction

  3. Working

  4. Charge on Electron
• Positive-Rays or Canal Rays


  • Introduction
  • Construction
  • Working
  • Properties of Positive Rays
• Discovery of Neutrons


  • Introduction
  • Experiment by James Chadwick
  • Properties of neutrons
• The Discovery of Nucleus (Rutherford's Experiment, 1910-11)


  1. Experiment

  2. Rutherford's Conclusions (Rutherford's Atomic Model)

  
• The discovery of nucleus (Rutherford's experiment, 1910-11)
• Rutherford's experiment


  Bombarding α-particles from a radioactive element on a thin metallic foil

  
• Rutherford's atomic model


  • An atom consists of a small heavy positively charged portion called Nucleus
  • There is a negatively charged portion which surround the nucleus containing electrons called extra-nuclear portion or planetary
  • The number of protons in the nucleus is equal to the no of electrons in the planetary
  • The electrons revolve around the nucleus
  • The centripetal force is equal to the electrostatic force
  • Only a very small volume is occupied by the nucleus

  
• Drawbacks in Rutherford's atomic model
• According to the classical electromagnetic theory the revolving electron around nucleus should lose energy continuously and ultimately it should fall into the nucleus
• If electron emits energy continuously, it should form continuous spectrum. But actually line spectrum is obtained
• Bohr's atomic theory


  Another possible structure of atom proposed by Neil Bohr in 1913

  
• Bohr's atomic model


  • Electrons revolve around the nucleus in definite energy levels called orbits or shells
  • As long as an electron remain in a shell it never gains or losses energy
  • The gain or loss of energy occurs within orbits only due to jumping of electrons from one energy level to another energy level
  • Angular momentum (mvr) of an electron is equal to nh/2π
  • The angular momentum of an orbit depends upon its quantum number and it is an integral multiple of the factor h/2π

  
• Derivation of radius of an orbit of an atom

  1. Centripetal force = Electrostatic force

  2. Radius of a moving electron is inversely proportional to the square of its velocity

  3. Velocity 'v' is determined from angular momentum

  
• Radius of an orbit = n^2 * a_0, where a_0 is a constant

  
• Bohr's atomic model is applicable to one electron system
• Bohr's theory cannot explain the origin of the spectrum of multi-electrons or polyelectronic systems like He, Li, Be etc.
• Bohr's theory cannot explain the fine structure of hydrogen spectrum
• Bohr's theory cannot explain the Zeeman effect and Stark effect
• According to Heisenberg's uncertainty principle, both the exact position and velocity of electron cannot be measured simultaneously
• Schrodinger gave a wave equation for hydrogen atom and according to him, the probability of finding an electron can be ascertained

  
• Plank's quantum theory


  Proposed by Max Plank in 1900 about the nature of light
• Plank's quantum theory


  • Energy is not emitted or absorbed continuously, it is emitted or absorbed in the form of wave packets or quanta
  • The amount of energy associated with quantum of radiation is directly proportional to the frequency (ν) of radiation
  • Energy is inversely proportional to wavelength
  • Energy is directly proportional to wave number
• The display of spectral lines of hydrogen is called hydrogen spectrum
• The first spectral lines were discovered in 1887 by Lyman and Balmer
• Origin of hydrogen spectrum on the basis of Bohr's model

  1. Excitation: Electrons of hydrogen atoms are excited to high energy levels

  2. De-excitation: Excited electrons come back to lower energy levels and emit energy
• Bohr's prediction of wave number in hydrogen emission spectrum

  Wave number = R(1/n1^2 - 1/n2^2), where n1 and n2 are the energy levels
• Spectral series in hydrogen spectrum

  • Lyman series
  • Balmer series
  • Paschen series
  • Brackett series
  • Pfund series
• Lyman series lies in the ultraviolet region, Balmer series in the visible region, and Paschen, Brackett and Pfund series in the infrared region
• Wave numbers


  Values lie in the UV region of the spectrum
• Limiting line


  Difference between first level and infinite level is the ionization energy of hydrogen atom
• Lyman series


  • All lines have close values
  • Appear in the form of a group
• Balmer series


  Electrons fall back to n=2
• Paschen series


  • Electrons from higher levels fall back to n=3
  • Lines appear in IR region
• Brackett series


  • Electrons from higher levels fall back to n=4
  • Lines lie in IR region
• Pfund series


  • Electrons from higher energy levels fall back to n=5
  • Lines lie in IR region
• 1. rays


  Electromagnetic rays of very short wavelength produced when cathode rays hit a heavy metal target
• 1. rays were discovered by Wilhelm Roentgen accidentally in 1895

  
• 1. rays (name)


  Since these rays were new, and of unknown nature and origin, Wilhelm Roentgen called them X-rays (X = unknown)