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Electrons In Atoms

Electrons In Atoms. Where are they?. Development of Atomic Models. Plum Pudding Model (1897) J.J. Thomson Electrons scattered in a “sea” of positive charges. Development of Atomic Models. Rutherford’s Model (1911) Discovered nucleus (disproves Plum Pudding)

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Electrons In Atoms

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  1. Electrons In Atoms Where are they?

  2. Development of Atomic Models • Plum Pudding Model (1897) • J.J. Thomson • Electrons scattered in a “sea” of positive charges

  3. Development of Atomic Models • Rutherford’s Model (1911) • Discovered nucleus (disproves Plum Pudding) • Electrons orbit nucleus like planets around the sun • Cannot explain many of the properties of atoms

  4. Development of Atomic Models • Bohr Model (1913) • Electrons move around nucleus in circular orbits at specific allowed distances • These distances relate to allowable energy levels • Energy levels – fixed energies an e- can have • Quantum of energy – energy needed to move an electron from one E level to another

  5. Development of Atomic Models • More Bohr Model • Electrons can gain or lose energy • Ground state – lowest energy level available • Excited state – higher energy level • Absorb energy (gain E) • Go from lower to higher E levels • Emit energy (lose E) • drop from higher to lower E levels • Give off E in the form of radiation (quanta of light)

  6. Development of Atomic Models • More Bohr Model • Energy levels get closer together as they get farther from the nucleus • Problem: Works well with the hydrogen atom but not much else

  7. Development of Atomic Models • Quantum Mechanical Model (1926) • Modern description of electrons in atoms • Cloud model or Quantum Theory • Schrodinger – developed mathematical equation to predict atomic behavior • Electrons NOT in exact path • Heisenberg Uncertainty Principle • Impossible to know both location and energy of an electron • Can measure one or the other – NOT both • Exact motion of electron unknown

  8. Development of Atomic Models • More Quantum Mechanical Model • Determines allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus • These locations are called principal energy levels • Within these energy levels are sublevels • Sublevels are subdivided into atomic orbitals

  9. Development of Atomic Models • More Quantum Theory • Atomic Orbitals • Region of space in which there is a high probability of finding an electron • 4 types of orbitals – s, p, d, and f • Different orbitals have different shapes • Each orbital can hold up to 2 electrons • 2 electrons in same orbital must have opposite spins (+1/2 and -1/2)

  10. s and p orbitals

  11. d sublevels

  12. f orbitals

  13. Quantum Numbers • 4 numbers used to describe electron location • Principal Energy Level (Principal Quantum Level) • n = 1, 2, 3… • Energy Sublevel Number • specifies s, p, d, or f sublevel • l = 0 to n-1 • l = 0 s sublevel • l = 1 p sublevel • l = 2 d sublevel • l = 3 f sublevel

  14. Quantum Numbers • Orbital quantum number (m) • m = -l to +l • Specifies which orbital within a sublevel the electron is located • Within sublevels, orbitals differ only in spatial orientation, not energy • Spin quantum number (ms) • ms = +1/2 or -1/2 • 1st electron in orbital has + spin

  15. Energy Levels, Sublevels, and Orbitals *** Remember: Each orbital can hold 2 electrons ***

  16. Orbitals and Electrons s sublevel – 1 orbital, 2 electronsp sublevel – 3 orbitals, 6 electrons d sublevel – 5 orbitals, 10 electronsf sublevel – 7 orbitals, 14 electrons Maximum # electrons / energy level = 2n2 where n = energy level

  17. Electron Configuration • The way in which electrons are arranged in various orbitals around the nucleus of an atom • Aufbau Principle • Electrons occupy the orbitals of lowest energy first

  18. Electron Configuration • Pauli Exclusion Principle • An atomic orbital may describe at most two electrons • Opposite spins • Boxes represent orbitals and arrows represent electrons • 3s sublevel with 1 electron  • 4s sublevel with 2 electrons 

  19. Electron Configuration • Hund’s Rule • Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin as large as possible • Orbitals of equal energy each get 1 electron before any pair up 2p sublevel

  20. Electron Configuration • Diagonal rule • s holds 2 • p holds 6 • d holds 10 • f holds 14

  21. Atomic Structure Practice • 3a) 1s22s22p3 Nitrogen ( 7 electrons) • b) 1s22s2 Beryllium (4 electrons) • c) 1s22s22p63s23p3 Phosphorus (15 electrons) • d)1s22s22p63s23p63d54s2 Manganese (25 electrons) e) Potassium (19) f) Zirconium (40) g) Promethium (61) h) Selenium (34)

  22. Atomic Structure Practice • 4a)Cu0 (29 e-) 1s22s22p63s23p64s23d9 • Cu+ (28 e-) 1s22s22p63s23p64s23d8 • Cu2+ (27 e-) 1s22s22p63s23p64s23d7 • 4b) Al0 (13 e-) 1s22s22p63s23p1 • Al3+ (10 e-) 1s22s22p6

  23. Electron Configuration • Short Cut Method • Rare Gas Configuration, Noble Gas Configuration, or Inert Gas Configuration (Either name OK) • Relate back to the previous rare gas • Put that element in [ ] • Start at s sublevel using whatever period the element is in • Nickel 1s22s22p63s23p64s23d8 or[Ar]4s23d8

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