![Electron Configurations (ions)
• F-, 10, Flouride: [1s2 2s2 2p6 ]• Cl-, 18, Chloride : [1s2 2s2 2p6 3s2 3p6]• Na+, 10, Sodium ion: [1s2 2s2 2p6 ]+](https://image.slidesharecdn.com/1-140308201016-phpapp01/85/1-4-atomic-structure-part1-72-320.jpg)
1.4 atomic structure part1
- 1. Learning Outcomes • Emission and absorption spectra of the hydrogen atom . • Balmer series in the emission spectrum as an example. • Line spectra as evidence for energy levels. • Energy sub-levels. • Viewing of emission spectra of elements using a spectroscope or a spectrometer.
- 3. Spectrscope In a light spectroscope, light is focused into a thin beam of parallel rays by a lens, and then passed through a prism or diffraction grating that separates the light into a frequency spectrum.
- 5. Continuous spectrum A Spectrum in which all wavelengths are present between certain limits.
- 8. Spectrum lines When light from an unknown source is analyzed in a spectroscope, the different patterns of bright lines in the spectrum reveal which elements emitted the light. Such a pattern is called an emission spectrum.
- 10. Emission Spectrum • Shows that atoms can emit only specific energies (discrete wavelengths, discrete frequencies)
- 11. hypothesis: if atoms emit only discrete wavelengths, maybe atoms can have only discrete energies
- 14. A turtle sitting on a staircase can take on only certain discrete energies
- 15. energy is required to move the turtle (electron) up the steps (energy levels) (absorption) energy is released when the turtle (electron) moves down the steps (energy levels) (emission)
- 16. energy staircase diagram for atomic hydrogen
- 17. bottom step is called the ground state higher steps are called excited states
- 18. Balmer Series • Balmer analysed the hydrogen spectrum and found that hydrogen emitted four bands of light within the visible spectrum: • Wavelength (nm) Color • 656.2 red • 486.1 blue • 434.0 blue-violet • 410.1 violet
- 19. Flame Test • Flame Test The following metals emit certain colours of light when their atoms are excited. • Metal Colour • Sodium (Na) Yellow • Lithium (Li) Pink/Red • Potassium (K) Purple • Copper (Cu) Green • Calcium (Ca) Pink • Barium (Ba) Yellow/Orange • Strontium (Sr) Red/Orange
- 20. Learning Outcomes • Energy levels in atoms. • Organisation of particles in atoms of elements nos. 1–20 (numbers of electrons in each main energy level). • Classification of the first twenty elements in the periodic table on the basis of the number of outer electrons.
- 21. Bohr
- 22. Bohr’s theory • Electrons revolve around nucleus in orbits • Electron in orbit has a fixed amount of energy • Orbits called energy levels • If electron stays in level it neither gains nor loses energy
- 23. Bohr • Atom absorbs energy • Electron jumps to higher level • Atom unstable at higher levels. Electron falls back to a lower level • Atom loses or emits energy of a particular frequency.
- 24. quantisation • Electrons can have only certain particular values of energy
- 25. EVIDENCE FOR ENERGY LEVELS • In Hydrogen electron in lowest (n=1) level; ground state • Energy given; electron jumps to higher level excited state • Falls back and emits a definite amount of energy • Energy appears as a line of a particular colour
- 26. colours • Energy emitted depends on the jumps • Different jumps emit different amounts of energy and hence different colours
- 27. Main energy levels (shells) • Spectroscopic notation for shells . • N shell name 1=K 2=L 3=M 4=N
- 29. Bohr Diagram
- 30. Bohr Diagrams To draw Bohr Diagrams: 1.Draw the nucleus as a solid circle. 2.Put the number of protons (atomic number) in the nucleus with the number of neutrons (atomic mass – atomic number) under it. 3.Place the number of electrons (same as protons) in orbits around the nucleus by drawing circles around the nucleus. Remember, 1st shell – 2 electrons, 2nd shell – 8 electrons, 3rd shell – 8 electrons, 4th shell – 18 electrons.
- 31. Valency & Groups
- 32. Valencies
- 34. Learning Outcomes • • • • • • • • Energy sub-levels. Heisenberg uncertainty principle. Wave nature of the electron. (Non-mathematical treatment in both cases.) Atomic orbitals. Shapes of s and p orbitals. Building up of electronic structure of the first 36 elements. Electronic configurations of ions of s- and p-block elements only. Arrangement of electrons in individual orbitals of p-block atoms.
- 36. Heisenberg • We cannot know both the position and speed of an electron • Therefore we cannot describe how an electron moves in an atom
- 37. . Einstein • . •
- 38. De Broglie • Matter has wave characteristics
- 39. . 2-slit expt.
- 40. Expected Result if light and : electrons are particles
- 41. Actual result for light and electrons : – demonstrates their wavelike nature
- 42. Electrons were both particles and waves Same for all sub-atomic particles Matter exists as particles and waves at the same time.
- 43. The electron as a wave
- 44. Orbital • A region in space where the probability of finding an electron of a particular is high
- 45. Electrons moving
- 46. Electron paths
- 47. Main levels AND THE NUMBER OF ELECTRONS • • • • 1 = 2e 2 = 8e 3 = 18e 4 = 32e
- 48. Sub-levels • • • • • • Each main level has sub-levels 1has s sub-level only 2 has s and p sub-levels 3 has s,p and d sub-levels 4 has s,p,d and f sub-levels Energy of sub-levels spd
- 49. 1s
- 50. 2s
- 51. 2p
- 52. 3d
- 53. Electrons in sub-levels • • • • s = 2e p = 6e d = 10e f = 14e
- 54. Sub-levels • 1 = s(2e) • 2 = s(2e) + p(6e) = 8e • 3 = s(2e) + p(6e) + d(10e) = 18e
- 56. The "p" orbital is dumb belled shaped and each P sub level is made of three "p" orbitals (because the P sub level can hold 6 electrons and every orbital holds 2 electrons)
- 57. P-orbitals
- 58. P-orbitals
- 59. Electrons in orbitals • S holds 2e • 3 p orbitals each holds only 2e • 5 d orbitals each holds only 2e
- 60. Pauli’s exclusion principle • Orbital can only hold 2electrons and these electrons must have opposite spins
- 62. Aufbau principle • Electrons fill levels in a specific order. • 1s 2s 2p 3s 3p 4s 3d 4p
- 63. AUFBAU
- 66. Hunds rule • When filling up the orbitals in a sublevel electrons fill then singly at first.
- 67. 5 electrons
- 68. 6 electrons Hund’s rule
- 70. Electron Configurations • • • • He, 2, helium : 1s2 Ne, 10, neon: 1s2 2s2 2p6 Ar, 18, argon : 1s2 2s2 2p6 3s2 3p6 Kr, 36, krypton : 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
- 71. Exceptions to Electron configuration rules • Cr • Half-filled orbitals give greater stability • 1s2 2s2 2p6 3s2 3p6 3d4 4s2 1s2 2s2 2p6 3s2 3p6 3d5 4s1 • Cu • Full 3d sub-level gives greater stability • 1s2 2s2 2p6 3s2 3p6 3d9 4s2 1s2 2s2 2p6 3s2 3p6 3d10 4s1
- 72. Electron Configurations (ions) • F-, 10, Flouride: [1s2 2s2 2p6 ]• Cl-, 18, Chloride : [1s2 2s2 2p6 3s2 3p6]• Na+, 10, Sodium ion: [1s2 2s2 2p6 ]+