1 / 40

Atoms, Molecules and Ions History

Explore the fascinating journey of atomic theory, from Democritus and Aristotle to Dalton, Thomson, Rutherford, and beyond. Learn about the discovery of subatomic particles and the development of the modern atomic model.

mestas
Download Presentation

Atoms, Molecules and Ions History

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Atoms, Molecules and IonsHistory • Greeks • Democritus 460-370 BC “atomos” • Aristotle- elements. • Alchemy • 1660 - Robert Boyle- experimental definition of element. • Lavoisier (1734-1794)- Father of modern chemistry. He wrote the book -1789.

  2. Laws • Conservation of Mass • Law of Definite Proportion- compounds have a constant composition. • Multiple Proportions- When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

  3. Multiple What??? • Water has 8 g of oxygen per g of hydrogen. • Hydrogen peroxide has 16 g of oxygen per g of hydrogen. • 16/8 = 2/1 • Small whole number ratios.

  4. Dalton’s Atomic Theory1803-1807 1) Elements are made up of atoms 2) Atoms of each element are identical. Atoms of different elements are different. 3) Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. 4) Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

  5. A Helpful Observation • Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. • Avogadro- interpreted that to mean nat the same temperature and pressure, equal volumes of gas contain the same number of particles. • (called Avogadro’s Hypothesis)

  6. Thomson’s Experiment (1897) • Used the Cathode Ray Tube(CRT) to discover the electron.

  7. Thomson’s Experiment Passing an electric current makes a beam appear to move from the negative (cathode) to the positive end (anode).

  8. Proving Electrons had Mass • Thomson devised a cathode ray tube with a paddle wheel built inside. When the high voltage electricity was turned on the paddle wheel began to rotate and move away from the cathode and towards the anode.

  9. Determining Charge • Thomson concluded from this evidence and from his previous experiments that tiny particles were being emitted from the atoms of the cathode. These tiny particles were negatively charged. He called these particles "electrons."

  10. Thomson’s Model of the Atom • Given these experimental results, Thomson proposed in 1897 what has since been referred to as the "plum pudding" model of the atom. This model depicts the atom as a diffuse cloud of positive charge with the negative electrons embedded randomly in it, like plums in pudding.

  11. Millikan (1909) - Mass of the Electron “(Oil Drop Experiment)” • Using this information he calculated the mass of the electron as 9.11 X 10-28 grams.

  12. Radiation • In 1896, the French scientist, Henri Becquerel, found that a piece of a mineral containing uranium could produce its image on a photographic plate in the absence of light. He attributed this phenomenon to a spontaneous emission of something which he called "radiation" which originated from the uranium.

  13. Radioactivity • Discovered by accident • Bequerel • Marie and Pierre Curie isolated the radioactive components at Bequerel’s suggestion • Three types • –alpha- helium nucleus (+2 charge, large mass) • __beta- high speed electron • –gamma- high energy light

  14. Rutherford (1911) • A fluorescent screen would detect radiation by flashing whenever it was struck by the radiation. • Rutherford observed that the radiation was diffracted into three beams by the charged plates.

  15. Rutherford “Gold Leaf Experiment” • Rutherford set up the apparatus which would bombard thin gold foil with alpha particles. These particles would then be detected by the fluorescent screen.He anticipated that all of the alpha particle detection would occur on the screen directly behind the foil.

  16. Rutherford got surprisingly different results • "It was about as credible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you!"

  17. Atomic Model Revised • Rutherford suggested that the atom was mostly empty space with a highly charged center. Most of the particles pass through the atom undisturbed, but a few get too close to the center and are deflected.

  18. "Planetary Model." • To account for these results Rutherford proposed a new model of the atom in 1913. This model had the following characteristics:The atom is mostly empty space with the majority of its mass concentrated in the center of the atom which he called the "nucleus."

  19. Mass Number and Atomic Mass • By the early 1930's the major subatomic particles had been discovered and their physical properties had been described. • James Chadwick (1932) - discovers Neutrons. Particles Mass (grams) Relative Mass (amu) Relative Charge Proton 1.67262 X 10-27 1.007 +1 Electron 9.10939 X 10-31 5.486 X 10-4 (~ 0) -1 Neutron 1.67493 X 10-27 1.009 0

  20. Atomic Mass Unit • amu - the unit that we use to measure atoms. • 1 amu ~ mass of one mole of the following: ~hydrogen ~ 1 proton ~ 1 neutron • Actual masses: • Particle Charge Mass (amu) • Proton positive(+1) 1.0073 • Neutron none(neutral) 1.0087 • Electron negative (-1) 5.486 x 10-4 • Hydrogen none (neutral) 1.0079

  21. Measuring Atoms o Angstrom (A) - a convenient non-SI unit of length used to express atomic dimensions. • 1 Angstrom = 1 x 10-10 meters • most atoms have diameters between 1 x 10-10 m and 5 x 10 -10 m, or between 1 - 5A. o

  22. Isotopes • All atoms have the same number of protons. • The number of neutrons may vary for a given element. • Isotope - atoms of a given element that differ in the number of neutrons and consequently the mass. • May be written as the symbol 12C or simply carbon-12 as opposed to the isotope 14C or carbon-14. 6 6

  23. Atomic Numbers and Mass Numbers • Atomic Number - is the number of protons, which is shown as the subscript. (it is also the number of electrons in a neutral atom.) • Mass Number - is the total number of protons plus neutrons in the atom. (which represent essentially all the mass of the atom.) • ex. 12C • has 6 Protons, 6 Electrons and 6 Neutrons • Number of Neutrons = Mass # - Atomic # 6 6 = 12 - 6

  24. Symbols • Contain the symbol of the element, the mass number and the atomic number. Mass number X Atomic number

  25. Nuclides • Nuclide - an atom of a specific isotope • ex. 14C or carbon-14 • 6 protons, 6 electrons and 8 neutrons • # of neutrons = atomic mass - atomic number • 8 = 14 - 6 6

  26. Symbols • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass Number • Name 24 Na 11

  27. Symbols • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass Number • Name 80 Br 35

  28. Symbols • if an element has 91 protons and 140 neutrons what is the • Atomic number • Mass number • number of electrons • Complete symbol • Name

  29. Symbols • if an element has 78 electrons and 117 neutrons what is the • Atomic number • Mass number • number of protons • Complete symbol • Name

  30. Atomic Mass • How heavy is an atom of oxygen? • There are different kinds of oxygen atoms. • More concerned with average atomic mass. • Based on abundance of each element in nature. • Don’t use grams because the numbers would be too small.

  31. Measuring Atomic Mass • Unit is the Atomic Mass Unit (amu) • One twelfth the mass of a carbon-12 atom. • 6 p+ and 6 n0 • Each isotope has its own atomic mass • we get the average using percent abundance.

  32. Calculating averages Average = (% as decimal x mass) + (% as decimal x mass) + (% as decimal x mass )

  33. Atomic Mass • Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of 62.93 amu and the rest has a mass of 64.93 amu.

  34. Atomic Mass • Magnesium has three isotopes. 78.99% magnesium 24 with a mass of 23.9850 amu, 10.00% magnesium 25 with a mass of 24.9858 amu, and the rest magnesium 25 with a mass of 25.9826 amu. What is the atomic mass of magnesium? • If not told otherwise, the mass of the isotope is the mass number in amu

  35. Periodic Table • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties.

  36. Groups These five groups are known by their names.

  37. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H).

  38. Periodic Table Metalloids border the stair-step line (with the exception of Al and Po).

  39. Periodic Table Metals are on the left side of the chart.

  40. Diatomic Molecules These seven elements occur naturally as molecules containing two atoms.

More Related