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Electron Configurations

Electron Configurations. Locating Electrons. Newtonian Mechanics – describes objects at ordinary velocities (classical mechanics) Quantum Mechanics – describes particles at velocities near that of light (subatomic particles) Quanta – a packet of nrg. Heisenberg’s Uncertainty Principle.

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Electron Configurations

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  1. Electron Configurations Locating Electrons

  2. Newtonian Mechanics – describes objects at ordinary velocities (classical mechanics) • Quantum Mechanics – describes particles at velocities near that of light (subatomic particles) • Quanta – a packet of nrg

  3. Heisenberg’s Uncertainty Principle • It is impossible to know both the location and velocity of an electron at the same time. • To see an e- we would have to bounce light off of it which would change its velocity and in turn its position. • Ex: helium-filled balloon in a dark room

  4. Schrödinger (1887-1961) • In 1926, he treated e- as waves, giving us the e- cloud model. • Radial Probability of Electrons The area of highest probability forms the e- cloud.

  5. Locating Electrons • Principle Quantum Number (n) • Sublevel (l) • Orbital (m) • Spin (s)

  6. Principle Quantum Number (n) • Energy levels are a particular distance from the nucleus # e- = 2 8 18 32 50 n = 1 2 3 4 5

  7. Principal Quantum Numbers (n) • The maximum number of electrons in each nrg level is 2n2 • At n = 1, there can be 2(1)2 = 2 e- • At n = 2, there can be 2(2)2 = 8 e- • At n = 3, there can be 2(3)2 = 18 e-

  8. Sublevel (l) • Tells the shape • Each nrg level has a # of sublevels = to n

  9. Orbital (m) • The 3rd quantum number (m) tells which orbital and electron occupies. • One pair (2e-) of electrons can occupy each orbital • s sublevels have 1 orbital (2e-) • p sublevels have 3 orbitals (6e-) • d sublevels have 5 orbitals (10e-) • f sublevels have 7 orbitals (14e-) • ** each orbital can hold UP TO 2 e-**

  10. “s” and “p” orbitals

  11. “d” orbitals

  12. Spin (s) • Indicates direction of spin of e- • -1/2 , +1/2 (clockwise, counterclockwise) • Pauli Exclusion Principle states that no two electrons in an atom can have the same set of 4 quantum numbers. • The two e- in an orbital must have opposite spins.

  13. Electron Configuration Notation • Helium has 2 electrons • n = 1 • l = s • m = 1 • s = 1 up, 1 down • Helium’s electron configuration would be: # of e- Principle Quantum # Sublevel

  14. Examples • Li • N • Ne • Na

  15. Degenerateorbitals have the same nrg

  16. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f • Sublevels fill in order of increasing nrg • 1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p

  17. Predicting electron configuration from the Periodic Table

  18. Example • What is the electron configuration for: • Cl – 17 • Pb - 82

  19. AUFBAU Principle • e- first occupy the lowest possible nrg level available. • Electron Dot Notation – show only valence e-, those in the outer most nrg level • ONLY UP TO 8e- • 8 e- = stable • Valence electrons – e- in the highest nrg levels • These e- are what form bonds

  20. Electron Dot Notation • Examples:

  21. Rules • Only show valence electrons • Dots are either placed 1 on each side or in pairs. • Never more than 2 per side • This is why lithium has only one dot and why carbon can have 2 dot notations.

  22. Example • What would the electron dot notation be for titanium? • Ti = 22e- BUT only 2 valence e- • Electron Configuration Notation • Electron Dot Notation Valence e- • d’s are NEVER valence e-, they ALWAYS fill after a high nrg level • Same for f’s

  23. Orbital Notation • Show all orbitals with electrons • Electrons represented as up and down arrows • Arrows must be opposite within orbitals Nitrogen (7) 1s 2s 2p Fluorine (9) 1s 2s 2p

  24. Hund’s Rule • Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals.

  25. Electrons • Quantum Mechanical Model – describes and electron as having a particular amount of energy, depending on its location. • Electron clouds give you the highest probability of locating and electron.

  26. Ions • Excited state – electrons in a higher than normal energy state. • Nitrogen: 1s2 2s2 2p3 (ground state) 1s2 2s2 2p2 3s1 (excited state)

  27. Ions • Lose or gain e- • Anions – are negatively charged, having gained e- • Cations – are positively charged, having lost e- ** atoms will gain or lose e- to become more stable**

  28. Examples • Na: 1s2 2s2 2p6 3s1 • Na+: 1s2 2s2 2p6 • Alkali metals, like Na, want to lose their 1 valence e- to become stable. • Cl: 1s2 2s2 2p6 3s2 3p5 • Cl-: 1s2 2s2 2p6 3s2 3p6 • Halogens, like Cl, want to gain a valence electron to become stable.

  29. Exceptions • Filled and half-filled sublevels are more stable than partially filled sublevels. • This Cr takes an e- from 4s to put one e- in each of its 3d orbitals and Cu takes a 4s to fill each of its 3d orbitals • Orbitals are stable when either full or half-full 1s 2s 2p 3s 3p 4s 3d

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