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Atomic Theory
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  1. Atomic Theory Everything written in black has to go into your notebook Everything written in blue should already be in there

  2. Atom: Smallest part of any element, which can take part in a chemical reaction • Element: Pure substance consisting of one type of atom only • Compound: Substance consisting of two or more elements (e.g. H2O) • Mixture: Consists of two or more substances but they are not joined together chemically (e.g. sand and iron filings)

  3. Molecule: Smallest part of a substance which can exist on its own (e.g. Cu, H2, H2O) • There are monatomic, diatomic or triatomic molecules

  4. 1 a.m.u. = 1.66 x 10-27 kg

  5. The Bohr Model Protons and neutrons are located in the nucleus of the atom Electrons orbit the nucleus in ‘shells’ or ‘energy levels’

  6. 1st energy level can hold 2 electrons 2nd energy level can hold 8 electrons 3rd energy level can hold 18 electrons

  7. 7 3 3P 4N E.g. Lithium Li

  8. 23 11 11P 12N E.g. sodium Na

  9. Isotopes: Atoms with the same number of protons but different numbers of neutrons (e.g. carbon 12, carbon 13 and carbon 14)

  10. Isotopes: Atoms having the same number of protons but different numbers of neutrons (e.g. carbon 12, carbon 13 and carbon 14) • Atomic Number (Z): The number of protons in an atom • Mass Number (A): The number of protons and neutrons in an atom • Relative atomic mass: The mass of an atom compared to of the mass of an atom of carbon

  11. An energy level is a specific level of energy which an electron has in an atom

  12. The emission spectrum is the series of lines which are given out by excited atoms of an element

  13. E2 Higher energy level Photon of light emitted Lower energy level E1

  14. E2 – E1 = hf E2 = energy of electron in higher energy level E1 = energy of electron in lower energy level h = Planck’s constant f = frequency of light emitted

  15. Flame tests (page 126)

  16. The ground state is the energy level of the electron before it gains energy The excited state is the energy level of the electron after it gains energy

  17. The Aufbau Principal • Electrons always fill the lowest energy level available when the atom is in the ground state • Remember that lower energy levels are nearer the nucleus, and higher energy levels are further away from the nucleus

  18. 2 8 18 • 1st shell (n =1) holds up to electrons • 2nd shell (n=2) holds up electrons • 3rd shell (n=3) holds up to electrons • 4th shell (n=4) holds up to electrons • Each of these main energy levels (shells) contains “sub-levels” 32

  19. Consider it like 5th year is split into 3 classes; so an energy level is split into ‘sub-energy levels’

  20. 4f 4d N=4 4p s holds 2 p holds 6 d holds 10 f holds 14 4s 3d N=3 3p 3s 2p N=2 2s N=1 1s

  21. They are filled in this order 4f 4d 4p 3d 3d “jumps” up! 4s 3p 3s 2p 2s 1s

  22. Writing electronic configurations: Example 2 Cobalt (Co) 1s2 2s2 2p6 3s2 3p6 4s2 3d7

  23. Copper and chromium are the only exceptions to the trend

  24. 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 4s1, 3d5 3d4

  25. 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 4s1, 3d10 3d9

  26. Atomic Orbitals • An atomic orbital is the region around a nucleus where there is a high probability of finding an electron

  27. Main energy level Atomic orbitals Sub-level s s px py pz

  28. s s px py pz

  29. s s px py pz s px py pz

  30. Every orbital can hold a maximum of 2 electrons

  31. s sub-level has 1 orbital • spherical in shape

  32. p sub-level has 3 orbitals • dumb bell shape px py pz

  33. d sub-level has 5 orbitals

  34. Hund’s Rule of Maximum Multiplicity • When two or more orbitals of equal energy are available to electrons the electrons occupy them singly, before filling them in pairs Sometimes called the“Bus Seat Rule”

  35. Pauli’s Exclusion Principle • No more than two electrons can occupy an orbital, and they can only do this if they have opposite spin

  36. Example • Oxygen • O

  37. 1s2 2s2 2p4

  38. Quantum Numbers • A code consisting of 4 numbers, which give the full information about any one electron in an atom

  39. Question • Give the 4 quantum numbers for each of the electrons in the Berylium atom

  40. electronic configuration: 1s2 2s2 1st electron: 1, 0, 0, ½ 2nd electron: 1, 0, 0, -½ 3rd electron: 2, 0, 0, ½ 4th electron: 2, 0, 0, -½

  41. The Periodic Table • Groups go from top to bottom • Elements in the same group have the same number of electrons in their outer shell, and so have similar chemical properties • NB: Group 1 (Alkali metals) • Group 2 (Alkaline earth metals) • Group 7 (Halogens) • Group 8 (Noble gases)

  42. Periods go from left to right • Elements in the same period have the same number of electron shells occupied

  43. A group is a vertical column in the Periodic Table A period is a horizontal row in the Periodic Table