1 / 48

Chapter 6

Chapter 6. Arrangement of Electrons in Atoms. 0r…. Matter waves and waves that don’t matter. The nature of light. Dual nature of light Wave characteristics   Particle characteristics. Wave nature of light. Electromagnetic radiation. Electromagnetic (EM) radiation.

quemby-kent
Download Presentation

Chapter 6

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 6 • Arrangement of Electrons in Atoms

  2. 0r…. • Matter waves and waves that don’t matter

  3. The nature of light • Dual nature of light • Wave characteristics  Particle characteristics

  4. Wave nature of light

  5. Electromagnetic radiation

  6. Electromagnetic (EM) radiation • Form of E w/ wavelength (l) behavior • Speed = 3.0 x 1010 cm/s (speed of light) • Wavelength (l)  distance between pts. on a wave

  7. Frequency (n) • # of waves that pass a given pt. in a specific time

  8. frequency • C = ln • Therefore, as l decreases, n increases • C = speed of light, (186,000 miles/s, or 299,792,458 m/ s)

  9. Continuous spectrum • All ls in a given range included

  10. Electromagnetic (EM) spectrum • All EM radiation

  11. Particle nature of light

  12. Photoelectric effect • Emission of e- by certain metals when light shines on them

  13. Max Planck (1900) • When a hot object loses E, it is lost in sm. Specific amts. Called quanta

  14. Quantum • Finite quantity of E that can be gained or lost by an atom

  15. Photon • Individual quantum of light

  16. Albert Einstein (1905) • Higher n = higher E • Absorb. Of photons of a specific E explains photoelectric effect

  17. Dual (wave-particle) nature of light

  18. Important formulas E = hn • h (Plank’s constant) = 6.626 x 10 -34 J . S • n (frequency) c = ln c (speed of light) = 186,000 miles/s, or 299,792,458 m/ s

  19. Hydrogen atom spectrum • Pass high voltage through H2 gas  gas glows pass light through prism bright line spectrum

  20. Bright line spectrum • Due to e-s boosted to high E state (excited state), then dropping to the ground state • Lines represent E given off when e-s drop to ground state

  21. Hydrogen spectrum • E of photon = difference between ground and excited state

  22. spectroscope

  23. Flame test

  24. Bohr Model of the atom (1913) • The Hydrogen e- can circle the H nucleus only in certain orbits (like rungs of a ladder) • Definite orbits occupied by electron particles • Worked w/ H atom only

  25. According to this theory an electron moving between orbits would disappear from one and reappear instantaneously in another without visiting the space between  “Quantum leap”

  26. “An electron doesn’t exist until it is observed” “Until it is observed an electron must be regarded as being at once everywhere and nowhere” Dennis Overbye

  27. Schrödinger Model (1926) • Wave properties of atoms • Worked w/ all atoms • e- in orbitals • e- clouds • Can not pinpoint location of e- and path at a given instant  immutable property of the universe

  28. Quantum numbers • “Electron address” • Location of e-s in the atom

  29. Quantum number 1“Pennsylvania” • Principle quantum number (main energy level) • n= 1,2,3……7

  30. Quantum number 2“Hollidaysburg” • Orbital quantum number (shape of orbital) • s,p,d,f

  31. Quantum number 3“N. Montgomery St.” • Magnetic quantum number (orientation of orbital about the nucleus)

  32. Quantum number 4“1510” • Spin Quantum number (two possible states of electron) • +1/2 or -1/2

  33. Arrangement of electrons

  34. Arrangement of electrons • Main energy levels: 1,2,3….. • Sublevels: s,p,d,f • Orbitals • Each s has 1 • Each p has 3 • Each d has 5 • Each f has 7 • Each orbital can hold a max of 2 e-

  35. Orbital notation • Unoccupied orbital ___ • Orbital with 1 e- #or $ • Orbital with 2 e- #$ • e.g. H #He #$ 1s 1s

  36. Orbital notation

  37. Electron configuration notation • Uses superscripts instead of lines • e.g. H 1s1 or He 1s2

  38. Electron dot notation • Uses only e- in highest (outermost) main energy levels • e.g. .Na .He.

  39. Aufbau (building up) principle • Electrons occupy lowest energy orbital that will receive them, e.g. hydrogen’s electron goes into the 1s orbital

  40. Hund’s rule • Orbitals of equal E are each occupied by one e- before 2nd e- is added, all e- in singly occupied orbitals must have same spin

  41. Pauli exclusion principle • No two e- in same atom have the same set of four quantum numbers

  42. Electron fill chart

  43. Shorthand notation

  44. exceptions • e.g. copper • [Ar] 4s1 3d10

More Related