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The Life and Times of Atom

The Life and Times of Atom. A story of one atom’s coming of age. BIRTH. 1809 - Dalton: pictured the atom as a tiny indestructible sphere Atom’s baby picture. Early Childhood (The awkward years). 1897-Thomson: discovered very light weight negatively charged particles (electrons)

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The Life and Times of Atom

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  1. The Life and Times of Atom A story of one atom’s coming of age

  2. BIRTH • 1809 - Dalton: pictured the atom as a tiny indestructible sphere Atom’s baby picture

  3. Early Childhood (The awkward years) • 1897-Thomson: discovered very light weight negatively charged particles (electrons) • Chemists determined that the negative charge must be balanced by a positive charge: the raisin bun model

  4. Early Adolescence • 1911 -Rutherford (McGill University) - publishes the results from the famous gold-foil experiment

  5. The Gold-Foil Experiment

  6. Shocking Results!!! • Until this point, atoms were thought to be solid throughout • Most of the alpha particles went right through the foil! • Some alpha particles curved when they went through • Only a few alpha particles deflected back (This was the expected result - think of running into a solid wall)

  7. Gold Foil Conclusions • The atom is made up of mostly empty space • Alpha particles are positive, they curved if they got too close to the small nucleus • Only alpha particles that hit the nucleus were deflected back, since this rarely happened, the nucleus must be very small!

  8. Atom’s Troubled Teen-aged Years • An entirely positive nucleus would explode (+ charges repel) • The total mass of the atom couldn’t be accounted for • 1932 - Atom gets a girlfriend! The neutron is discovered

  9. Rutherford’s Model of the Atom • The nucleus is small and made up of protons and neutrons • The electrons circle around the nucleus

  10. Problems in Paradise??? Rutherford’s model doesn’t quite work: • Electrons should lose energy and crash into the nucleus (this clearly doesn’t happen) • 19th century physics dictates that a body in motion must continuously give off energy - seen as a continuous spectrum through a spectroscope - but we see a line spectrum

  11. Bohr’s Addition to the Atom • 1913 - Bohr explains why a line spectrum is seen instead of a continuous spectrum • Electrons are only giving off certain frequencies of light • Electrons travel in defined spaces called orbitals, which have a defined energy

  12. How does a line spectrum tell us all that? • When an electron is excited (given energy) it jumps from one orbital to a higher orbital • The electron does not stay excited and eventually goes back to its ground state (original orbital) • A wave of light is emitted (photon) from this process which can be seen as a line on a line spectrum

  13. Problems with Bohr’s Theory • Bohr couldn’t explain why lines appeared in ones, threes, fives and sevens - more on this later! • Physicist Max Planck supported Bohr’s idea that atoms can absorb or emit only discrete quantities of energy called quantums • Einstein called these ‘packets’ of energy photons

  14. Adulthood • 1926 - Schrodinger - derived the quantum mechanical model of the atom • Described electrons as having wave-like properties • Mathematically determined the shape of orbitals and the probability of an electron being in a certain place at a certain time - orbitals are not just spheres anymore! • 1927 - Heisenburg - Heisenburg Uncertainty Principle: Although the shape of the orbital is predictable, the exact location of an e- can not be determined

  15. Atom’s Portrait 1927

  16. The Four Quantum Numbers (which are actually letters)

  17. Why Use Quantum Theory? • Quantum is the ‘new and improved’ Bohr-Rutherford diagram • This model shows e- placement which helps us determine valence e- and stability of an atom, this allows us to predict atom behaviour • Each orbital can hold a maximum of 2e-

  18. Orbital Shapes & Orientation s is a sphere shape - 1 orientation = 1 orbital = 2e- p is a figure eight - 3 orientations = 3 orbitals = 6e-

  19. d orbitals have a ‘flower’ shape - 5 orientations = 5 orbitals = 10 e- f orbitals have many shapes - 7orientations in = 7 orbitals = 14 e-

  20. Rules for Quantum • Aufbau Principle - each e- is added into the subshell with the lowest E orbital available • Hund’s Rule - Each orbital subshell gets a single electron first and then e- can pair. All e- are ‘up’ when single • Pauli Exclusion Principle - no e- can have the same 4 quantum #s in an atom - e- sharing an orbital have opposite spins

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