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The Atom

The Atom. From Philosophical Idea to Scientific Theory. Dalton’s Model of the Atom. Solid sphere: Solid indivisible sphere. Dalton’s Atomic Theory. 1. All matter is composed of extremely small particles called atoms.

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The Atom

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  1. The Atom From Philosophical Idea to Scientific Theory

  2. Dalton’s Model of the Atom • Solid sphere: • Solid indivisible sphere

  3. Dalton’s Atomic Theory • 1. All matter is composed of extremely small particles called atoms. • 2. Atoms of a given element are identical in size, mass, and other properties… different elements have different properties • 3. Atoms cannot be subdivided, created, or destroyed

  4. Dalton’s Atomic Theory • 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds • 5. In chemical reactions, atoms are combined, separated, or rearranged.

  5. Modern Atomic Theory • (what we believe about the atom today) • Keeps some of Dalton’s postulates but others have been disproved. Now we know: 1. Atoms are divisible even into smaller particles 2. A given element can have atoms with different masses

  6. Cathode-ray experiments • Result of experiments between electricity and matter • Carried out in cathode ray tubes: Many of these experiments involved passing electric current through gases at low pressures

  7. CATHODE RAY TUBE • Picture – Label Parts

  8. CATHODE RAYS cont. • Experiments to Investigate Relationship between Energy and Matter • Procedure: Pass current from the cathode to the anode

  9. 1st Subatomic Particle Discovery • RESULTS & HYPOTHESIS: Current passed through tube surface of the tube directly opposite the cathode glowed Hypothesized glow caused by a stream of particles, which they called a cathode ray

  10. EXPERIMENTS TO TEST INITIAL HYPOTHESIS MAIN RESULTS • 1. Object placed between cathode and opposite end of tube cast a shadow on the glass • 2. A paddle wheel placed on rails between the electrodes rolled along the rails from cathode toward anode.

  11. RESULTS CONTINUED • 3. Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge • 4. The rays were deflected away from a negatively charged object.

  12. HYPOTHESIS FROM THE OTHER EXPERIMENTS • The particles that compose cathode rays are negatively charged

  13. JJ Thomson • Supported Hypothesis (that cathode rays are negatively charged particles) • Measured the ratio of the charge of cathode-ray particles to their mass: found it was always the same Therefore concluded that all cathode rays are composed of identical negatively charged particles (called electrons)

  14. CHARGE & MASS OF ELECTRON • Thomson’s Experiment Showed that the electron has a very large charge for its tiny mass

  15. Thomson’s Model of the Atom • Plum Pudding

  16. Millikan • Oil-drop Experiment

  17. ROBERT A. MILLIKAN’S EXPERIMENTS (1909) • He showed that the mass of the electron is about one two-thousandth (1/2000) the mass of the simplest type of hydrogen atom • More accurate experiments show the mass is actually 1/1837 the mass of H

  18. DISCOVERY OF THE ATOMIC NUCLEUS • Rutherford’s Experiment (Draw Picture)

  19. STARTLING RESULTS • Expected the alpha particleds to pass through with only a slight deflection (assumed mass and charge were uniformly distributed throughout the atoms of the gold foil) • Mostly true but about 1 in 8000 was redirected back toward source

  20. RUTHERFORD’S CONCLUSIONS • 1. Rebounded alpha particles must have experienced some powerful force within the atom • 2. The force must occupy a very small amount of space in the atom. The atom must be mostly empty space

  21. RUTHERFORD’S CONCLUSIONS • The force must be caused by a very densely packed bundle of matter with a positive electric charge. • He called this bundle of matter the nucleus.

  22. QUESTIONS LEFT TO PONDER • WHERE WERE THE ELECTRONS? Rutherford suggested that the electrons surrounded the nucleus like planets around the sun but did not know what kept the electrons in motion around the nucleus

  23. INSIDE THE ATOM – PART 2 SUBATOMIC PARTICLES ISOTOPES AVERAGE ATOMIC WEIGHT

  24. Subatomic particles • Protons (p+) • Nucleus • Determine the identity of the atom • Moseley’s organization of periodic table • Neutrons (n0) • Nucleus • Electrons (e-) • Most important in determining element’s properties

  25. Using Periodic Table • Periodic Table – Element information Atomic # # protons Element Symbol Average Atomic Weight (amu)

  26. Using Periodic Table • Periodic Table – Element information Atomic # # protons Element Symbol Average Atomic Weight (amu)

  27. Using Periodic Table • Subatomic particle values: • 1. Atomic # • 2. Atomic Weight • 3. Mass #

  28. Using Periodic Table

  29. Isotopes • 1. Definition • Isotope: atoms of the same element with different numbers of neutrons • 2. Notation • Symbol • Hyphen

  30. Average atomic weight • 1. Definition • Average atomic weight: weighted average of all naturally occurring isotopes of a given element • (Found on the periodic table)

  31. Avg. Atomic Mass Average atomic mass • 2. Formula:

  32. Average atomic weight • Example (Round to 2 decimal places) • Silver exists as 51.84% 107Ag and 48.16% 109Ag. The actual mass of 107Ag is 106.90509 amu and the actual mass of 109Ag is 108.90476 amu. What is the average atomic mass?

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