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MODERN ATOMIC THEORY

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  1. MODERN ATOMIC THEORY Chapter 10

  2. ANCIENT GREEKS’ VIEW OF MATTER • About 400 B.C. , Aristotle thought all matter was made of four “elements” : • earth • air • fire • water

  3. ANCIENT GREEKS’ VIEW OF MATTER • At about the same time another Greek philosopher, Democritus, said that matter was made of tiny, indivisible particles called atoms. • Atomos is the Greek word for indivisible.

  4. Modern View of the Atom • Tiny, dense, positively charged nucleus made up of positive protons and neutral neutrons. • Negatively charged electron shells enclose the nucleus and contain negative electrons.

  5. Atomic Spectra and Bohr One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit. • 1. Any orbit should be possible and so is any energy. • 2. But a charged particle moving in an electric field should emit energy. • End result should be destruction! Electron orbit +

  6. Similarity of Elements • Elements are grouped together in vertical columns (Groups) that have similar properties. • Alkali Metals -- Li, Na, K, Rb, & Cs • Halogens -- F2, Cl2, Br2, & I2 • Noble Gases -- He, Ne, Ar, Kr, Xe, & Rn

  7. Electromagnetic Radiation • Radiant energy that exhibits wave-like behavior and travels through space at the speed of light in a vacuum.

  8. Electromagnetic Radiation wavelength Visible light Amplitude wavelength Node Ultraviolet radiation

  9. Waves • Waves have 3 primary characteristics: • 1.Wavelength:distance between two peaks in a wave. • 2.Frequency:number of waves per second that pass a given point in space. • 3.Speed:speed of light is 2.9979  108 m/s.

  10. As the wavelength () decreases, the frequency () increases.

  11. The electromagnetic spectrum.

  12. Wavelength and frequency can be interconverted. • =c/ • = frequency (s1, Hz, cyc/s, or waves/s) • = wavelength (m) • c= speed of light (m/s)

  13. Huygens thought light travels as waves, while Newton believed it travels as particles.

  14. Photons • Photons -- tiny particle of electromagnetic radiation -- a bundle of light energy. • Ground state -- electrons are at their lowest energy state in an atom. • Excited state -- electrons have absorbed energy by jumping up to a higher energy state in the atom.

  15. Larger energy jumps by electrons produce shorter wavelength (more energetic) light.

  16. Line Spectra of Excited Atoms • Visible lines in H atom spectrum are called the BALMER series. High E Short  High  Low E Long  Low 

  17. Line Spectra of Excited Atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element.

  18. Atomic Spectrum of Hydrogen • Continuous spectrum:Containsallthe wavelengths of light. • Bright Line (discrete) spectrum:Containsonly someof the wavelengths of light.

  19. The diagrams above present evidence for discrete energy levels about a nucleus. Electrons can only be found in certain energy levels with certain energies.

  20. Atomic Line Spectra and Niels Bohr • Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the BRIGHT LINE SPECTRA of excited atoms. Niels Bohr (1885-1962)

  21. Bohr’s Model • Bohr’s Model was incorrect. • Replaced by QUANTUM or WAVE MECHANICS MODEL. • e- can only exist in certain discrete orbitals. • e- is restricted to QUANTIZED energy states. • e- can not be exactly located--location based upon probability.

  22. Quantum or Wave Mechanics • de Broglie (1924) proposed that all moving objects have wave properties. L. de Broglie (1892-1987)

  23. Quantum or Wave Mechanics • Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. E. Schrodinger 1887-1961

  24. Failure of the Bohr Model The Bohr Model of the atom paved the way for the Quantum Mechanical Theory, but current theory is in no way derived from the Bohr Model of the atom. The Bohr Model of the Atom was fundamentally incorrect-- atoms do not move in circular orbits about the nucleus.

  25. 1s Orbital

  26. 2s Orbital

  27. p Orbitals The three p orbitals lie 90o apart in space A p orbital

  28. 2px Orbital

  29. 2py Orbital

  30. 2pz Orbital

  31. 3px Orbital

  32. 3dxy Orbital

  33. 3dxz Orbital

  34. 3dyz Orbital

  35. 3dyz Orbital

  36. 3dx2- y2 Orbital

  37. Quantum Numbers (QN) • 1.Principal QN (n = 1, 2, 3, . . .) - related to size and energy of the orbital. • 2.Angular Momentum QN -- l(s, p, d, & f) - relates to shape of the orbital. • 3.Magnetic QN -- ml(x, y, or z plane) - relates to orientation of the orbital in space relative to other orbitals. • 4.Electron Spin QN -- ms(+1/2, 1/2) - relates to the spin states of the electrons-- clockwise or counterclockwise.

  38. Electron Arrangement • Level Sublevel # Orbitals # electrons • 1-7 s 1 2 • 2-7 p 3 6 • 3-7 d 5 10 • 4-7 f 7 14

  39. Energy Levels and Orbitals • n = the number of the energy level. • n2 = the number of orbitals in an energy level. • 2n2 = the number of electrons in an energy level.

  40. Pauli Exclusion Principle • In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms). • Therefore, an orbital can hold only two electrons, and they must have opposite spins.

  41. Aufbau Principle • As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.

  42. Electron Filling Order--Aufbau

  43. Hund’s Rule • The lowest energy configuration for an atom is the one having themaximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. Orbitals half-fill before they completely fill.

  44. Writing Atomic Electron Configurations S • Two ways of writing configs. One is called the electron configuration notation. Electron configuration notation for H, atomic number = 1 1 no. of s 1 electrons value of l value of n Electron-dot symbol is H.

  45. Writing Atomic Electron Configurations • Two ways of writing configs. Other is called the orbital box notation. Quantum numbers are an energy address instead of a positional address. Electron-dot symbol is He:

  46. Lithium • Group 1A • Atomic number = 3 • 1s22s1 ---> 3 total electrons • Li.

  47. Beryllium • Group 2A • Atomic number = 4 • 1s22s2 ---> 4 total electrons • Be:

  48. Boron • Group 3A • Atomic number = 5 • 1s2 2s2 2p1 ---> • 5 total electrons

  49. Carbon • Group 4A • Atomic number = 6 • 1s2 2s2 2p2 ---> • 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

  50. Nitrogen • Group 5A • Atomic number = 7 • 1s2 2s2 2p3 ---> • 7 total electrons