Atomic structure


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Atomic structure

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Atomic structure

  1. 1. ATOMIC STRUCTURE -: made easy:- by M.,Anwar Sohail Bachelor of Science and Education Master of Science (Organic Chemistry) CHEMISTRY EDUCATOR Pelham High08/19/12 Mr. Sohails 1
  2. 2. Democritus’ atom [Hypothetical / Not based on experiments] Democritus proposed that matter is composed of tiny indivisible particles called ‘atom’ The word ‘atom’ means unable to be divided.08/19/12 Mr. Sohails 2
  3. 3. Dalton’s atomic theory (1808) [Based on experiments] Every element is made of tiny, unique particles called atoms that cannot be subdivided. Atoms of the same element are exactly alike. Atoms of different elements can join to form molecules.08/19/12 Mr. Sohails 3
  4. 4. Discovery of fundamental or subatomic particles The electrons, protons and neutrons are called fundamental particles or fundamental subatomic particles.08/19/12 Mr. Sohails 4
  5. 5. Canal Rays and Protons (1886)  Eugene Goldstein noted streams of positively charged particles in cathode rays in 1886. – Particles move in opposite direction of cathode rays. – Called “Canal Rays” because they passed through holes (channels or canals) drilled through the negative electrode.  Canal rays must be positive. – Goldstein postulated the existence of a positive fundamental particle called the “proton”.08/19/12 Mr. Sohails 5
  6. 6. Discovery of Electrons (1897) Electrons are discovered by J.J. Thompson when high voltage is applied across a sealed glass tube called the ‘discharge tube’ or CRT at very low pressure. The Discharge He found that what was called as Tube cathode rays until his time was not “rays” but “particles” travelling from cathode to anode. He called them electrons.08/19/12 Mr. Sohails 6
  7. 7. Discovery of Neutrons (1932) James Chadwick in 1932 analyzed the results of α-particle scattering on thin Be films. Chadwick recognized existence of massive neutral particles which he called neutrons. – Chadwick discovered the neutron.08/19/12 Mr. Sohails 7
  8. 8. Characteristics of subatomic particles at a glance08/19/12 Mr. Sohails 8
  9. 9. Thomson’s Atomic model(1898) (Also called Plum-pudding model) Thomson puts together the subatomic particles and comes forward with his atomic model. In the atom the mass and the positive charge is evenly distributed throughout the atom (like pudding) and the negatively charged electrons are embedded in it like the plum. He could not experimentally prove his model.08/19/12 Mr. Sohails 9
  10. 10. Rutherford’s Alpha particles scattering experiment (1911)  Rutherford bombarded Alpha particles on a very thin(0.00006cm) gold foil.  Most of the particles passed through, some deflected at large angles and 1 in 20000 deflected back to its own path.08/19/12 Mr. Sohails 10
  11. 11. Inference from Rutherford’s experiment  Almost all the Alpha particles passed through the gold foil means Most of the atom is empty space.  Some of the + charged alpha particles are deflected at large angles because there is a very tiny dense core of mass and + charge located in the atom. (Called nucleus)08/19/12 Mr. Sohails 11
  12. 12. Rutherford’s Alpha Rays Scattering Experiment Most of the Alpha Particles passed Through.Alpha OneParticles in every 20,000 deflected back on its own path.Source Gold Foil Some Deflected at large angles.08/19/12 Mr. Sohails 12
  13. 13. Rutherford’s Atomic model (1920) Based on his experiment he postulated a model. The important postulates are:  The atom is mostly hollow.  The mass and the positive charge (protons & neutrons) are located at the center at a very small portion called nucleus.  The electrons revolve around the Also called Planetary nucleus like the planets revolve model around the sun.08/19/12 Mr. Sohails 13
  14. 14. Atomic (z)# and Mass #  Atomic # is the # of protons present inside the nucleus of an atom.  It’s unique to each element therefore, its the identity of an element.  No two elements can have the same atomic number.  Elements are listed in the periodic table in the increasing order of their atomic numbers.08/19/12 Mr. Sohails 14
  15. 15. Mass # (A)  Mass # is the sum of the # of protons and neutrons present inside the nucleus of an atom. Therefore, must be a whole # not decimal.  A = p + n OR A = Z + n  The periodic table lists the average atomic mass not the mass #.  Atomic mass rounded to nearest whole number is the Mass #.  Mass # has no units.08/19/12 Mr. Sohails 15
  16. 16. Isotopes  Atoms of the same element with different mass #.  They have 1. Same atomic # 2. Same symbol 3. Same # of protons and electrons 4. Different # of neutrons & mass #08/19/12 Mr. Sohails 16
  17. 17. Isotopic Symbol Net ChargeMass # (A)Atomic # (Z) Symbol08/19/12 Mr. Sohails 17
  18. 18. The 3 Nuclie of H isotopesP P PN N NZ Z ZM M Me e e Complete Isotopic Symbol Worksheet08/19/12 Mr. Sohails 18
  19. 19. Average Atomic Mass  Weighted average of the atomic masses of all the naturally occurring isotopes of an element is called average atomic mass.  It is measured in amu (atomic mass unit)  1 amu is 1/12th of the mass of C-12 atom.08/19/12 Mr. Sohails 19
  20. 20. Calculating Average Atomic Mass (Mass of A*%) + (Mass of B*%) 100 Complete Average Atomic Mass – 1 Worksheet08/19/12 Mr. Sohails 20
  21. 21. Isotones Isotones are atoms of different elements with same # of neutrons. Examples: S – 32 and P – 31 Ca – 40 and K – 3908/19/12 Mr. Sohails 21
  22. 22. Problems with Rutherford’s model  As per the classical laws of Physics: if a particle (electron) is revolving around oppositely charged particle (positive nucleus), the revolving particle loses its energy continuously and finally falls in to the central particle. Therefore the atom should collapse.But this is not happening in nature.  If the negatively charged electron is revolving around positively charged nucleus, the atomic spectra should be a band spectrum but in nature the atomic spectrum is line spectrum.08/19/12 Mr. Sohails 22
  23. 23. Band spectrum  When white light is passed through a prism, it splits in to 7 different colors and they appear as bands of 7 colors on a film or screen. (Example in nature: Rainbow)  This is called a band spectrum. It is not a characteristic of an atomic spectrum08/19/12 Mr. Sohails 23
  24. 24. Line Spectrum or Atomic emission spectrum  When electricity is passed through a tube filled with a gas (Ex.CRT), light will be emitted. If the emitted light is passed through the prism and its image is recorded on a film it appears as ‘sharp lines on black background’. This is called line spectrum or “atomic emission spectrum” .Every element has a characteristic emission spectrum of its own.08/19/12 Mr. Sohails 24
  25. 25. Absorption spectrum  An absorption spectrum is formed by shining a beam of white light through a sample of gas. – Absorption spectra indicate the wavelengths of light that have been absorbed by the gas. – It appears as dark lines on bright background.08/19/12 Mr. Sohails 25
  26. 26. Characteristics of Light08/19/12 Mr. Sohails 26
  27. 27. Characteristics of Light Velocity (c): Distance traveled by light in 1 second. It’s a constant 3.00 x 10 8 m/s c=ν λ Wave length (λ): Distance between any two similar points on successive waves. Measured in m or nm (nano meters) 1nm = 10 – 9 m = 10 – 7 cm λ =c/ν Frequency (ν ): # of waves that cross a given point in 1 second. Measured in Hertz (Hz) or cycles per Hz second (cps) ν = c / λ cps Amplitude: Height of a crest or depth of a trough. Refers to the intensity of light.Energy (E): Energy contained in a wave. Measured in Joules (J) E = h νWhere h is Planck’s constant (6.626 x 10-34 J)08/19/12 Complete Characteristics of light Worksheet Mr. Sohails 27
  28. 28. Frequency, Wavelength & Energy relationshipsWhen frequency increases: Energy increases Wavelength decreasesWhen Wavelength increases: Energy decreases Frequency decreasesWhen amplitude decreases: intensity (brightness of light) decreases08/19/12 Mr. Sohails 28
  29. 29. Bohr’s atomic model Neils Bohr presented his atomic model retaining the basic idea of Rutherford’s model. The important postulates are: 1. Electrons revolve around the nucleus in definite, closed, circular paths called orbits. 2. Each orbit is associated with a definite amount of energy therefore also called as energy level. 3. These orbits or energy levels are numbered 1,2,3,4….. or K,L,M,N…. from inside onwards. Bigger the orbit, They are also called principal greater is the energy quantum levels, represented by associated with it. ‘n’.08/19/12 Mr. Sohails 29
  30. 30. Bohr’s model Continued:- 4. More than one energy levels are possible for an electron. However, as long as an electron is in an energy level its energy remains constant. 5. When an electron gains energy it jumps from lower energy level to higher. 6. When it jumps back from higher energy level to lower, it loses energy in 7. The energy released ( ∆E )can be the form of light. calculated by: Where ‘h’ is Planck’s constant, ‘v’ is the frequency of light emitted.08/19/12 Mr. Sohails 30
  31. 31. Modern Model or Wave Mechanical Model or Quantum Mechanical Model of the atom08/19/12 Mr. Sohails 31
  32. 32. Particle nature of light(1901) Max Plancks Quantum Theory  Max Planck studied the radiation emitted by various objects at high temperatures and came to a conclusion that:  Light is absorbed or emitted by matter in the form of discrete packets of energy. Each energy packet is called Photon and the energy it holds is termed Quantum.  The energy contained in each PHOTON of light is directly proportional to its frequency and can be calculated by the equation: E=hv  Where h is Planck’s constant (6.626 x 10-34 J)  Planck’s quantum theory helped understanding the phenomenon of Photoelectric effect (ejection of electrons from the surface of metal when light of a certain frequency08/19/12 falls on it.) Mr. Sohails 32
  33. 33. DeBroglie’s Dual nature of electron  Based on Planck’s quantum theory and Bohr’s quantized orbits, De Broglie suggested that:  every moving particle exhibits a wave nature so also the electrons.  electrons behave more like waves on a vibrating string than like particles.  The wave length of any particle wave can be calculated by the equation: λ =h/mv (De Broglie’s equation)08/19/12 Mr. Sohails 33
  34. 34. Heisenberg’s uncertainty principle Its impossible to find out both the position and the speed of an electron accurately at the same time. It is because to locate an electron, light, having wave length shorter than the size of an electron should fall on it and be reflected. Light with such a short wave length will have very high energy, which will energize the electron. Therefore, its velocity around the nucleus will increase. 08/19/12 Mr. Sohails 34
  35. 35. Splitting of Bohr’s Spectral lines Introduction to Orbitals/Quantum #s With a regular spectroscopePrincipal Quantum # =(n) sharp principal diffused fundamentalUnder High resolutionspectroscopeAzimathul Quantum # =(l) Under Magnetic FieldMagnetic Quantum # =(m)Under Electric FieldSpin Quantum # =(s) 08/19/12 Mr. Sohails 35
  36. 36. Principle Quantum # Bohr’s spectral lines. = Energy Level or Bohr’s atomic orbits. Values: any # 1,2,3,….so on from inside outwards. Value can’t be zero The total # of electrons that can be accommodated in an energy level is given by 2n2 where n is the Principle Quantum # or energy level.08/19/12 Mr. Sohails 36
  37. 37. Azimathul Quantum # l Splitting of Bohr’s spectral lines under high resolution spectroscope. = sub energy Levels or orbitals. Values: 0 to n – 1 Total # of “l ” values will be equal to n Ex: For n=4 l values will be: 0,1,2,3 Figure out the l values for 1st, 2nd & 3rd energy levels l = 0 : s orbital l = 1 : p orbital l = 2 : d orbital l = 3 : p orbital 08/19/12 Mr. Sohails 37
  38. 38. Magnetic Quantum # (m) Splitting of high resolution lines in magnetic field. Also called angular momentum Q #. # of m values for each l value = 2l +1(How many?) Value ranges from – l to 0 to + l(What are they?)Practice:08/19/12 Mr. Sohails 38
  39. 39. Figure out the m values1. How many m values are there for s orbital?2. What are they?3. How many m values are there for p orbital?4. What are they?5. How many m values are there for d orbital?6. What are they?7. How many m values are there for f orbital?8. What are they? 08/19/12 Mr. Sohails 39
  40. 40. Spin Quantum # (s) Indicates electron spin in the orbital or electron cloud (either clockwise or counterclockwise) Values: for each m value there are 2 s values; they are +1/2 and -1/2 This indicates that in each m there are 2 electrons one spinning clockwise and the other counterclockwise.08/19/12 Mr. Sohails 40
  41. 41. Atomic Orbitals As it is impossible to locate an electron’s exact position at a given time, therefore: The area around the nucleus where the probability of finding an electron is maximum is called an orbital. There are 4 atomic orbitals discovered so far. They are s,p,d,f The s orbital is spherical shaped electron cloud, the p orbital is a dumbbell shaped electron cloud and the d orbital is a double dumbbell and the f orbital is an 8 lobbed dumbbell.08/19/12 Mr. Sohails 41
  42. 42. Orbitals and Electrons Energy level Types of # of Orbitals Electrons Total # of electrons in thePrincipal Quantum Orbitals Magnetic Quantum Spin Quantum # (s) energy level # (n) Azimuthal Quantum # (m) # (l)1 s 1 2 22 s 1 2 8 p 3 63 s 1 2 18 p 3 6 d 5 104 s 1 2 32 p 3 6 d 5 1008/19/12 f Mr. Sohails 7 14 42
  43. 43. s Orbitals There is one s Orbitals in each energy level Each one can hold 2 electrons.08/19/12 Mr. Sohails 43
  44. 44. p Orbitals There are 3 p Orbitals in each energy level (from 2nd energy level on wards) Each one can hold 2 electrons therefore 6 electrons in each p sublevel08/19/12 Mr. Sohails 44
  45. 45. d Orbitals There are 5 p Orbitals in each energy level (from 3rd energy level on wards)Each one can hold 2 electrons therefore 10 electrons in each d sublevel08/19/12 Mr. Sohails 45
  46. 46. What does the modern atom look like?08/19/12 Mr. Sohails 46
  47. 47. Electron Configuration Arrangement of electrons in the various orbitals of the atom of an element is called electron configuration. It is governed by 3 laws: Aufbau Principle: Electrons occupy the lowest energy orbital available. Pauli’s exclusion principle: No more than two electrons in each orbital. Hunds rule: When degenerate orbitals are available, Pairing of electrons takes place after half filling.08/19/12 Mr. Sohails 47
  48. 48. Atomic Orbital Energy Diagram08/19/12 Mr. Sohails 48
  49. 49. End of Pr esentation Remember the Atomic Structure. It is the key to learn Chemistry08/19/12 Mr. Sohails 49