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The Development of Atomic Theory

The Development of Atomic Theory. Larry Scheffler Lincoln High School Portland OR. 1. The Atom. The term atom is derived from the Greek word atomos (atomos) meaning invisible Democritius (470-370 BC ) suggested that all matter was made up of invisible particles called atoms. 2.

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The Development of Atomic Theory

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  1. The Development of Atomic Theory Larry Scheffler Lincoln High School Portland OR 1

  2. The Atom • The term atom is derived from the Greek wordatomos(atomos) meaning invisible • Democritius (470-370 BC ) suggested that all matter was made up of invisible particles called atoms 2

  3. Law of Constant Composition A compound always contains atoms of two or More elements combined in definite proportions by mass Example: Water H2O always contains 8 grams of oxygen to 1 gram of hydrogen 3

  4. Law of Multiple Proportions Atoms of two or more elements may combine in different ratios to produce more than one compound. Examples: NO NO2 N2O N2O5 4

  5. Dalton’s Atomic Theory • All elements are composed of indivisible and indestructible particles called atoms. • Atoms of the same element are exactly alike, They have the same masses. • Atoms of different elements have different masses. • Atoms combine to form compounds in small whole number ratios.. 5

  6. Objections to Dalton’sAtomic Theory • Atoms are not indivisible. They are • composed of subatomic particles. • Not all atoms of a particular element • have exactly the same mass. • Some nuclear transformations • alter (destroy) atoms 6

  7. Crookes Experiment Crookes found that passing an electrical current through a gas at very low pressure caused the gas to glow. Putting a magnet next to the beam caused it to be deflected. 7

  8. The Electron • The electron was the first subatomic particle to be identified. • In 1897 J.J Thomson used a cathode ray tube to establish the presence of a charged particle known as the electron • Thomson established the charge to mass ratio E/m = 1.76 x 108 coulombs/gram 8

  9. A Cathode Ray Tube Thomson found that an electrical field would also deflect an electron beam. He surmised that the ratio of charge to mass is constant.

  10. Thomson’s Charge to Mass Ratio E/m = 1.76 x 108 coulombs/gram

  11. Thomsen’s Plum Pudding Model Thompson proposed that an atom was made up of electrons scattered unevenly through out an elastic sphere. These charges were surrounded by a sea of positive charge to balance the electron's charge like plums surrounded by pudding. This early model of the atom was called The Plum Pudding Model. A more contemporary American label might be the “chocolate chip cookie” model 11

  12. Millikan’s Experiment • By varying the charge on the plates, Millikan found that he could suspend the oil drops or make them levitate. 12

  13. Millikan’s Experiment Millikan used his data tomeasure the charge of an electron and then to calculate the mass of the electron from Thomson’s charge to mass ratio. Given the charge = 1.60 x 10-19 coulomb and the ratio of E/m = 1.76 x 108 coulombs/gram it is possible to calculate the mass Mass = 9.11 x 10-28 gram 13

  14. Protons First observed by E. Goldstein in 1896 J.J. Thomson established the presence of positive charges. The mass of the proton is 1.673 x 10-24 grams 14

  15. Rutherford’s Experiment Rutherford oversawGeiger and Marsden carrying out his famous experiment. They fired high speed alpha particles (Helium nuclei) at a piece of gold foil which was only a few atoms thick. They found that although most of them passed through. About 1 in 10,000 hit and were deflected 1910 Ernest Rutherford 15

  16. Rutherford’s Experiment 16

  17. Rutherford’s Experiment 17

  18. Rutherford’s Experiment By studying this pattern, Rutherford concluded that atoms have a very dense nucleus, but there are mostly empty space. 18

  19. Subatomic Particles The diameter of a single atom ranges From 0.1 to 0.5 nm. (1 nm = 10-9 m). Within the atom are smaller particles: Electrons Protons Neutrons 19

  20. Neutrons Discovered by James Chadwick in 1932 Slightly heavier than a proton Mass of a neutron = 1.675 x 10-24 grams 20

  21. The Bohr Model Niels Bohr proposed the Planetary Model in 1913. Electrons move in definite orbits around the nucleus like planets moving around the nucleus. Bohr proposed that each electron moves in a specific energy level. 21

  22. Aspects of the Bohr Model • Bohr put together Balmer’s and Plank’s discoveries to form a new atomic model • In Bohr’s model: • Electrons can orbit only at certain allowed distances from the nucleus. • Electrons that are further away from the nucleus have higher energy levels (explaining the faults with Rutherford’s model). 22

  23. The Electromagnetic Spectrum

  24. Wave Characteristics Energy of a wave E = hn Frequency = n = number of peaks per unit of time Speed of light c = nl

  25. Emission Spectra 25

  26. Flame Tests

  27. According to Bohr Atoms radiate energy whenever an electron jumps from a higher-energy orbit to a lower-energy orbit. Also, an atom absorbs energy when an electron gets boosted from a low-energy orbit to a high-energy orbit. 27

  28. Problems with the Bohr Model • The Bohr model provided a model that gave precise results for simple atoms like hydrogen. • Using the Bohr model precise energies could be calculated for energy level transitions in hydrogen. • Unfortunately these calculations did not work for atoms with more than 1 electron. 28

  29. Weakness of the Bohr Model • According to the Bohr model electrons could be found in orbitals with distinct energies. • When the data for energies measured using spectral methods where compared to the values predicted by the Rydberg equation, they were accurate only for hydrogen. • By the 1920s, further experiments showed that Bohr's model of the atom had some difficulties. Bohr's atom seemed too simple to describe the heavier elements. 29

  30. Modern View of the Atom The wave mechanical model for the atom was developed to answer some of the objections that were raised about the Bohr model. It is based on the work of a number of scientists and evolved over a period of time The quantum theorists such as Maxwell Planck suggested that energy consists of small particles known as photons. These photons can have only discreet energies Maxwell Planck 30

  31. Modern View of the Atom Albert Einsteindemonstrated the equivalence of matter and energy. Hence matter and energy in Einstein’s theory were not different entities but different expressions of the same thing Einstein then proposed the equivalence of Matter and Energy given by his famous equation E = mc2 31

  32. Modern View of the Atom Louis de Broglie suggested that if energy could be thought of as having particle properties, perhaps matter could be thought of as having wave like characteristics Louis de Broglie 32

  33. Modern View of the Atom Louis de Broglie proposed that an electron is not just a particle but it also has wave characteristics. E = mc2 = hn 33

  34. Modern View of the Atom The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa. --Heisenberg, Uncertainty paper, 1927 Heisenberg proposed that it was impossible to know the location and the momentum of a high speed particle such as an electron. 34

  35. Modern View of the Atom The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa. --Werner Heisenberg, Uncertainty paper, 1927 The atom cannot be defined as a solar system with discreet orbits for the electrons. The best that we could do was define the probability of finding an electron in a particular location. 35

  36. Modern View of the Atom Edwin Schroedinger proposed that the electron is really a wave. It only exists when we identify its location. Therefore the electrons are best thought of probability distributions rather than discreet particles. 36

  37. Modern View of the Atom The modern view of the atom suggests that the atom is more like a cloud. Atomic orbitals around the nucleus define the places where electrons are most likely to be found. 37

  38. Wave Mechanical Model The location of the electron in a hydrogen atom is a probability distribution. 38

  39. Progression of Atomic Models Our view of the atom has changed over time 39

  40. ATOMIC STRUCTURE Particle Charge Mass proton + charge 1 neutron No charge 1 electron - charge 0 40

  41. ATOMIC NUMBER AND MASS NUMBER Mass Number He 4 the number of protons and neutrons in an atom 2 Atomic Number the number of protons in an atom Number of electrons = Number of protons in a neutral atom 41

  42. Atomic Mass The atomic mass of an atom is a relative number that is used to compare the mass of atoms. An atomic mass unit is defined as 1/12 of the mass of an atom of carbon 12. The atomic masses of all other atoms are a ratio to carbon 12 42

  43. Isotopes • Many elements have atoms that have multiple forms • Different forms of the same element having different numbers of neutrons are called isotopes. • For example: Carbon exists as both Carbon 12 and Carbon 14 • Carbon 12 Carbon 14 • 6 electrons 6 electrons • 6 protons 6 protons • 6 neutrons8 neutrons 43

  44. Isotopes and Atomic Mass Many elements have atoms that have multiple isotopes. Isotopes vary in abundance. Some are quite common while others are very rare. The atomic mass that appears in the periodic table is a weighted average taking into account the relative abundance of each isotope. 44

  45. or Na-24 or Na-23 Isotope:one of two or more atoms having the same number of protons but different numbers of neutrons

  46. Measuring Atomic Mass--the Mass Spectrometer The mass spectrometer can be used to determine the atomic mass of isotopes.

  47. Mass Spectrum of Neon • The mass spectrum neon shows three isotopes with the isotope at atomic mass = 20 accounting for more than 90% of neon.

  48. Mass Spectrum of Germanium • The mass spectrum of germanium shows 5 peaks at relative atomic masses of 70, 72,73,74, and 75

  49. Calculating the average relative atomic mass • The average atomic mass that is shown in the periodic table is really the weighted average of the atomic masses of each of the elements isotopes. Germanium has 5 isotopes whose relative atomic masses are shown in the table Mass Number% Abundance 70 20.55 72 27.37 73 7.67 74 36.74 75 7.67

  50. Calculating the Average Relative Atomic Mass • To calculate the average atomic mass multiply the atomic mass of each isotope by its abundance (expressed as a decimal fraction) Mass Number% Abundance 70 20.55 72 27.37 73 7.67 74 36.74 75 7.67 Average atomic mass = (0.2055)(70) + (0.2737)(72) + (0.0767)(73) + (0.3674)(74)+ (0.0767)(75) = 72.36 Note: atomic masses are ratios so they do not have real units although they are sometimes called atomic mass units or amu

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