1. The Life of Atom http://www.youtube.com/watch?v=h hbqIJZ8wCM

2. Birth In 1809 Dalton came up with the first picture of the atom. Tiny indestructible sphere

3. Childhood In 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. Adolescence In 1911 Rutherford published the results from the gold-foil experiment

5. The Results… Until now atoms were thought to be solid 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 (which is what was expected)

6. Gold-Foil Conclusions The atom is made up of mostly empty space Alpha particles are positive and curved when they got too close to the small nucleus Only alpha particles that hit the nucleus got deflected back (rarely happened so must be small)

7. Teen Years Entirely positive nucleus would explode (+ charges repel) The total mass number of the atom could not be accounted for In 1932 Atom gets a girlfriend…the neutron is discovered!

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

9. Trouble in paradise? Rutherford’s model doesn’t quite work Electrons should lose energy and crash into the nucleus but this doesn’t happen 19 th 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

10. Bohr’s Addition to Atom In 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 defined energy

11. How does a line spectrum say all that? When an electron is excited it jumps from one orbital to a higher one 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 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 In 1926 Schrodinger derived the quantum mechanical model of the atom Described atoms as having wave-like properties which came from de Broglie’s hypothesis Mathematically determined the shape of the orbitals and the probably of an electron being in a certain place at a certain time- orbitals are not just spheres anymore

15. In 1927 Heisenburg stated: Although the shape of the orbital is predictable, the exact location of an electron cannot be determined. This is the Heisenburg uncertainty principle

16. Quantum Theory Quantum is the new and improved Bohr- Rutherford model Model shows electron placement and helps to determine valence electrons and stability of the atom Each orbital can hold a maximum of 2 electrons

17. Orbital Shapes and Orientation S is a sphere shape – 1 orientation = 1 orbital = 2 electrons P is a figure 8 or dumbbell shape – 3 orientations = 3 orbitals = 6 electrons

18. D orbitals have a clover shape – 5 orientations = 5 orbitals = 10 electrons

19. F orbitals have many shapes – 7 orientations = 7 orbitals = 14 electrons

20. Rules for Quantum 1.Aufbau Principle - each electron is added into the subshell with the lowest energy orbital available 2.Hund’s Rule - Each orbital subshell gets a single electron first and then electron can pair. All electron are ‘up’ when single 3.Pauli Exclusion Principle - no electron can have the same 4 quantum #s in an atom - electron sharing an orbital have opposite spins

21. Quantum Number Principle quantum number ‘n’ Describes the orbital’s energy level and relative size Orbital shape quantum number ‘l’ Describes the orbital’s shape, energy of subshells Magnetic quantum number ‘m l ’ Describes the orientation in space Spin quantum number ‘m s ’ Describes the behaviour of a specific electron in an orbital

22. Summary of Quantum numbers of electrons in atoms NameSymbolAllowed ValuesProperty Principal (Shell)nPositive integers (1, 2, 3, etc.) Orbital size and energy level Orbital shape (subshell) lIntegers from 0 to (n-1) Orbital shape (l values 0, 1,2 and 3 correspond to s, p, d and f orbitals) Magneticmlml Integers from –l to +l Orbital orientation Spinmsms +1/2 or -1/2Spin orientation