1. Electrons in Atoms Chapter 5…the truth about electrons

2. Atomic Models  1803  Dalton's model was that the atoms were tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior that was determined by what kind of element they were.

3. J.J. Thomson  “plum pudding” model  Discovered the “electron”  Electron embedded in a sphere of positive electrical charge  1897

4. Hantaro Nagaoka  1904  Suggested that an atoms has a central nucleus  No evidence to back claim

5. Rutherford’s experiment  Gold foil and alpha particles 1911

6. Alpha particle (+) and Gold foil  The above diagram shows what we would expect the result of Rutherford's experiment to be if the "plum pudding" model of the atom is correct.  The above indicates the actual result. Most of the alpha particles are only slightly deflected, as expected, but occasionally one is deflected back towards the source.

7. Rutherford conclusions  Atom mostly empty space (most of the alpha particles passed right through)  Tiny positive nucleus  Nucleus has large mass (alpha particle did not move nucleus, it bounced back so mass must be big)  Electrons move around nucleus

8. Niels Bohr “the old quantum theory”  Electrons found in specific circular paths, planetary model  Energy levels, energy increases to valence electrons  Quantum is the amount of energy it takes to move and energy level  1913

9. Quantum  These ladder steps are some what like energy levels. (a) In an ordinary ladder, the rungs are equally spaced. (b) The energy levels in atoms are unequally spaced, like the rungs in this ladder. The higher energy levels are closer together so the amount of energy is not always the same as you move.

10. Ground State vs. Excited State  Ground State= lowest energy configuration “normal”  Excited State=higher energy configurations than normal  Q: Why does hamburger have lower energy than steak?  A: Because it's in the ground state.

11. Why Bohr was not perfect -  The Bohr model gave results in agreement with experiment for the hydrogen atom. However, it still failed in many ways to explain the energies absorbed and emitted by atoms with more than one electron.

12. De Broglie’s particle-wave dulaity  In 1924 de Broglie's proposed that all moving particles has a wavelength is inversely proportional to momentum and that the frequency is directly proportional to the particle's kinetic energy ; λ = h/pde Broglie  This concept leads to one of the most important ideas in 20th century science:  The small, light, fast moving electron also exhibits wave-particle duality. It can be conceived of as a particle or as a wave.  This development leads to atomic and molecular orbitals.

13.  Edwin Schrödinger knew of de Broglie's proposal that a electrons exhibited wave- particle duality. With this idea in mind, he devised/constructed a differential equation for a wavelike electron resonating in three dimensions about a point positive charge.

14. Erwin Schrodinger  In the 1920s, a whole new theory of physics, called quantum mechanics, presented an even more radical picture of the atom. The electrons cannot be pinpointed but exist as a sort of cloud of probability outside the nucleus. quantum mechanicscloud of probability  The airplane propeller is somewhere in the blurry region it produces in this picture, but the picture does not tell you its exact position at any instant. Similarly, the electron cloud of an atom represents the locations where an electron is likely to be found.

15. Quantum Mechanical Model  Determines the allowed energies an electron can have and how likely it is (probability!) to find the electron in various locations around nucleus. Mathematical theory!  n= the Principle Quantum Number or Principle energy level (energy levels the electrons are found in)

16.  THE DISCOVERY OF THE ELECTRON 

17. Atomic Orbitals  Region in space with a high probability to find an electron

18. Energy Sublevel  Each energy sublevel corresponds to an orbital to a different shape.  Different atomic orbitals are denoted by letters.  S= spherical, P= dumbbell D & F more complex

19. Summary of Principal energy levels, sublevels, & orbital's Principal energy level# of sublevelsType of sublevels n=11 1s (1 orbital) n=22 2s(1orbital), 2p(3orbitals) n=33 3s(1 orbital), 3p(3orbitals), 3d(5orbitals) n=44 4s(1 orbital), 4p(3orbitals), 4d(5orbitals), 4f(7orbitals)

20. James Chadwick – 1932 Nucleus mystery solved  Rutherford speculated in 1920 that there existed electrically neutral particles with the protons that make up the missing mass but no one accepted his idea at the time.

21. Complete model!  James Chadwick discovered a new type of radiation that consisted of neutral particles. It was discovered that these neutral atoms came from the nuclei of the atom. This last discovery completed the atomic model.

22. Quarks….where are we now?  1964 Quarks are proposed by Murray Gell- Mann and George Zweig (math based theory)  quarks are subatomic particles thought to be elemental and indivisible. They are one of the two kinds of spin-½ fermions (the other being the leptons). Objects made up of quarks are known as hadrons; well known examples are protons and neutrons.subatomic particlesspinfermionsleptonshadrons protonsneutrons